SCHEME OF WORK
WEEK:
1. Revision (i) saturated hydrocarbons: alkaline e.g methane preparation, properties and uses. (ii) Isomerism. IUPAC nomenclature.
2. Unsaturated hydrocarbons Alkenes e.g ethane, (C2H4) – Nomenclature, preparation, properties and uses.
3. Unsaturated hydrocarbons: Alkynes e.g (C2H4) – Nomenclature, preparation, properties and uses.
4. (i). Aromatics hydrocarbons (ii). Benzene: Structure, properties and uses (iii). Derivative of benzene e.g methy benzene.
5. (A). Alkanols: Sources, general molecular formula, nomenclature, classification, types, preparation, properties and uses. (B).Test for Alkanols.
6. (A). Alkanoic (carbonxylic) acids: sources, nomenclature, structure, preparation, properties and uses. (B). Alkanoates: general moleculear formula, nomenclature, preparation and uses.
7. Fats and oils as higher esters, sources, properties and uses. Detergents and soap: structure, their mode and action.
8. Natural and Synthesis Polymers. (i). Polymerization (Addition and condensation). (ii) Plastics: Thermoplastic and thermosetting polymers. (ii) Resins.
9. (a) Carbohydrates: Sources, general molecules formula, classification, properties and uses. Test for carbohydrate. (b) Biotechnology: Food processing, fermentation including production of garri, bread etc.
POTEINS: source, structure, properties and uses. Test for protein.
10. Amines and Amides: General molecular structure preparation, properties and uses.
11. Revision
1ST TERM
WEEK 1
TOPIC: STRUCTURE OF AN ATOM
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Define an atom and draw a simple atom.
b. Show the gross feature of an atom diagrammatically.
c. Describe sub-atomic particles: protons, neutrons, electrons, their differences charge, mass and functions.
d. Explain atomic number/proton number, number of neutrons, isotopes and isotopy, atomic mass.
e. Define and explain relative atomic mass (Ar) and relative molecular mass (Mr) based on carbon-12 scale.
f. Calculate the relative molecular masses (Mr) OF COMPOUNDS.
g. Explain atoms, molecules and ions and how the three are related.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
EVALUATION:
1. Define an atom; sketch a simple atom showing the nucleus and electrons.
2. State the functions of proton, neutrons, electrons in an atom, their masses, charges and location.
3. Define atomic number (proton number), mass number, number of neutrons and the calculations involving the three of them.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: ISOTOPY AND ATOMIC MASS
EVALUATION:
1. Explain and define relative atomic mass (Ar) and relative molecular mass (Mr) based on carbon-12 scale.
ASSIGNMENT:
1. Calculate the Mr of
a. CUSO4.5H2O
b. CaSO4
c. NaCl
(where Ar: Cu = 63.5, S= 32; O= 16;H=1;Ca=40)
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: ATOMS, MOLECULES AND IONS
EVALUATION:
I. Define and explain atoms, molecules and ions (anions and cations).
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Define an atom and draw a simple atom.
b. Show the gross feature of an atom diagrammatically.
c. Describe sub-atomic particles: protons, neutrons, electrons, their differences charge, mass and functions.
d. Explain atomic number/proton number, number of neutrons, isotopes and isotopy, atomic mass.
e. Define and explain relative atomic mass (Ar) and relative molecular mass (Mr) based on carbon-12 scale.
f. Calculate the relative molecular masses (Mr) OF COMPOUNDS.
g. Explain atoms, molecules and ions and how the three are related.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
EVALUATION:
1. Define an atom; sketch a simple atom showing the nucleus and electrons.
2. State the functions of proton, neutrons, electrons in an atom, their masses, charges and location.
3. Define atomic number (proton number), mass number, number of neutrons and the calculations involving the three of them.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: ISOTOPY AND ATOMIC MASS
EVALUATION:
1. Explain and define relative atomic mass (Ar) and relative molecular mass (Mr) based on carbon-12 scale.
ASSIGNMENT:
1. Calculate the Mr of
a. CUSO4.5H2O
b. CaSO4
c. NaCl
(where Ar: Cu = 63.5, S= 32; O= 16;H=1;Ca=40)
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: ATOMS, MOLECULES AND IONS
EVALUATION:
I. Define and explain atoms, molecules and ions (anions and cations).
WEEK 2
TOPIC: NUCLEAR CHEMISTRY
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Define radioactivity and list its characteristics
b. List the three main types of radiation and their properties/characteristics
c. Define x-rays
d. Explain the methods of detecting radiation
e. Define half life
f. List the uses of radioisotopes and explain.
g. Explain artificial transmutation
h. Use of Einstein's equation E=mc2 in radioactivity
i. Define binding energy
j. Define and differentiate between nuclear fission and nuclear equation
k. Explain the principle of nuclear reactor based on fission reactions.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
Nuclear Chemistry: is the study of the nuclear reactions taking place in the nucleus/nuclei of a atom(s) of element(s).
Radioactivity: is the spontaneous emission of radiation by an element. Such an element is called a RADIOACTIVE ELEMENT.
The isotopes of an element may be stable and radioactive isotopes e.g 12C and 14C are stable and radioactive isotopes of carbon respectively.
CHARACTERISTICS OF RADIOACTIVITY
1. A radioactive substance emits radiation continually and spontaneously.
2. Temperature and pressure have no effect on the rate at which this radiation is emitted.
3. Radioactive radiation, unlike rays, can penetrate through opaque matter.
4. However, like visible light rays, radioactive radiation affects photographic plates.
5. Radioactive radiation ionizes the gases through which it passes, causing fluorescence in certain substances e.g Zinc sulphide, and leaving tracks in a cloud chamber.
6. Radioactivity is always associated with a release of energy which is about a million times as great as that liberated during any chemical reaction.
Radioactivity energy (i.e. energy released by radioactivity) is called NUCLEAR ENERGY.
TYPES OF RADIATION
Radioactive radiation consists of three main components with different penetrating power namely
a. Alpha-rays (α-rays)
b. Beta-rays (B-rays)
c. Gamma rays (r rays). These three rays can be separated and identified by their behaviour in an electrostatic field.
ALPHA-RAY BETA-RAY GAMMA-RAY
Nature Helium nuclei; 42He Fast moving stream of electrons 0-1e Electromagnetic radiation
Electric charge +2 -1 No charge
Mass 4 units 1/1840 units No mass
Velocity About 1/20 the speed of light Varies (from 3-99% of the speed of light) Speed of light
Relative penetration 1 100 10000
Absorber Thin paper Aluminum foil (metal paper) Large lead block
Deflection in an electrostatic field Lightly deflected towards the negative plate Strongly deflected towards the positive plate Not deflected at all
EVALUATION:
1. Define radioactivity and radioactive element.
2. Give examples of stable and unstable isotopes of an element.
3. List the characteristics of radioactivity
4. List the three types of radioactive radiation.
5. List the characteristic differences in the three main types of radioactive radiation
ASSIGNMENT:
a. Show with the aid of diagram, how an electrostatic field can be used to separate the three main types of radiation.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: RADIOACTIVITY
X-RAYS: X-rays are electromagnetic waves, like visible light, but with a shorter wavelength and are produced by the movement of electrons from the shells of an atom. X-rays can penetrate easily through most solid substances which are opaque to visible light, such as metal foils, flesh, wood and paper. Hard x-rays have a greater penetrating power than soft x-rays. Soft x-rays are used in medicine to photograph human body parts (produce a shadow photograph of the bones). Hard x-rays are used for destroying cancerous cells.
METHODS OF DETECTING RADIOACTIVE RADIATION
There are 3 methods of detecting radioactive radiation:
a. The Gieger-muller counter: is based on the ionizing effect of radiation on gases. This produces pulse current which is amplified and may be detected with
(i) audible click from a loud speaker
(ii) Movement of a rate meter
(iii) Reading recorded in the register of a scaler
b. Scintillation counter: This uses the fluorescence of certain minerals (e.g Zinc sulphide) which glows when exposed to radiation and this glow contains tiny flashes of light or scintillations which may be seen under a microscope or counted with a suitable device.
c. Diffusion cloud chamber: This is used for detecting alpha and beta particles which are allowed to pass through a gas super-saturated with water vapour; this forms ions which look like visible vapour trail that can be photographed.
Radioactive disintegration
This is the spontaneous decay of the nucleus of an atom. The original substance undergoing decay is called the parent nucleus and it decays to form the daughter nucleus. This process is called TRASMUTATION of an atom. Rate of decay depends on the radioactive material and varies widely from material to material. During decay, either an alpha particle or beta particle is emitted. Sometimes, gamma rays accompany the emission.
ALPHA DECAY: Involves a radioactive substance decaying to produce helium nucleus 4 2 HE, called alpha particle, by losing two protons and four units of mass.
EVALUATION:
a. Define x-ray and explain soft and hard x-rays with their uses.
b. List and explain the three methods of detecting radiation
c. Explain radioactive decay and use equations to represent alpha and beta rays.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: RATE OF RADIOACTIVE DECAY (HALF-LIFE) AND USES OF RADIOACTIVITY
HALF-LIFE: The half-life of a radioactive element is the time taken for half of the original radioactive substance to decay e.g. half-life (t1/2) of Ra-226 is 1622 years; if we have 10g of Ra-226 at the beginning, then at the end of 1622 years we will have 5g of Ra-226 left.
USES OF RADIOACTIVITY:
1. Medical uses in curing cancer using 60-60. Iodine-131 is used in treating cancer of thyroid gland and P-32 for treating leukemia.
2. Sterilization of equipment (e.g. surgical equipment) makes use of gamma radiation
3. Industrial uses where beta and gamma radiations are used to monitor and control thickness of sheet material and also to test for leaks in pipes.
4. Agricultural purposes such as inducing mutations (modifying genetic constitution) in plants and animals to obtain new and improved varieties. Radioactivity is also used in insect and pest control.
5. Radioactive tracers in tracing metabolic processes such as photosynthesis.
6. Dating techniques such as using carbon-14 to determine the age of archaeological excavations.
EVALUATION:
a) Define half-life and solve calculations on half-life
b) List and explain the uses of radioactivity
ASSIGNMENT:
a. The half-life of an element X is 15days. If we have 5g of x initially, what is the mass of X after (i) 5 days (ii) 20 days (iii) 40 days?
TOPIC: NUCLEAR FISSION AND NUCLEAR FUSION
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
Nuclear Fission is a process in which the nucleus of a heavy element is split into two nuclei of nearly equal mass with a release of energy and radiation.
Nuclear fission is used in making atomic or fission bomb and the atomic pile or nuclear reactor.
Nuclear Fusion is a process in which two or more light nuclei fuse or combine to form a heavier nucleus with a release of energy.
Nuclear fusion is also called THERMONUCLEAR REACTION because large energy is given to the positively charged nuclei to overcome the strong repulsion between them; and this is believed to be the source of energy of the sun and stars.
Mass Energy Relationship
This can be explained by using Einstein's equation in his law of relativity i.e.,
E = mc2
E= energy; m = mass; c= speed of light = 3 x 108 ms-1
EVALUATION:
i) Explain nuclear fission and nuclear fusion with examples shown with equations
ii) Write the formula expressing the relationship between energy and mass
ASSIGNMENT:
a) Explain what a chain reaction is
b) How much energy is released when 0.004g of mass is annihilated in a nuclear reaction?
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Define radioactivity and list its characteristics
b. List the three main types of radiation and their properties/characteristics
c. Define x-rays
d. Explain the methods of detecting radiation
e. Define half life
f. List the uses of radioisotopes and explain.
g. Explain artificial transmutation
h. Use of Einstein's equation E=mc2 in radioactivity
i. Define binding energy
j. Define and differentiate between nuclear fission and nuclear equation
k. Explain the principle of nuclear reactor based on fission reactions.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
Nuclear Chemistry: is the study of the nuclear reactions taking place in the nucleus/nuclei of a atom(s) of element(s).
Radioactivity: is the spontaneous emission of radiation by an element. Such an element is called a RADIOACTIVE ELEMENT.
The isotopes of an element may be stable and radioactive isotopes e.g 12C and 14C are stable and radioactive isotopes of carbon respectively.
CHARACTERISTICS OF RADIOACTIVITY
1. A radioactive substance emits radiation continually and spontaneously.
2. Temperature and pressure have no effect on the rate at which this radiation is emitted.
3. Radioactive radiation, unlike rays, can penetrate through opaque matter.
4. However, like visible light rays, radioactive radiation affects photographic plates.
5. Radioactive radiation ionizes the gases through which it passes, causing fluorescence in certain substances e.g Zinc sulphide, and leaving tracks in a cloud chamber.
6. Radioactivity is always associated with a release of energy which is about a million times as great as that liberated during any chemical reaction.
Radioactivity energy (i.e. energy released by radioactivity) is called NUCLEAR ENERGY.
TYPES OF RADIATION
Radioactive radiation consists of three main components with different penetrating power namely
a. Alpha-rays (α-rays)
b. Beta-rays (B-rays)
c. Gamma rays (r rays). These three rays can be separated and identified by their behaviour in an electrostatic field.
ALPHA-RAY BETA-RAY GAMMA-RAY
Nature Helium nuclei; 42He Fast moving stream of electrons 0-1e Electromagnetic radiation
Electric charge +2 -1 No charge
Mass 4 units 1/1840 units No mass
Velocity About 1/20 the speed of light Varies (from 3-99% of the speed of light) Speed of light
Relative penetration 1 100 10000
Absorber Thin paper Aluminum foil (metal paper) Large lead block
Deflection in an electrostatic field Lightly deflected towards the negative plate Strongly deflected towards the positive plate Not deflected at all
EVALUATION:
1. Define radioactivity and radioactive element.
2. Give examples of stable and unstable isotopes of an element.
3. List the characteristics of radioactivity
4. List the three types of radioactive radiation.
5. List the characteristic differences in the three main types of radioactive radiation
ASSIGNMENT:
a. Show with the aid of diagram, how an electrostatic field can be used to separate the three main types of radiation.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: RADIOACTIVITY
X-RAYS: X-rays are electromagnetic waves, like visible light, but with a shorter wavelength and are produced by the movement of electrons from the shells of an atom. X-rays can penetrate easily through most solid substances which are opaque to visible light, such as metal foils, flesh, wood and paper. Hard x-rays have a greater penetrating power than soft x-rays. Soft x-rays are used in medicine to photograph human body parts (produce a shadow photograph of the bones). Hard x-rays are used for destroying cancerous cells.
METHODS OF DETECTING RADIOACTIVE RADIATION
There are 3 methods of detecting radioactive radiation:
a. The Gieger-muller counter: is based on the ionizing effect of radiation on gases. This produces pulse current which is amplified and may be detected with
(i) audible click from a loud speaker
(ii) Movement of a rate meter
(iii) Reading recorded in the register of a scaler
b. Scintillation counter: This uses the fluorescence of certain minerals (e.g Zinc sulphide) which glows when exposed to radiation and this glow contains tiny flashes of light or scintillations which may be seen under a microscope or counted with a suitable device.
c. Diffusion cloud chamber: This is used for detecting alpha and beta particles which are allowed to pass through a gas super-saturated with water vapour; this forms ions which look like visible vapour trail that can be photographed.
Radioactive disintegration
This is the spontaneous decay of the nucleus of an atom. The original substance undergoing decay is called the parent nucleus and it decays to form the daughter nucleus. This process is called TRASMUTATION of an atom. Rate of decay depends on the radioactive material and varies widely from material to material. During decay, either an alpha particle or beta particle is emitted. Sometimes, gamma rays accompany the emission.
ALPHA DECAY: Involves a radioactive substance decaying to produce helium nucleus 4 2 HE, called alpha particle, by losing two protons and four units of mass.
EVALUATION:
a. Define x-ray and explain soft and hard x-rays with their uses.
b. List and explain the three methods of detecting radiation
c. Explain radioactive decay and use equations to represent alpha and beta rays.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT: RATE OF RADIOACTIVE DECAY (HALF-LIFE) AND USES OF RADIOACTIVITY
HALF-LIFE: The half-life of a radioactive element is the time taken for half of the original radioactive substance to decay e.g. half-life (t1/2) of Ra-226 is 1622 years; if we have 10g of Ra-226 at the beginning, then at the end of 1622 years we will have 5g of Ra-226 left.
USES OF RADIOACTIVITY:
1. Medical uses in curing cancer using 60-60. Iodine-131 is used in treating cancer of thyroid gland and P-32 for treating leukemia.
2. Sterilization of equipment (e.g. surgical equipment) makes use of gamma radiation
3. Industrial uses where beta and gamma radiations are used to monitor and control thickness of sheet material and also to test for leaks in pipes.
4. Agricultural purposes such as inducing mutations (modifying genetic constitution) in plants and animals to obtain new and improved varieties. Radioactivity is also used in insect and pest control.
5. Radioactive tracers in tracing metabolic processes such as photosynthesis.
6. Dating techniques such as using carbon-14 to determine the age of archaeological excavations.
EVALUATION:
a) Define half-life and solve calculations on half-life
b) List and explain the uses of radioactivity
ASSIGNMENT:
a. The half-life of an element X is 15days. If we have 5g of x initially, what is the mass of X after (i) 5 days (ii) 20 days (iii) 40 days?
TOPIC: NUCLEAR FISSION AND NUCLEAR FUSION
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
Nuclear Fission is a process in which the nucleus of a heavy element is split into two nuclei of nearly equal mass with a release of energy and radiation.
Nuclear fission is used in making atomic or fission bomb and the atomic pile or nuclear reactor.
Nuclear Fusion is a process in which two or more light nuclei fuse or combine to form a heavier nucleus with a release of energy.
Nuclear fusion is also called THERMONUCLEAR REACTION because large energy is given to the positively charged nuclei to overcome the strong repulsion between them; and this is believed to be the source of energy of the sun and stars.
Mass Energy Relationship
This can be explained by using Einstein's equation in his law of relativity i.e.,
E = mc2
E= energy; m = mass; c= speed of light = 3 x 108 ms-1
EVALUATION:
i) Explain nuclear fission and nuclear fusion with examples shown with equations
ii) Write the formula expressing the relationship between energy and mass
ASSIGNMENT:
a) Explain what a chain reaction is
b) How much energy is released when 0.004g of mass is annihilated in a nuclear reaction?
WEEK 3
TOPIC: ELECTRONIC ENERGY LEVELS
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Describe the nature of light as dual nature
b) Explain how light behaves like a wave
c) Use equation to show the relationship between the velocity of light and wavelength with frequency
d) Explain the particulate nature of light which is known as the Quantum Theory of light.
e) Explain photon and planck's constant (6.62 x 10-34 js) from the equation Ephoton = hv
f) Use the wave mechanics model to explain Heisenberg's uncertainty principle and also explain an ORBITAL
g) Explain the four quantum numbers
h) State Paulieclusion principle
i) State Hund's rule
j) Draw orbital box diagram showing electron sublevels
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
Nature of Light: Light has a dual nature. This means that light can behave as a wave (wave theory) and as a particle (Quantum theory)
Light as a wave: The wave theory explained the properties of light such as
1) Reflection
2) Refraction due to a change in the speed of light as it proceeds from one medium to another
3) Diffraction which is the spreading of light waves round corners and through apertures
4) Interference which is the superimposition of light waves
5) Polarization
Experiments showed that light must be an electromagnetic wave. X-rays, visible light, infra red rays, UV rays and radio waves are all forms of electromagnetic radiation that travel through space as waves at a constant velocity, C (the velocity of light, C=3.00 x 108ms-1). These waves have wavelength and frequency related to each other by the equation c = ⋋v
Light As a Particle (Quantum theory Of Light)
Apart from wave properties of light, light also have particle properties. Light can thus behave as if it were composed of tiny packets or quanta, of energy called photons. When a photon collides with an electron, the electron accepts energy from it and if this energy is sufficiently large, the electron becomes removed from the surface of the metal. The energy of the photon emitted or absorbed by a substance is proportional to the frequency of light, v i.e, Ephoton = hv
Where h is called PLANCK'S CONSTANT and is equal to 6.6 x 10-34Js. This is the quantum theory of light.
EVALUATION:
1) Describe the dual nature of light
2) List the wave characteristics of light
3) Explain the quantum theory of light
ASSIGNMENT:
1. Define photon
2. If the velocity of light is 3 x 108ms-1 and wavelength is 1000m, calculate the frequency showing how the unit attached is obtained in your working.
CONTENT: WAVE MECHANICS MODEL
Experiments have shown the wave mechanics model to assume that electrons have wave-like properties which make electrons to be very elusive. This led to the uncertainty principle postulated by Heisenberg (in 1927).
THE UNCERTAINTY PRINCIPLE states that if we measure the momentum of an electron accurately, then we cannot know its position with certainty and vice-versa.
Electrons move around the nucleus at certain energy levels of circular orbits. These energy levels were identified by principal quantum number, n.
ORBITAL is a region around the nucleus where there is a possibility of finding an electron with a certain amount of energy (that is given by a wave function).
EVALUATION: The students should be able to:
i) State Heisenberg's uncertainty principle.
ii) Define orbits and explain energy levels of electrons
ASSIGNMENT:
i) Define orbital
ii) Explain principal quantum number, n.
CONTENT: THE FOUR QUANTUM NUMBERS
The energy of an electron may be characterized by four quantum numbers.
1) The Principal Quantum Number, n, has integral values 1,2,3,4, etc which determines the energy levels or shells. The orbit with n=1 is called the K shell, the orbit with n= 2 is known as the L shell, and so on. The maximum possible number of electrons in a shell is given by 2n2.
2) The Subsidiary or azimuthal quantum number, l, has integral values ranging from 0 to (n-1). The electrons with subsidiary quantum numbers 0,1,2 and 3 are called s,p,d and f electrons respectively. S, p, d and f mean sharp, principal, diffuse and fundamental, as derived from spectroscopic terms. When l=0, the orbital is spherical and is called s-orbital. When l=1, the orbital is dumbbell shaped and is called a p-orbital. When l=2, the orbital is double dumbbell shaped and is called a d-orbital and when l=3, a more complicated f-orbital results.
3) The Magnetic Quantum number, m, has values ranging from -1 through 0 to +1. Thus when l=1 (p-sub shell), m= -1, 0, +1, named Px, Py, Pz.
When l=2, m= -2, -1, 0, +1,+2 for the d sub shell.
4) The Spin Quantum number, s, has values -1/2 and +1/2 which shows that each energy level in a sub shell can hold two electrons of opposite sign
Table: The energy sublevels
Name of electron shell Value of n Values of l Number of values l Number of sublevels Names of the sublevels
K 1 0 One One S
L 2 0 and 1 Two Two S and P
M 3 0,1 and 2 Three Three S, P and D
N 4 0,1,2 and 3 Four Four S,P,D and F
EVALUATION:
i) List the four quantum numbers.
ii) Briefly explain the four quantum numbers and their formular representations.
ASSIGNMENT:
i) Sketch the lobes of P-orbital showing Px, Py and Pz sub-orbital
ii) Explain the importance of spin quantum number.
CONTENT: PAULI EXCLUSION PRINCIPLE
Pauli Exclusion Principle states that two electrons in the same atom cannot have the same values for all four quantum numbers. In other words, two electrons in an atom cannot behave in the same manner e.g.
Element Orbital Electrons
Boron 1S22S22P1
Nitrogen 1S22S22P3
Fluorine 1S22S22P5
Neon 1S22S22P6
HUND'S RULE OF MULTIPLICITY states that the sub-orbitals (e.g. 2Px2Py2Pz) are first singly filled with electrons before the electrons will start pairing up in opposite directions in the sub-orbitals.
EVALUATION:
1) State Pauli exclusion principle.
2) Use box diagram of electron sublevels to explain Pauli exclusion principle.
ASSIGNMENT:
1) State Hund's rule of multiplicity
2) Write the orbital electronic configuration of oxygen and draw the box diagram of its electron sublevels.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Describe the nature of light as dual nature
b) Explain how light behaves like a wave
c) Use equation to show the relationship between the velocity of light and wavelength with frequency
d) Explain the particulate nature of light which is known as the Quantum Theory of light.
e) Explain photon and planck's constant (6.62 x 10-34 js) from the equation Ephoton = hv
f) Use the wave mechanics model to explain Heisenberg's uncertainty principle and also explain an ORBITAL
g) Explain the four quantum numbers
h) State Paulieclusion principle
i) State Hund's rule
j) Draw orbital box diagram showing electron sublevels
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION) AND ADDISON-WESLEY COLLEGE CHEMISTRY.
CONTENT:
Nature of Light: Light has a dual nature. This means that light can behave as a wave (wave theory) and as a particle (Quantum theory)
Light as a wave: The wave theory explained the properties of light such as
1) Reflection
2) Refraction due to a change in the speed of light as it proceeds from one medium to another
3) Diffraction which is the spreading of light waves round corners and through apertures
4) Interference which is the superimposition of light waves
5) Polarization
Experiments showed that light must be an electromagnetic wave. X-rays, visible light, infra red rays, UV rays and radio waves are all forms of electromagnetic radiation that travel through space as waves at a constant velocity, C (the velocity of light, C=3.00 x 108ms-1). These waves have wavelength and frequency related to each other by the equation c = ⋋v
Light As a Particle (Quantum theory Of Light)
Apart from wave properties of light, light also have particle properties. Light can thus behave as if it were composed of tiny packets or quanta, of energy called photons. When a photon collides with an electron, the electron accepts energy from it and if this energy is sufficiently large, the electron becomes removed from the surface of the metal. The energy of the photon emitted or absorbed by a substance is proportional to the frequency of light, v i.e, Ephoton = hv
Where h is called PLANCK'S CONSTANT and is equal to 6.6 x 10-34Js. This is the quantum theory of light.
EVALUATION:
1) Describe the dual nature of light
2) List the wave characteristics of light
3) Explain the quantum theory of light
ASSIGNMENT:
1. Define photon
2. If the velocity of light is 3 x 108ms-1 and wavelength is 1000m, calculate the frequency showing how the unit attached is obtained in your working.
CONTENT: WAVE MECHANICS MODEL
Experiments have shown the wave mechanics model to assume that electrons have wave-like properties which make electrons to be very elusive. This led to the uncertainty principle postulated by Heisenberg (in 1927).
THE UNCERTAINTY PRINCIPLE states that if we measure the momentum of an electron accurately, then we cannot know its position with certainty and vice-versa.
Electrons move around the nucleus at certain energy levels of circular orbits. These energy levels were identified by principal quantum number, n.
ORBITAL is a region around the nucleus where there is a possibility of finding an electron with a certain amount of energy (that is given by a wave function).
EVALUATION: The students should be able to:
i) State Heisenberg's uncertainty principle.
ii) Define orbits and explain energy levels of electrons
ASSIGNMENT:
i) Define orbital
ii) Explain principal quantum number, n.
CONTENT: THE FOUR QUANTUM NUMBERS
The energy of an electron may be characterized by four quantum numbers.
1) The Principal Quantum Number, n, has integral values 1,2,3,4, etc which determines the energy levels or shells. The orbit with n=1 is called the K shell, the orbit with n= 2 is known as the L shell, and so on. The maximum possible number of electrons in a shell is given by 2n2.
2) The Subsidiary or azimuthal quantum number, l, has integral values ranging from 0 to (n-1). The electrons with subsidiary quantum numbers 0,1,2 and 3 are called s,p,d and f electrons respectively. S, p, d and f mean sharp, principal, diffuse and fundamental, as derived from spectroscopic terms. When l=0, the orbital is spherical and is called s-orbital. When l=1, the orbital is dumbbell shaped and is called a p-orbital. When l=2, the orbital is double dumbbell shaped and is called a d-orbital and when l=3, a more complicated f-orbital results.
3) The Magnetic Quantum number, m, has values ranging from -1 through 0 to +1. Thus when l=1 (p-sub shell), m= -1, 0, +1, named Px, Py, Pz.
When l=2, m= -2, -1, 0, +1,+2 for the d sub shell.
4) The Spin Quantum number, s, has values -1/2 and +1/2 which shows that each energy level in a sub shell can hold two electrons of opposite sign
Table: The energy sublevels
Name of electron shell Value of n Values of l Number of values l Number of sublevels Names of the sublevels
K 1 0 One One S
L 2 0 and 1 Two Two S and P
M 3 0,1 and 2 Three Three S, P and D
N 4 0,1,2 and 3 Four Four S,P,D and F
EVALUATION:
i) List the four quantum numbers.
ii) Briefly explain the four quantum numbers and their formular representations.
ASSIGNMENT:
i) Sketch the lobes of P-orbital showing Px, Py and Pz sub-orbital
ii) Explain the importance of spin quantum number.
CONTENT: PAULI EXCLUSION PRINCIPLE
Pauli Exclusion Principle states that two electrons in the same atom cannot have the same values for all four quantum numbers. In other words, two electrons in an atom cannot behave in the same manner e.g.
Element Orbital Electrons
Boron 1S22S22P1
Nitrogen 1S22S22P3
Fluorine 1S22S22P5
Neon 1S22S22P6
HUND'S RULE OF MULTIPLICITY states that the sub-orbitals (e.g. 2Px2Py2Pz) are first singly filled with electrons before the electrons will start pairing up in opposite directions in the sub-orbitals.
EVALUATION:
1) State Pauli exclusion principle.
2) Use box diagram of electron sublevels to explain Pauli exclusion principle.
ASSIGNMENT:
1) State Hund's rule of multiplicity
2) Write the orbital electronic configuration of oxygen and draw the box diagram of its electron sublevels.
WEEK 4
TOPIC: PERIODIC CHEMISTRY
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Define and explain periodicity
b. Define and explain periodic law
c. Explain groups and periods
d. Draw the periodic table of the first twenty elements
e. Explain periodic trends and the periodic table.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: PERIODICITY AND THE PERIODIC LAW
PERIODICITY of the elements is the occurrence of successive groups of elements showing strong similar chemical properties because of their similar outer electron shells (e.g. 3L1 (2,1), 11Na (2,8,1) and 19K (2,8,8,1) all belong to group I because of ie- in the outermost shell and they will all show similar chemical properties.
THE PERIODIC LAW: states that the properties of the elements are a periodic function of their atomic numbers.
Atomic number (proton number) of a neutral atom is equal to the number of electrons and this can be arranged in terms of electronic configuration. Going by this, the elements with the same valence electrons will fall into the same group while those with the same number of shells will fall in the same period in order of increasing atomic number.
The GROUPS are vertical columns of elements in the periodic table containing elements with the same number of shells in order of increasing atomic number. There are seven periods labeled PERIODS 1 to 7.
EVALUATION:
1. Define and explain periodicity
2. Define and explain periodic law
3. Define and explain groups and periods.
ASSIGNMENT:
1) Write a periodic table of the first twenty elements.
2) Draw the electronic structures of beryllium and magnesium and use these to show how they belong to the same group.
CONTENT: PERIODIC TRENDS (ATOMIC RADIUS)
The ATOMIC RADIUS of an element is the size of the atom of that element. Two trends are noticeable with atomic radius in the periodic table:
a) Atomic radius increases down a group because of increase in the number of shells down a group while still maintaining same valence electron.
b) Atomic radius decreases across a period from left to right because the number of shells remain the same but the valence electrons increase, leading to a greater attractive force to the nucleus thereby reducing the size of the atom from left to right across the same period.
EVALUATION:
i) Define atomic radius
ii) Explain the two trends noticeable about the atomic radius in the periodic table.
ASSIGNMENT:
(1) An element x has atomic number 11 and mass number of 23.
a. How many neutrons does an atom of x contain?
b. Write the electric configuration in an atom of x
c. Will x react in an electrovalent or covalent manner?
d. Give reasons for your answer.
CONTENT: IONIZATION ENERGY
Ionization Energy is the energy required to remove a valence electron from an atom of element to become a positive ion.
Two trends are noticeable for ionization energy:
(a) Ionization energy decreases down a group (because the number of shells increases down a group although the valence electrons remain the same; thus less attraction between valence electrons and nucleus down the group, leading to less energy needed to pull out a valence electron)
(b) Ionization energy increases across a period (because the valence electrons increase across a period although the number of shells remains the same, thus more energy is needed to pull out a valance electron).
ELECTRO-NEGATIVITY
Electro-negativity is the ability of an atom to gain electron to become negatively charged.
(a) Electro-negativity increases from left to right across a period (because valence electrons increase and the tendency is to gain electrons to attain stability as valence electrons increase).
(b) Electro-negativity decreases down a group (because the number of shells increases down a group and this means less attraction between the nucleus and outer shell, thus decreasing electro-negativity).
The most electronegative group is Group VII known as the HALOGENS (fluorine, chlorine, bromine, iodine, astatine).
EVALUATION:
1. Explain ionization energy
2. Explain the two trends noticeable about ionization energy in the periodic table.
ASSIGNMENT:
1. Define electro-negativity
2. Relate the electro-negativity to electron affinity
3. Explain the trends noticeable about electro-negativity in the periodic table.
CONTENT: ELECTRO-POSITIVITY
Electro-positivity is the ability of an atom to lose electrons to become positively charged.
a) Electro-positivity increases down a group (because the number of shells increases down a group leading to less attractive force between nucleus and valence electrons, thus making it easier to lose electron down a group.
b) Electro-positivity decreases across a period from left to right (because the number of shells remain the same while the valence electrons increase, leading to greater attractive force between nucleus and valence electrons, thus making it more difficult to lose electrons across a period from left to right).
The most electro-positivity group is the Group I elements known as the ALKALI METALS (Lithium, Sodium, Potassium, Rubidium, Cesium, Francium).
After the group I elements, the next most electro-positive group is the Group II elements called the ALKALINE EARTH METALS (beryllium, magnesium, calcium, strontium, barium, radium).
EVALUATION:
i) Define and explain electro-positivity
ii) Explain the trends noticeable about electro-positivity in the periodic table of elements.
ASSIGNMENT:
1. Explain why Group I elements are the most electro-positive elements.
2. Explain why Group VII elements are the most electro-negative elements.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Define and explain periodicity
b. Define and explain periodic law
c. Explain groups and periods
d. Draw the periodic table of the first twenty elements
e. Explain periodic trends and the periodic table.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: PERIODICITY AND THE PERIODIC LAW
PERIODICITY of the elements is the occurrence of successive groups of elements showing strong similar chemical properties because of their similar outer electron shells (e.g. 3L1 (2,1), 11Na (2,8,1) and 19K (2,8,8,1) all belong to group I because of ie- in the outermost shell and they will all show similar chemical properties.
THE PERIODIC LAW: states that the properties of the elements are a periodic function of their atomic numbers.
Atomic number (proton number) of a neutral atom is equal to the number of electrons and this can be arranged in terms of electronic configuration. Going by this, the elements with the same valence electrons will fall into the same group while those with the same number of shells will fall in the same period in order of increasing atomic number.
The GROUPS are vertical columns of elements in the periodic table containing elements with the same number of shells in order of increasing atomic number. There are seven periods labeled PERIODS 1 to 7.
EVALUATION:
1. Define and explain periodicity
2. Define and explain periodic law
3. Define and explain groups and periods.
ASSIGNMENT:
1) Write a periodic table of the first twenty elements.
2) Draw the electronic structures of beryllium and magnesium and use these to show how they belong to the same group.
CONTENT: PERIODIC TRENDS (ATOMIC RADIUS)
The ATOMIC RADIUS of an element is the size of the atom of that element. Two trends are noticeable with atomic radius in the periodic table:
a) Atomic radius increases down a group because of increase in the number of shells down a group while still maintaining same valence electron.
b) Atomic radius decreases across a period from left to right because the number of shells remain the same but the valence electrons increase, leading to a greater attractive force to the nucleus thereby reducing the size of the atom from left to right across the same period.
EVALUATION:
i) Define atomic radius
ii) Explain the two trends noticeable about the atomic radius in the periodic table.
ASSIGNMENT:
(1) An element x has atomic number 11 and mass number of 23.
a. How many neutrons does an atom of x contain?
b. Write the electric configuration in an atom of x
c. Will x react in an electrovalent or covalent manner?
d. Give reasons for your answer.
CONTENT: IONIZATION ENERGY
Ionization Energy is the energy required to remove a valence electron from an atom of element to become a positive ion.
Two trends are noticeable for ionization energy:
(a) Ionization energy decreases down a group (because the number of shells increases down a group although the valence electrons remain the same; thus less attraction between valence electrons and nucleus down the group, leading to less energy needed to pull out a valence electron)
(b) Ionization energy increases across a period (because the valence electrons increase across a period although the number of shells remains the same, thus more energy is needed to pull out a valance electron).
ELECTRO-NEGATIVITY
Electro-negativity is the ability of an atom to gain electron to become negatively charged.
(a) Electro-negativity increases from left to right across a period (because valence electrons increase and the tendency is to gain electrons to attain stability as valence electrons increase).
(b) Electro-negativity decreases down a group (because the number of shells increases down a group and this means less attraction between the nucleus and outer shell, thus decreasing electro-negativity).
The most electronegative group is Group VII known as the HALOGENS (fluorine, chlorine, bromine, iodine, astatine).
EVALUATION:
1. Explain ionization energy
2. Explain the two trends noticeable about ionization energy in the periodic table.
ASSIGNMENT:
1. Define electro-negativity
2. Relate the electro-negativity to electron affinity
3. Explain the trends noticeable about electro-negativity in the periodic table.
CONTENT: ELECTRO-POSITIVITY
Electro-positivity is the ability of an atom to lose electrons to become positively charged.
a) Electro-positivity increases down a group (because the number of shells increases down a group leading to less attractive force between nucleus and valence electrons, thus making it easier to lose electron down a group.
b) Electro-positivity decreases across a period from left to right (because the number of shells remain the same while the valence electrons increase, leading to greater attractive force between nucleus and valence electrons, thus making it more difficult to lose electrons across a period from left to right).
The most electro-positivity group is the Group I elements known as the ALKALI METALS (Lithium, Sodium, Potassium, Rubidium, Cesium, Francium).
After the group I elements, the next most electro-positive group is the Group II elements called the ALKALINE EARTH METALS (beryllium, magnesium, calcium, strontium, barium, radium).
EVALUATION:
i) Define and explain electro-positivity
ii) Explain the trends noticeable about electro-positivity in the periodic table of elements.
ASSIGNMENT:
1. Explain why Group I elements are the most electro-positive elements.
2. Explain why Group VII elements are the most electro-negative elements.
WEEK 5
TOPIC: PERIODIC GRADATION OF GROUP SEVEN ELEMENTS
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Explain the meaning of halogens.
b) Show how halogens are the most reactive non-metals known.
c) Explain the great similarity in the properties of the halogens.
d) Explain the trends exhibited by the halogens as we go down the group from fluorine to iodine.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: PERIODIC GRADATION OF PROPERTIES OF GROUP SEVEN ELEMENTS (THE HALOGENS)
The Group 7 elements are called the HALOGENS because halogen means SALT FORMER. Halogens are the most reactive non-metals known and they are so reactive with the other non-metals in the environment that in nature, they exist mainly as salts instead of as free elements.
The halogens have the same number of valence electrons (seven) and this makes them to show great similarity in their properties.
SIMILARITIES OF HALOGENS
1) They are all non-metals
2) They exist as diatomic molecules
3) They are coloured
4) They ionize to form univalent negative ions which react with metallic ions to form electrovalent compounds.
5) Their hydrides are covalent gases at room temperature and dissolve readily in water to form acids.
EVALUATION:
1) Explain why Group 7 elements are called halogens.
2) Give the reason why the halogens have great similarity in their properties.
ASSIGNMENT:
Briefly explain the five similarities of the halogens.
CONTENT: TRENDS EXHIBITED BY THE HALOGENS
The trends exhibited by the halogens are as follows:-
1) A change in state (at 25℃) fluorine and chlorine are gases; bromine is a liquid; and iodine is a solid.
2) Their colours become darker; fluorine is yellow; chlorine is greenish yellow; bromine is reddish black; and iodine is black.
3) Their melting and boiling points increase progressively.
4) The ease with which they ionize to form negative ions decrease, i.e. their electronegativity decreases. This is reflected by a decrease in their chemical reactivity.
For example
a. In their reactions with hydrogen to form hydrides - fluorine reacts explosively even in the dark; chlorine slowly diffuse light but explosively in bright light; bromine, slowly in bright light; and iodine slowly and incompletely even in bright light.
b. In their displacement reactions, where each halogen displaces the one following it from a solution of the latter's salt.
STRUCTURAL PROPERTIES OF THE HALOGENS
HALOGEN F2 Cl2 Br2 I2
ATOMIC NUMBER 9 17 35 53
ELECTRON SHELLS K,L K,L,M K,L,M,N K,L,M,N,O
ELECTRONS IN OUTERMOST SHELLS 7 7 7 7
EVALUATION:
(1) Explain the three physical trends shown by the halogens
(2) Explain the chemical trends shown by the halogens (i.e. electronegativity and group displacement reactions).
ASSIGNMENT:
(1) Use chemical equations to show the chemical trend of group displacement reactions as exhibited by the halogens.
CONTENT: PROPERTIES OF CHLORINE AS A TYPICAL HALOGEN
PHYSICAL PROPERTIES:
1) Chlorine is greenish-yellow gas with an unpleasant choking smell.
2) It is moderately soluble in water. About 2.3cm3 of it will dissolve in 1cm3 of water at S.T.P
3) It is about 2.5 times denser than air.
4) It can easily be liquefied under a pressure of about 6 atmospheres.
5) It is poisonous. As little as 20 parts per million of it in the air can damage the mucous lining of our lungs.
EVALUATION:
(1) List the physical properties of chlorine
(2) Explain the solubility and density of chlorine
ASSIGNMENT:
(1) Briefly explain the physical properties of chlorine.
CONTENT: CHEMICAL PROPERTIES OF CHLORINE AS A TYPICAL HALOGEN
1. With its seven valence electrons, chlorine is very reactive and forms negative ion, Cl- by gaining electrons in electrovalent compounds ( e.g. NaCl, CaCl2) or by sharing a pair of electrons in a single covalent bond with another atom of fairly similar electro negativity (e.g. as in Cl-Cl and hydrogen chloride, H-Cl)
2. Displacement reactions down the group e.g.
Cl2(g) + 2NaBr 2Nacl(aq) + Br2(l)
Cl2(g) + 2HI(aq) 2Hcl(aq) + I2(s)
Cl2(g) + 2NaI(aq) 2Nacl(aq) + I2(s)
3. (i) Combination with non metals H2(g) + Cl2(g) 2Hcl(s)
(ii) Combination with metals (directly)
2Nac(S) + Cl2(g) 2Nacl(s)
Zn(s) + Cl2(g) ZnCl2(s)
2Fe(s) + 3Cl2(g) 2FeCl3(s)
4. Reactions with hydrogen
a. With hydrocarbons. C10H16(l) + 8Cl2(g) 10C(s) + 16HCl(g)
b. CH4(g) + Cl2(g) CH3Cl(g) + HCl(g)
c. With ammonia 2NH3(g) + 3Cl2(g) N2(g) + 6HCl(g)
6HCl(g) + 6NH3(g) 6NH4Cl(s)
d. With H2S. H2S(g) + Cl2(g) 2HCl(g) + S(s)
e. With water. Cl2(g) + H2O(l) Hcl(aq) + HOCl(aq)
2HOCl(aq) sunlight 2HCl(aq) + O2(g)
f. As oxidizing agent: 2FeCl2(aq) + Cl2(g) 2FeCl3(aq)
g. As a bleaching agent: HOCl(aq) HCl(aq) + [0]
Dye + [0] (Dye + 0)
EVALUATION:
(i) List the chemical properties of chlorine
(ii) Show with chemical equations, four of the chemical behaviour of chlorine as a typical halogen.
ASSIGNMENT:
1) Comment on the relative reactivities of fluorine, chlorine, bromine and iodine.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Explain the meaning of halogens.
b) Show how halogens are the most reactive non-metals known.
c) Explain the great similarity in the properties of the halogens.
d) Explain the trends exhibited by the halogens as we go down the group from fluorine to iodine.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: PERIODIC GRADATION OF PROPERTIES OF GROUP SEVEN ELEMENTS (THE HALOGENS)
The Group 7 elements are called the HALOGENS because halogen means SALT FORMER. Halogens are the most reactive non-metals known and they are so reactive with the other non-metals in the environment that in nature, they exist mainly as salts instead of as free elements.
The halogens have the same number of valence electrons (seven) and this makes them to show great similarity in their properties.
SIMILARITIES OF HALOGENS
1) They are all non-metals
2) They exist as diatomic molecules
3) They are coloured
4) They ionize to form univalent negative ions which react with metallic ions to form electrovalent compounds.
5) Their hydrides are covalent gases at room temperature and dissolve readily in water to form acids.
EVALUATION:
1) Explain why Group 7 elements are called halogens.
2) Give the reason why the halogens have great similarity in their properties.
ASSIGNMENT:
Briefly explain the five similarities of the halogens.
CONTENT: TRENDS EXHIBITED BY THE HALOGENS
The trends exhibited by the halogens are as follows:-
1) A change in state (at 25℃) fluorine and chlorine are gases; bromine is a liquid; and iodine is a solid.
2) Their colours become darker; fluorine is yellow; chlorine is greenish yellow; bromine is reddish black; and iodine is black.
3) Their melting and boiling points increase progressively.
4) The ease with which they ionize to form negative ions decrease, i.e. their electronegativity decreases. This is reflected by a decrease in their chemical reactivity.
For example
a. In their reactions with hydrogen to form hydrides - fluorine reacts explosively even in the dark; chlorine slowly diffuse light but explosively in bright light; bromine, slowly in bright light; and iodine slowly and incompletely even in bright light.
b. In their displacement reactions, where each halogen displaces the one following it from a solution of the latter's salt.
STRUCTURAL PROPERTIES OF THE HALOGENS
HALOGEN F2 Cl2 Br2 I2
ATOMIC NUMBER 9 17 35 53
ELECTRON SHELLS K,L K,L,M K,L,M,N K,L,M,N,O
ELECTRONS IN OUTERMOST SHELLS 7 7 7 7
EVALUATION:
(1) Explain the three physical trends shown by the halogens
(2) Explain the chemical trends shown by the halogens (i.e. electronegativity and group displacement reactions).
ASSIGNMENT:
(1) Use chemical equations to show the chemical trend of group displacement reactions as exhibited by the halogens.
CONTENT: PROPERTIES OF CHLORINE AS A TYPICAL HALOGEN
PHYSICAL PROPERTIES:
1) Chlorine is greenish-yellow gas with an unpleasant choking smell.
2) It is moderately soluble in water. About 2.3cm3 of it will dissolve in 1cm3 of water at S.T.P
3) It is about 2.5 times denser than air.
4) It can easily be liquefied under a pressure of about 6 atmospheres.
5) It is poisonous. As little as 20 parts per million of it in the air can damage the mucous lining of our lungs.
EVALUATION:
(1) List the physical properties of chlorine
(2) Explain the solubility and density of chlorine
ASSIGNMENT:
(1) Briefly explain the physical properties of chlorine.
CONTENT: CHEMICAL PROPERTIES OF CHLORINE AS A TYPICAL HALOGEN
1. With its seven valence electrons, chlorine is very reactive and forms negative ion, Cl- by gaining electrons in electrovalent compounds ( e.g. NaCl, CaCl2) or by sharing a pair of electrons in a single covalent bond with another atom of fairly similar electro negativity (e.g. as in Cl-Cl and hydrogen chloride, H-Cl)
2. Displacement reactions down the group e.g.
Cl2(g) + 2NaBr 2Nacl(aq) + Br2(l)
Cl2(g) + 2HI(aq) 2Hcl(aq) + I2(s)
Cl2(g) + 2NaI(aq) 2Nacl(aq) + I2(s)
3. (i) Combination with non metals H2(g) + Cl2(g) 2Hcl(s)
(ii) Combination with metals (directly)
2Nac(S) + Cl2(g) 2Nacl(s)
Zn(s) + Cl2(g) ZnCl2(s)
2Fe(s) + 3Cl2(g) 2FeCl3(s)
4. Reactions with hydrogen
a. With hydrocarbons. C10H16(l) + 8Cl2(g) 10C(s) + 16HCl(g)
b. CH4(g) + Cl2(g) CH3Cl(g) + HCl(g)
c. With ammonia 2NH3(g) + 3Cl2(g) N2(g) + 6HCl(g)
6HCl(g) + 6NH3(g) 6NH4Cl(s)
d. With H2S. H2S(g) + Cl2(g) 2HCl(g) + S(s)
e. With water. Cl2(g) + H2O(l) Hcl(aq) + HOCl(aq)
2HOCl(aq) sunlight 2HCl(aq) + O2(g)
f. As oxidizing agent: 2FeCl2(aq) + Cl2(g) 2FeCl3(aq)
g. As a bleaching agent: HOCl(aq) HCl(aq) + [0]
Dye + [0] (Dye + 0)
EVALUATION:
(i) List the chemical properties of chlorine
(ii) Show with chemical equations, four of the chemical behaviour of chlorine as a typical halogen.
ASSIGNMENT:
1) Comment on the relative reactivities of fluorine, chlorine, bromine and iodine.
WEEK 6
TOPIC: ELEMENTS OF THE FIRST TRANSITION SERIES
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Define transition elements in terms of their position in the periodic table
b) Define transition elements in terms of 3d-orbital.
c) List and explain the location of the first transition series.
d) List and explain the properties of the transition elements (first transition series).
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: ELEMENTS OF THE FIRST TRANSITION SERIES
TRANSITION ELEMENTS:
The transition elements are all metals which are of economic importance (e.g. gold and silver are used as jewelries, iron in steel for construction of bridges and building of automobiles; aluminum alloy used in building air crafts). They are found in the d-block of the periodic table between Groups 2 and 3. They occupy 3 rows, with 10 elements in each row.
DEFINITION: Transition elements are metallic elements which have partially filled d orbitals.
FIRST TRANSITION SERIES
The first transition series is from Scandrum (21Sc) to Zinc (30Zn). There are ten d-block metals in the first transition series which occur in Period 4. Potassium and calcium are the s-block metals in this period. The transition metals have very similar properties ad are quite different from the reactive s-block metals of Groups 1 and 2.
First transition series Sc Ti V Cr Mn Fe Co Ni Cu Zn
Number of D electrons 1 2 3 4 5 6 7 8 9 10
EVALUATION:
(i) List the economic importance of the transition elements.
(ii) Describe the location of the transition elements in the periodic table.
(iii) Define transition elements
(iv) Describe the location of the first transition elements in the periodic table and also list the ten transition elements in the first transition series with their atomic numbers.
ASSIGNMENT:
Write the electronic configuration of:
1) Scandrum (21Sc)
2) Zinc (30Zn)
CONTENT: PHYSICAL PROPERTIES OF THE TRANSITION ELEMENTS (FIRST TRANSITION SERIES REFERENCE)
The physical properties of the transition elements are:
1. They are typical metals with high boiling and melting points.
2. They are hard and dense
3. They are lustrous (i.e. can be polished)
4. They are sonorous (i.e. gives a note when hit)
5. They are malleable (i.e. can be beaten into shape)
6. They are ductile (i.e. can be drawn into wire)
7. They are good conductors of heat and electricity.
These properties of the transition metals indicate the presence of strong metallic bonding. This is due to the 3 d electrons available for bonding in the atoms of these metals in addition to the 4 s electrons. Thus, the presence of these extra electrons makes the metallic bonds in the transition elements very strong. In comparison, potassium (Group 1) and calcium (Group II) are soft and have lower melt points, because their atoms have only one or two electrons in their outermost shells for forming metallic bonds. Thus, the metallic bonds in s-block metals are nnot as strong as those in d-block metals. Consequently, the strong metallic bonds and the small sizes of the atoms account for the high melting and boiling points, densities and tensile strengths of the transition metals. The densities of the transition metals increase across the series because their relative atomic masses increase progressively while their atomic sizes remain fairly constant.
EVALUATION: The students should be able to:
1) List seven physical properties of the transition elements
2) Use the 3 d and 4 s electrons of the transition metals to explain their pronounced physical properties over the s-block metals.
ASSIGNMENT:
1. Explain the difference in the atomic size of the s-block and d-block metals.
CONTENT: ELECTRONIC CONFIGURATION OF THE FIRST TRANSITION SERIES/CHEMICAL PROPERTIES
The atoms of the first transition series metals have one or two 4 s electrons like the Groups I and II metals in the same period 4 but in addition, they have partially filled 3 d orbitals which are responsible for the special properties of the transition metals. Exceptions include:
(1) Zinc atom (Zn), Zinc ion (Zn2+), copper atom (Cu) and copper ion (Cu+) which have completely filled 3 d orbitals and
(2) Scandium ion (Sc3+) which does not have any electron in its 3 d orbitals.
Zinc and the forms of copper and scandium given above are usually considered as NON-TRANSITIONAL.
Chromium (24Cr) and copper (29Cu) atoms have only one electron in the 4 s orbital. This is because a special stability is associated with the 3 d5 and 3d10 electronic configurations, i.e. when all the five 3 d orbitals are singly or doubly filled.
CHEMICAL REACTIVITY
Unlike the S-block and P-block elements, the chemical properties of the transition metals in the same period does not vary quite markedly, from left to right because electrons are added progressively to the inner orbitals, not the outermost orbitals, as in the s-block and p-block elements.
The transition elements are not as reactive as the s-block metals because the nuclei of the transition elements have a greater attraction on their electrons than the nuclei of the s-block metals. So the s-block metals have a lower ionization energy and are more reactive than the transition metals.
In general, transition metals are moderately reactive and their reactivities decrease across the series due to a corresponding increase in the value of their ionization energies.
EVALUATION:
1. Explain the electronic configuration of the first transition series metals and show those that are NON-TRANSITIONAL.
2. Explain why the s-block metals are more chemically reactive than the transition metals.
ASSIGNMENT:
Explain why the following are non-transitional, with the aid of electronic orbital box diagrams
30Zn; 30Zn 2+; 29Cu; 29Cu+; 21Sc3+
CONTENT: CHEMICAL PROPERTIES OF THE FIRST TRANSITION SERIES ELEMENTS
The chemical properties of the transition elements are
1) Variable oxidation states
2) Complex ion formation
3) Colour of transition metal ions
4) Catalytic activity
(1) VARIABLE OXIDATION STATES: Transition metals have variable oxidation states because the 3 d electrons are available for bond formation e.g. manganese with the following electronic configuration 1S2 2S2 2P6 3S2 3P6 3d5 4S2 can lose.
(a) Two 4S electrons to give a +2 oxidation state as in MnO3
(b) Two 4S and two 3d electrons to give a +4 oxidation state as in MnO2
(c) Two 4S and four 3d electrons to give a +6 oxidation state as in MnO42-
(d) Two 4S and five 3d electrons to give a +7 oxidation state as in MnO4-
(2) COMPLEX ION FORMATION: Transition metals have a outstanding ability to form complex ions. A complex ion has a central positive ion linked to several other atoms, ions or molecules called LIGANDS. The bonding between the central metal ion and the ligands may be either predominantly electrovalent or predominantly coordinate. An example of a complex ion is the blue tetra amine copper (II) ion, [Cu(NH3)4]2+, in which the central Cu2+ is linked by coordinate bonding to four NH3 molecules in a tetrahedral arrangement. Another example is hexacyanoferrate(III) ion, [Fe(CN)6]3-
(3) COLOUR OF TRANSITION METAL IONS: Transition metallic ions are usually coloured which serves as a useful guide in identifying a compound. The colours are associated with partially filled 3 d orbitals (i.e. 3 d1 to 3d9).
(4) CATALYTIC ACTIVITY: Because transition metals easily change their oxidation states, they are able to act as catalysts e.g. (2H202 MnO2 2H20 + O2); Haber process uses finely divided iron catalyst; contact process uses V2O5; hydrogenation of vegetable oil uses nickel catalyst.
EVALUATION:
(i) List the four chemical properties of transition elements.
(ii) Explain how the 3 d electrons account for the chemical properties of the transition elements.
ASSIGNMENT:
(1) Transition metals are characterized by variable oxidation state, coloured ions, catalytic activity and the formation of complexes. Illustrate this statement using two first transition series metals as examples.
(2) What is a complex ion? Name two metals that form complex ions, giving an example in each case.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Define transition elements in terms of their position in the periodic table
b) Define transition elements in terms of 3d-orbital.
c) List and explain the location of the first transition series.
d) List and explain the properties of the transition elements (first transition series).
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: ELEMENTS OF THE FIRST TRANSITION SERIES
TRANSITION ELEMENTS:
The transition elements are all metals which are of economic importance (e.g. gold and silver are used as jewelries, iron in steel for construction of bridges and building of automobiles; aluminum alloy used in building air crafts). They are found in the d-block of the periodic table between Groups 2 and 3. They occupy 3 rows, with 10 elements in each row.
DEFINITION: Transition elements are metallic elements which have partially filled d orbitals.
FIRST TRANSITION SERIES
The first transition series is from Scandrum (21Sc) to Zinc (30Zn). There are ten d-block metals in the first transition series which occur in Period 4. Potassium and calcium are the s-block metals in this period. The transition metals have very similar properties ad are quite different from the reactive s-block metals of Groups 1 and 2.
First transition series Sc Ti V Cr Mn Fe Co Ni Cu Zn
Number of D electrons 1 2 3 4 5 6 7 8 9 10
EVALUATION:
(i) List the economic importance of the transition elements.
(ii) Describe the location of the transition elements in the periodic table.
(iii) Define transition elements
(iv) Describe the location of the first transition elements in the periodic table and also list the ten transition elements in the first transition series with their atomic numbers.
ASSIGNMENT:
Write the electronic configuration of:
1) Scandrum (21Sc)
2) Zinc (30Zn)
CONTENT: PHYSICAL PROPERTIES OF THE TRANSITION ELEMENTS (FIRST TRANSITION SERIES REFERENCE)
The physical properties of the transition elements are:
1. They are typical metals with high boiling and melting points.
2. They are hard and dense
3. They are lustrous (i.e. can be polished)
4. They are sonorous (i.e. gives a note when hit)
5. They are malleable (i.e. can be beaten into shape)
6. They are ductile (i.e. can be drawn into wire)
7. They are good conductors of heat and electricity.
These properties of the transition metals indicate the presence of strong metallic bonding. This is due to the 3 d electrons available for bonding in the atoms of these metals in addition to the 4 s electrons. Thus, the presence of these extra electrons makes the metallic bonds in the transition elements very strong. In comparison, potassium (Group 1) and calcium (Group II) are soft and have lower melt points, because their atoms have only one or two electrons in their outermost shells for forming metallic bonds. Thus, the metallic bonds in s-block metals are nnot as strong as those in d-block metals. Consequently, the strong metallic bonds and the small sizes of the atoms account for the high melting and boiling points, densities and tensile strengths of the transition metals. The densities of the transition metals increase across the series because their relative atomic masses increase progressively while their atomic sizes remain fairly constant.
EVALUATION: The students should be able to:
1) List seven physical properties of the transition elements
2) Use the 3 d and 4 s electrons of the transition metals to explain their pronounced physical properties over the s-block metals.
ASSIGNMENT:
1. Explain the difference in the atomic size of the s-block and d-block metals.
CONTENT: ELECTRONIC CONFIGURATION OF THE FIRST TRANSITION SERIES/CHEMICAL PROPERTIES
The atoms of the first transition series metals have one or two 4 s electrons like the Groups I and II metals in the same period 4 but in addition, they have partially filled 3 d orbitals which are responsible for the special properties of the transition metals. Exceptions include:
(1) Zinc atom (Zn), Zinc ion (Zn2+), copper atom (Cu) and copper ion (Cu+) which have completely filled 3 d orbitals and
(2) Scandium ion (Sc3+) which does not have any electron in its 3 d orbitals.
Zinc and the forms of copper and scandium given above are usually considered as NON-TRANSITIONAL.
Chromium (24Cr) and copper (29Cu) atoms have only one electron in the 4 s orbital. This is because a special stability is associated with the 3 d5 and 3d10 electronic configurations, i.e. when all the five 3 d orbitals are singly or doubly filled.
CHEMICAL REACTIVITY
Unlike the S-block and P-block elements, the chemical properties of the transition metals in the same period does not vary quite markedly, from left to right because electrons are added progressively to the inner orbitals, not the outermost orbitals, as in the s-block and p-block elements.
The transition elements are not as reactive as the s-block metals because the nuclei of the transition elements have a greater attraction on their electrons than the nuclei of the s-block metals. So the s-block metals have a lower ionization energy and are more reactive than the transition metals.
In general, transition metals are moderately reactive and their reactivities decrease across the series due to a corresponding increase in the value of their ionization energies.
EVALUATION:
1. Explain the electronic configuration of the first transition series metals and show those that are NON-TRANSITIONAL.
2. Explain why the s-block metals are more chemically reactive than the transition metals.
ASSIGNMENT:
Explain why the following are non-transitional, with the aid of electronic orbital box diagrams
30Zn; 30Zn 2+; 29Cu; 29Cu+; 21Sc3+
CONTENT: CHEMICAL PROPERTIES OF THE FIRST TRANSITION SERIES ELEMENTS
The chemical properties of the transition elements are
1) Variable oxidation states
2) Complex ion formation
3) Colour of transition metal ions
4) Catalytic activity
(1) VARIABLE OXIDATION STATES: Transition metals have variable oxidation states because the 3 d electrons are available for bond formation e.g. manganese with the following electronic configuration 1S2 2S2 2P6 3S2 3P6 3d5 4S2 can lose.
(a) Two 4S electrons to give a +2 oxidation state as in MnO3
(b) Two 4S and two 3d electrons to give a +4 oxidation state as in MnO2
(c) Two 4S and four 3d electrons to give a +6 oxidation state as in MnO42-
(d) Two 4S and five 3d electrons to give a +7 oxidation state as in MnO4-
(2) COMPLEX ION FORMATION: Transition metals have a outstanding ability to form complex ions. A complex ion has a central positive ion linked to several other atoms, ions or molecules called LIGANDS. The bonding between the central metal ion and the ligands may be either predominantly electrovalent or predominantly coordinate. An example of a complex ion is the blue tetra amine copper (II) ion, [Cu(NH3)4]2+, in which the central Cu2+ is linked by coordinate bonding to four NH3 molecules in a tetrahedral arrangement. Another example is hexacyanoferrate(III) ion, [Fe(CN)6]3-
(3) COLOUR OF TRANSITION METAL IONS: Transition metallic ions are usually coloured which serves as a useful guide in identifying a compound. The colours are associated with partially filled 3 d orbitals (i.e. 3 d1 to 3d9).
(4) CATALYTIC ACTIVITY: Because transition metals easily change their oxidation states, they are able to act as catalysts e.g. (2H202 MnO2 2H20 + O2); Haber process uses finely divided iron catalyst; contact process uses V2O5; hydrogenation of vegetable oil uses nickel catalyst.
EVALUATION:
(i) List the four chemical properties of transition elements.
(ii) Explain how the 3 d electrons account for the chemical properties of the transition elements.
ASSIGNMENT:
(1) Transition metals are characterized by variable oxidation state, coloured ions, catalytic activity and the formation of complexes. Illustrate this statement using two first transition series metals as examples.
(2) What is a complex ion? Name two metals that form complex ions, giving an example in each case.
WEEK 7
TOPIC: CHEMICAL BONDING
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Define and explain chemical bonding
b) List the types of chemical bonding
c) Explain ionic bonding with its properties
d) Explain covalent bonding with its properties
e) Explain coordinate covalency (dative bonding and its properties.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: CHEMICAL BONDING
The rare gases have very stable electronic configuration which makes them to be thermodynamically and chemically stable. The tendency of the other elements is to try to attain this stable duplet (helium) or octet (neon, argon, krypton) structure possessed by the rare (noble gases. This is achieved during chemical combination. There are two main types of chemical combination (bonding).
1. Electrovalent combination (or ionic bonding)
2. Covalent combination which is classified into:
a) Ordinary covalent combination
b) Coordinate covalent combination
ELECTROVALENT BONDING (IONIC BONDING)
In electrovalent bonding, there is a transfer of outermost shell (valence) electrons from one atom (usually metallic) to another atom (usually non-metallic). Thus, we have atoms which act as donors of electrons and those which act as acceptors of electrons.
CHARACTERISTIC PROPERTIES OF ELECTROVALENT (OR IONIC) COMPOUNDS)
1) Ionic compounds do not contain molecules but instead they consist of aggregates of oppositely charged ions.
2) Ionic compounds are solids and do not vaporize easily.
3) Ionic compounds are good conductors of heat and electricity; thus ionic compounds are electrolytes which can conduct electricity when molten or when dissolved in water.
4) Ionic compounds generally dissolve in water and other polar solvents like ethanol but they do not dissolve in non-polar solvents like benzene and tetra chloromethane salts, alkalis and bases are electrovalent and acids when in solution in water also show electrovalency.
EVALUATION:
(1) Explain the staple duplet and octet configurations of the noble gases using their electronic configurations.
(2) Relate chemical bonding to the tendency of other elements to attain the stable configuration of the noble gases.
(3) List the properties of electrovalent bonding and how they are formed.
(4) List the other types of chemical bonding.
ASSIGNMENT:
1) Use ionic equations and electronic structures to explain how the bonding takes place in potassium oxide K2O.
CONTENT: ORDINARY COVALENT COMBINATION
In ordinary covalent bonding, there is no transfer of electrons between the atoms. Instead, there is a sharing of a pair of electrons between the two reacting atoms so that both can attain the stable octet structure. This pair of shared electrons is called the SHARED PAIR where each member (reacting atom) contributes an electron each. The unshared pair of electrons in the outermost shell are called the LONE-ELECTRON PAIRS. Diatomic molecules are formed by covalent bonding (e.g. hydrogen molecule, H2 and chlorine molecule Cl2). Water molecule, H2O and hydrogen chloride gas molecule, HCl(g) are formed covalently.
Organic compounds are also formed by this method, e.g. the formation of methane.
Conventionally, the shared pair of electrons is represented by a stroke known as single covalent bond between two atoms in the bonding, e.g. H-H and Cl-Cl.
EXAMPLES OF COVALENT BONDINGS
1. Formation of a hydrogen molecule, H2.
2. Formation of a chlorine molecule, Cl2
3. Formation of a water molecule, H2O
4. Formation of methane molecule, CH4
5. Formation of ammonia molecule, NH3
EVALUATION:
a. Explain ordinary covalent bonding
b. Explain shared pair of electrons and lone electron pair
c. Use electronic structure diagrams to show the covalent combination in H2, Cl2 and H2O
ASSIGNMENT:
Use electronic structure diagrams to show how two ammonia molecules are formed from three hydrogen molecules and two nitrogen molecules.
CONTENT: ORDINARY COVALENT BONDING
Molecules with double bonds.
1. Formation of an oxygen molecule
2. Formation of carbon (iv) oxide molecule
3. Ethene molecule
Molecules with triple covalent bonds
1. Formation of nitrogen molecule
2. Ethyne molecule
EVALUATION:
(1) Explain double bond and triple bond formations
(2) Use electronic structure diagrams to illustrate double bond in O2 and CO2 molecules.
(3) Use electronic structure diagrams to illustrate triple bond in N2 and ethyne molecules.
ASSIGNMENT:
(1) Distinguish between the bonds present in calcium oxide and carbon (iv) oxide.
CONTENT: CHARACTERISTIC PROPERTIES OF COVALENT COMPUNDS
1. Covalent compounds consist of molecules. They contain ions.
2. Simple covalent compounds are gases or volatile liquids e.g. ammonia, carbon (iv) oxide, ethanol. This is so because their molecules are electrically neutral and have little attractive force for each other.
3. Since covalent compounds do not contain ions, thus, covalent compounds cannot conduct electricity and so are non-electrolytes.
4. Covalent compounds are usually soluble in non polar solvents (i.e. covalent organic solvents such as benzene or carbon disulphide) and not in polar solvents like water.
EVALUATION:
(1) List the characteristic properties of ordinary covalent molecules in terms of structural existence, state of matter, solubility and conductivity.
(2) Explain coordinate covalency using ammonium ion as a typical example.
ASSIGNMENT:
1) Distinguish between the characteristic properties of covalent compounds and electrovalent compounds.
2) Use electronic structure diagram to show the coordinate covalency combination in hydroxonium ion, H3O+.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a) Define and explain chemical bonding
b) List the types of chemical bonding
c) Explain ionic bonding with its properties
d) Explain covalent bonding with its properties
e) Explain coordinate covalency (dative bonding and its properties.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: CHEMICAL BONDING
The rare gases have very stable electronic configuration which makes them to be thermodynamically and chemically stable. The tendency of the other elements is to try to attain this stable duplet (helium) or octet (neon, argon, krypton) structure possessed by the rare (noble gases. This is achieved during chemical combination. There are two main types of chemical combination (bonding).
1. Electrovalent combination (or ionic bonding)
2. Covalent combination which is classified into:
a) Ordinary covalent combination
b) Coordinate covalent combination
ELECTROVALENT BONDING (IONIC BONDING)
In electrovalent bonding, there is a transfer of outermost shell (valence) electrons from one atom (usually metallic) to another atom (usually non-metallic). Thus, we have atoms which act as donors of electrons and those which act as acceptors of electrons.
CHARACTERISTIC PROPERTIES OF ELECTROVALENT (OR IONIC) COMPOUNDS)
1) Ionic compounds do not contain molecules but instead they consist of aggregates of oppositely charged ions.
2) Ionic compounds are solids and do not vaporize easily.
3) Ionic compounds are good conductors of heat and electricity; thus ionic compounds are electrolytes which can conduct electricity when molten or when dissolved in water.
4) Ionic compounds generally dissolve in water and other polar solvents like ethanol but they do not dissolve in non-polar solvents like benzene and tetra chloromethane salts, alkalis and bases are electrovalent and acids when in solution in water also show electrovalency.
EVALUATION:
(1) Explain the staple duplet and octet configurations of the noble gases using their electronic configurations.
(2) Relate chemical bonding to the tendency of other elements to attain the stable configuration of the noble gases.
(3) List the properties of electrovalent bonding and how they are formed.
(4) List the other types of chemical bonding.
ASSIGNMENT:
1) Use ionic equations and electronic structures to explain how the bonding takes place in potassium oxide K2O.
CONTENT: ORDINARY COVALENT COMBINATION
In ordinary covalent bonding, there is no transfer of electrons between the atoms. Instead, there is a sharing of a pair of electrons between the two reacting atoms so that both can attain the stable octet structure. This pair of shared electrons is called the SHARED PAIR where each member (reacting atom) contributes an electron each. The unshared pair of electrons in the outermost shell are called the LONE-ELECTRON PAIRS. Diatomic molecules are formed by covalent bonding (e.g. hydrogen molecule, H2 and chlorine molecule Cl2). Water molecule, H2O and hydrogen chloride gas molecule, HCl(g) are formed covalently.
Organic compounds are also formed by this method, e.g. the formation of methane.
Conventionally, the shared pair of electrons is represented by a stroke known as single covalent bond between two atoms in the bonding, e.g. H-H and Cl-Cl.
EXAMPLES OF COVALENT BONDINGS
1. Formation of a hydrogen molecule, H2.
2. Formation of a chlorine molecule, Cl2
3. Formation of a water molecule, H2O
4. Formation of methane molecule, CH4
5. Formation of ammonia molecule, NH3
EVALUATION:
a. Explain ordinary covalent bonding
b. Explain shared pair of electrons and lone electron pair
c. Use electronic structure diagrams to show the covalent combination in H2, Cl2 and H2O
ASSIGNMENT:
Use electronic structure diagrams to show how two ammonia molecules are formed from three hydrogen molecules and two nitrogen molecules.
CONTENT: ORDINARY COVALENT BONDING
Molecules with double bonds.
1. Formation of an oxygen molecule
2. Formation of carbon (iv) oxide molecule
3. Ethene molecule
Molecules with triple covalent bonds
1. Formation of nitrogen molecule
2. Ethyne molecule
EVALUATION:
(1) Explain double bond and triple bond formations
(2) Use electronic structure diagrams to illustrate double bond in O2 and CO2 molecules.
(3) Use electronic structure diagrams to illustrate triple bond in N2 and ethyne molecules.
ASSIGNMENT:
(1) Distinguish between the bonds present in calcium oxide and carbon (iv) oxide.
CONTENT: CHARACTERISTIC PROPERTIES OF COVALENT COMPUNDS
1. Covalent compounds consist of molecules. They contain ions.
2. Simple covalent compounds are gases or volatile liquids e.g. ammonia, carbon (iv) oxide, ethanol. This is so because their molecules are electrically neutral and have little attractive force for each other.
3. Since covalent compounds do not contain ions, thus, covalent compounds cannot conduct electricity and so are non-electrolytes.
4. Covalent compounds are usually soluble in non polar solvents (i.e. covalent organic solvents such as benzene or carbon disulphide) and not in polar solvents like water.
EVALUATION:
(1) List the characteristic properties of ordinary covalent molecules in terms of structural existence, state of matter, solubility and conductivity.
(2) Explain coordinate covalency using ammonium ion as a typical example.
ASSIGNMENT:
1) Distinguish between the characteristic properties of covalent compounds and electrovalent compounds.
2) Use electronic structure diagram to show the coordinate covalency combination in hydroxonium ion, H3O+.
WEEK 8
TOPIC: SIMPLE MOLECULES AND THEIR SHAPES/METALLIC BONDING
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Explain metallic bonding, properties of metals, Van der waals forces.
b. Explain hydrogen bond with illustrated examples.
c. Relate metallic bonding, Van der Waals forces and hydrogen bond as inter-molecular bonding.
d. Draw and explain the shapes of simple covalent molecules such as methane, ammonia, water, carbon (iv) oxide and also explain the term hybridization.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: METALLIC BONDING
Metal atoms are held together in solid crystalline form by metallic bonding in something like the following way. The valence electrons of each metallic atom are only loosely held, being relatively distant from the nucleus, and they separate from particular nuclei to move at random through the crystal lattice. The residual ions, now positively charged by loss of valence electrons, tend to repel each other but are held together by the moving electron cloud ad some overlapping of residual electron orbits.
This type of bonding is very strong in some metals, e.g. iron, which are difficult to shatter, but it is much weaker in, e.g. sodium or potassium, which can be cut with a knife.
PHYSICAL PROPERTIES OF METALS AS EXPLAINED BY METALLIC BONDING
The physical properties of metals are:
1. High melting and boiling points.
2. Characteristic luster.
3. Malleable, i.e. can be beaten into shapes
4. Ductile, i.e. can be drawn into a thin wire
5. Sonorous, i.e. it gives off a note when hit
6. Hard but not brittle, with great tensile strength
7. Relatively high densities
8. Good conductors of heat and electricity
Some metals however, do not show all these properties, e.g. mercury is a liquid with a melting point of -39℃ while sodium and potassium are light, soft metals with low melting points of 97℃ and 63℃ respectively.
If an electric potential difference is applied to the ends of a metallic rod, the free electrons lose their random motion and move towards the positive end of the rod, being steadily replaced by more from the source of potential difference. Hence metals are good electrical conductors. As freely moving electrons can convey heat energy, metals are also good conductors of heat.
Since the bonding agent in a metal is mainly a moving electron cloud, the ions of most metals will usually slide relative to one another, under stress, without shattering the lattice and produce a new position of stability. This accounts for the malleability and ductility of many metals, though special temperature conditions may sometimes be necessary.
EVALUATION: The students should be able to:
(1) Explain metallic bonding with the aid of a diagram
(2) List the physical properties (8 of them) of metals and use metallic bonding to explain them.
ASSIGNMENT:
Give an explanation to account for why mercury is a liquid.
PERIOD: 2
CONTENT: CHEMICAL PROPERTIES OF METALS
Chemical properties of elements are dependent on the number of valence electrons present in their atomic structures.
A METAL is an element whose atoms ionize by electron loss to form positively charged ions (i.e CAT IONS).
The chemical properties of metals are:
1. Ionization behaviour: Metals ionize by losing electrons (i.e. they are electropositive) to become positive ions which enter into chemical combinations with negative ions or radicals (mainly composed of non-metals) to form electrovalent compounds.
2. Reducing agents: Metals are reducing agents by definition because they tend to donate their valence electrons readily during chemical reactions.
3. Reaction with acids: A metal which is more electropositive than hydrogen readily displaces the hydrogen ion, H+, from an acid.
4. Nature of the oxides: Most metals react with oxygen to form basic oxides which are mainly ionic compounds, soluble basic oxides form alkalis.
Some metals like aluminum and zinc, however, form amphoteric oxides (i.e. having both basic and acidic characteristics).
EVALUATION: The students should be able to:
1. Explain the basis of the chemical properties of elements.
2. Define a metal in terms of its ionization.
3. List the four chemical properties of metals and briefly explain them.
ASSIGNMENT:
1. Use chemical equation to show how magnesium forms a basic oxide and thus an alkali.
PERIOD: 3
CONTENT: VAN DER WAALS FORCES AND HYDROGEN BOND.
VAN DER WAALS FORCES
Most non-metallic elements form molecules by covalent combination, e.g. iodine as I2. In appropriate conditions of temperature and pressure, the atomic nuclei of one molecule and the electrons of another molecule attract each other sufficiently to bring about a close approach. As the molecules come together, the electrons of each begin to exert repulsive forces on each other. The forces of attraction and repulsion are balanced in the formation of a crystal. These are called Van der Waals forces which are rather weak and the crystals tend to have low melting points. Simple molecules of compounds e.g. naphthalene form crystalline solids by the operation of Van der waals forces in a similar way.
HYDROGEN BOND
The hydrogen bond is an intermolecular force which arises when hydrogen is covalently linked to elements like nitrogen, oxygen and fluorine which are strongly electronegative (i.e. they have affinities for electrons). They tend to pull the shared pair of electrons in the covalent bonds toward themselves, resulting in the formation of a DIPOLE where the hydrogen atom is partially positive, while the nitrogen, oxygen or fluorine atom is partially negative. An electrostatic attraction occurs between the two dipoles when the positive pole of one attracts the negative pole of another. This attractive force is known as the HYDROGEN BOND. Although hydrogen bond is weak, it has important effects on the physical properties of compounds like hydrogen fluoride and water.
EVALUATION: The students should be able to:
1. Explain Van der Waals forces
2. List two substances each held together by Van der Waals forces.
3. Explain hydrogen bond.
4. List two substances each in which hydrogen bond is operating
ASSIGNMENT:
Explain briefly in electronic terms why the ammonia molecule can participate readily in co-ordinate covalency.
PERIOD: 4
CONTENT: SHAPES OF SIMPLE COVALENT MOLECULES
Simple covalent molecules are normally represented as flat structures. But in reality, such molecules have definite shape. This is because atoms of molecules combine by sharing electrons and the single covalent bond formed is as a result of an overlap of two atomic orbitals, each occupied by a single electron. Thus, the bond will be orientated in the direction of the orbitals which provide the electrons, so the bond is directional. Such bonds influence the shape and structure of the resulting molecule. Electron pairs include shared (or bonding) pairs and lone (non-bonding) pairs. The electron clouds of these pairs are negatively charged. Within a molecule, such electron clouds mutually repel each other and stay as far apart as possible. This affects the shape of the molecule.
METHANE, CH4
For carbon to attain a valency of four, one electron in the 2S is “promoted” to occupy the empty 2IPz orbital, creating four unpaired electrons (one S and three P) in its valance shell.
HYBRIDIZATION is the promotion of an electron from one orbital to another of the same energy level so as to have a mixture of orbitals (or hybrid orbitals e.g. SP3) for bonding.
TETRAHEDRAL SHAPE OF CH4
a. Diagrammatic representation
b. Three-dimensional representation
TRIGONAL PYRAMIDAL SHAPE OF THE AMMONIA MOLECULE
a. Diagrammatical representation
b. Three-dimensional representation
THE ANGULAR SHAPE OF WATER MOLECULE
a. Diagrammatical representation
b. Three-dimensional representation
THE LINEAR SHAPE OF THE CARBON (IV) OXIDE MOLECULE
a. Diagrammatic representation
b. Three-dimensional representation
THE LINEAR SHAPES OF OXYGEN, HYDROGEN AND CHLORINE MOLECULES
a. Diagrammatical representation
b. Three-dimensional representations
EVALUATION: The students should be able to:
a. Briefly explain the definite shape of simple covalent molecules.
b. Use methane structure to explain hybridization
c. Draw the tetrahedral shape of CH4 three-dimensionally
d. Draw the trigonal pyramidical shape and the angular shape of ammonia and water molecules respectively in a three-dimensional representation.
e. Draw the linear shapes of CO2, O2, H2 and Cl2
ASSIGNMENT:
Draw the electronic orbital box structure of hybridized carbon showing the spin of the electrons.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Explain metallic bonding, properties of metals, Van der waals forces.
b. Explain hydrogen bond with illustrated examples.
c. Relate metallic bonding, Van der Waals forces and hydrogen bond as inter-molecular bonding.
d. Draw and explain the shapes of simple covalent molecules such as methane, ammonia, water, carbon (iv) oxide and also explain the term hybridization.
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: METALLIC BONDING
Metal atoms are held together in solid crystalline form by metallic bonding in something like the following way. The valence electrons of each metallic atom are only loosely held, being relatively distant from the nucleus, and they separate from particular nuclei to move at random through the crystal lattice. The residual ions, now positively charged by loss of valence electrons, tend to repel each other but are held together by the moving electron cloud ad some overlapping of residual electron orbits.
This type of bonding is very strong in some metals, e.g. iron, which are difficult to shatter, but it is much weaker in, e.g. sodium or potassium, which can be cut with a knife.
PHYSICAL PROPERTIES OF METALS AS EXPLAINED BY METALLIC BONDING
The physical properties of metals are:
1. High melting and boiling points.
2. Characteristic luster.
3. Malleable, i.e. can be beaten into shapes
4. Ductile, i.e. can be drawn into a thin wire
5. Sonorous, i.e. it gives off a note when hit
6. Hard but not brittle, with great tensile strength
7. Relatively high densities
8. Good conductors of heat and electricity
Some metals however, do not show all these properties, e.g. mercury is a liquid with a melting point of -39℃ while sodium and potassium are light, soft metals with low melting points of 97℃ and 63℃ respectively.
If an electric potential difference is applied to the ends of a metallic rod, the free electrons lose their random motion and move towards the positive end of the rod, being steadily replaced by more from the source of potential difference. Hence metals are good electrical conductors. As freely moving electrons can convey heat energy, metals are also good conductors of heat.
Since the bonding agent in a metal is mainly a moving electron cloud, the ions of most metals will usually slide relative to one another, under stress, without shattering the lattice and produce a new position of stability. This accounts for the malleability and ductility of many metals, though special temperature conditions may sometimes be necessary.
EVALUATION: The students should be able to:
(1) Explain metallic bonding with the aid of a diagram
(2) List the physical properties (8 of them) of metals and use metallic bonding to explain them.
ASSIGNMENT:
Give an explanation to account for why mercury is a liquid.
PERIOD: 2
CONTENT: CHEMICAL PROPERTIES OF METALS
Chemical properties of elements are dependent on the number of valence electrons present in their atomic structures.
A METAL is an element whose atoms ionize by electron loss to form positively charged ions (i.e CAT IONS).
The chemical properties of metals are:
1. Ionization behaviour: Metals ionize by losing electrons (i.e. they are electropositive) to become positive ions which enter into chemical combinations with negative ions or radicals (mainly composed of non-metals) to form electrovalent compounds.
2. Reducing agents: Metals are reducing agents by definition because they tend to donate their valence electrons readily during chemical reactions.
3. Reaction with acids: A metal which is more electropositive than hydrogen readily displaces the hydrogen ion, H+, from an acid.
4. Nature of the oxides: Most metals react with oxygen to form basic oxides which are mainly ionic compounds, soluble basic oxides form alkalis.
Some metals like aluminum and zinc, however, form amphoteric oxides (i.e. having both basic and acidic characteristics).
EVALUATION: The students should be able to:
1. Explain the basis of the chemical properties of elements.
2. Define a metal in terms of its ionization.
3. List the four chemical properties of metals and briefly explain them.
ASSIGNMENT:
1. Use chemical equation to show how magnesium forms a basic oxide and thus an alkali.
PERIOD: 3
CONTENT: VAN DER WAALS FORCES AND HYDROGEN BOND.
VAN DER WAALS FORCES
Most non-metallic elements form molecules by covalent combination, e.g. iodine as I2. In appropriate conditions of temperature and pressure, the atomic nuclei of one molecule and the electrons of another molecule attract each other sufficiently to bring about a close approach. As the molecules come together, the electrons of each begin to exert repulsive forces on each other. The forces of attraction and repulsion are balanced in the formation of a crystal. These are called Van der Waals forces which are rather weak and the crystals tend to have low melting points. Simple molecules of compounds e.g. naphthalene form crystalline solids by the operation of Van der waals forces in a similar way.
HYDROGEN BOND
The hydrogen bond is an intermolecular force which arises when hydrogen is covalently linked to elements like nitrogen, oxygen and fluorine which are strongly electronegative (i.e. they have affinities for electrons). They tend to pull the shared pair of electrons in the covalent bonds toward themselves, resulting in the formation of a DIPOLE where the hydrogen atom is partially positive, while the nitrogen, oxygen or fluorine atom is partially negative. An electrostatic attraction occurs between the two dipoles when the positive pole of one attracts the negative pole of another. This attractive force is known as the HYDROGEN BOND. Although hydrogen bond is weak, it has important effects on the physical properties of compounds like hydrogen fluoride and water.
EVALUATION: The students should be able to:
1. Explain Van der Waals forces
2. List two substances each held together by Van der Waals forces.
3. Explain hydrogen bond.
4. List two substances each in which hydrogen bond is operating
ASSIGNMENT:
Explain briefly in electronic terms why the ammonia molecule can participate readily in co-ordinate covalency.
PERIOD: 4
CONTENT: SHAPES OF SIMPLE COVALENT MOLECULES
Simple covalent molecules are normally represented as flat structures. But in reality, such molecules have definite shape. This is because atoms of molecules combine by sharing electrons and the single covalent bond formed is as a result of an overlap of two atomic orbitals, each occupied by a single electron. Thus, the bond will be orientated in the direction of the orbitals which provide the electrons, so the bond is directional. Such bonds influence the shape and structure of the resulting molecule. Electron pairs include shared (or bonding) pairs and lone (non-bonding) pairs. The electron clouds of these pairs are negatively charged. Within a molecule, such electron clouds mutually repel each other and stay as far apart as possible. This affects the shape of the molecule.
METHANE, CH4
For carbon to attain a valency of four, one electron in the 2S is “promoted” to occupy the empty 2IPz orbital, creating four unpaired electrons (one S and three P) in its valance shell.
HYBRIDIZATION is the promotion of an electron from one orbital to another of the same energy level so as to have a mixture of orbitals (or hybrid orbitals e.g. SP3) for bonding.
TETRAHEDRAL SHAPE OF CH4
a. Diagrammatic representation
b. Three-dimensional representation
TRIGONAL PYRAMIDAL SHAPE OF THE AMMONIA MOLECULE
a. Diagrammatical representation
b. Three-dimensional representation
THE ANGULAR SHAPE OF WATER MOLECULE
a. Diagrammatical representation
b. Three-dimensional representation
THE LINEAR SHAPE OF THE CARBON (IV) OXIDE MOLECULE
a. Diagrammatic representation
b. Three-dimensional representation
THE LINEAR SHAPES OF OXYGEN, HYDROGEN AND CHLORINE MOLECULES
a. Diagrammatical representation
b. Three-dimensional representations
EVALUATION: The students should be able to:
a. Briefly explain the definite shape of simple covalent molecules.
b. Use methane structure to explain hybridization
c. Draw the tetrahedral shape of CH4 three-dimensionally
d. Draw the trigonal pyramidical shape and the angular shape of ammonia and water molecules respectively in a three-dimensional representation.
e. Draw the linear shapes of CO2, O2, H2 and Cl2
ASSIGNMENT:
Draw the electronic orbital box structure of hybridized carbon showing the spin of the electrons.
WEEK 9
TOPIC: STOICHIOMETRY AND CHEMICAL REACTIONS
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Express the relationship between chemical quantities such as mole ratios and mass relationships.
b. Balance any given chemical equation stoichiometrically.
c. Determine the empirical formula and the molecular formula of chemical compounds.
d. Express the laws of chemical combination
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: STOICHIOMETRY OF REACTIONS
Stoichiometry of reaction is the mole ratio in which reactants combine and products are formed. Since a mole represents the molar mass of a substance in grams, we can use this to calculate the reacting mass of reactants and products formed from a balanced equation,
A MOLE is the amount of substance which contains as many elementary entities (atoms, molecules, ions) as there are carbon atoms in 12g of carbon -12.
CALCULATIONS
1. A metal X with relative atomic mass 56 forms an oxide with formula X2O3. How many grams of the metal will combine with 10g of oxygen?
SOLUTION:
4 X +3O2 2X2O3
4 : 3 : 2
(4 x 56g) (3 X 16 X 2)
224g 96g
Therefore: 96g of oxygen combined with 224g of X
Therefore: 10g of oxygen will combine with 224g X 10g
96g
= 23.33g of X
Therefore: 10g of oxygen will combine with 23.33g of X.
2. 12.5g of zinc trioxocarbonate (iv), ZnCO3, were heated very strongly to a constant mass and the residue treated with excess hydrochloric acid, HCl. Calculate the mass of Zinc chloride, ZnCl2, that would be obtained. (Zn= 65, C = 12, O = 16, H = 1, Cl = 35.5)
SOLUTION:
ZnCO3(s) ZnO(s) + CO2(g)
(65 + 12 + 48)g (65 + 16)g
= 125g 81g
ZnO (s) + 2HCl(aq) ZnCl2(aq) + H2O(l)
81g (65 + 71)g
81g 136g
ZnCO3(s) ZnO(s) + CO2(g)
125g 81g
125g of ZnZO3 produced 81g of ZnO
Therefore: 12.5g of ZnCO3 will produce 81g X 12.5g of ZnO
125g
ZnO(s) + 2HCl(aq) ZnCl2(aq) + H2O(l)
81g 136g
81g of ZnO produced 136g of ZnCl2
Therefore: 8.1g of ZnO will produce 136g X 8.1g of ZnCl2
81g
= 13.6g of ZnCl2
Therefore: 13.6g of ZnCl2 would be obtained.
EVALUATION:
1. Calculate the mass concentration in gdm-3 of the following solutions:
a. 1moldm-3 potassium hydroxide solution
b. 0.1moldm-3 trioxonitrate (V) acid
c. 0.05moldm-3 solution of sodium trioxocarbonate (IV)- decahydrate (K = 39, H=1, O= 16, N=14, Na=23, C=12)
ASSIGNMENT:
Explain the significance of stoichiometry in calculations involving chemical reactions.
CONTENT: CALCULATIONS INVOLVING MASS-VOLUME IN CHEMICAL REACTIONS
1. Find the volume of oxygen produced by 1mole of potassium trioxochlorate (V) at s.t.p in the following reaction. (K=39, Cl=35.5, O=16. Molar volume at s.t.p = 22.4dm3)
2KClO3(s) 2KCl(s) + 3O2(g)
SOLUTION:
2KClO3(s) 2KCl(s) + 3O2(g)
2moles 3moles
2 (39 + 35.5 + 48)gmol-1 3(22.4dm3)
2(122.5)gmol-1
From the equation,
2moles of KClO3 produced 3 moles of O2
Therefore: 2moles of KClO3 produced 3 X 22.4dm3 of O2
Therefore 1 mole of KClO3 produced 3 X 22.4dm3 X 1 of O2
2
= 67.2dm3 of O2 = 33.6dm3
2
Therefore : 33.6dm3 of O2 will be produced.
2. Ethane burns completely in oxygen according to the equation below
C2H6(g) + 7 O2(g) 2CO2(g) + 3H2O(l)
2
What is the amount in moles of carbon (IV) oxide that will be produced when 6.0g of ethane are completely burnt in oxygen? (O=16, C=12, H=1)
SOLUTION:
C2H6(g) + 7 O2(g) 2CO2(g) + 3H2O(l)
2
2C2H6(g) + 7 O2(g) 4CO2(g) + 3H2O(l)
2moles 4moles
Number of moles = 6g = 1 mol
30gmol-1 5
2 moles of C2H6 produced 4 moles of CO2
1/5 mol of C2H6 will produce 4 moles X 1 mol
2 moles 3
= 0.4 moles of O2
Therefore: 0.4 moles of CO2 will be produced.
EVALUATION:
a. Find the mass of sodium trioxocarbonate (IV) needed to give 22.4dm3 of carbon (IV) oxide at s.t.p in the reaction.
Na2CO3(s) + 2HCl(aq) 2NaCl(AQ) + CO2(g) + H2O(l)
(Na=23, O=16, C=12, molar volume of a gas at s.t.p =22.4dm3)
b. A excess of a divalent metal M was dissolved in a limited volume of hydrochloric acid. If 576cm3 of hydrogen were liberated at s.t.p. what was the mass of the metal that produced this volume of hydrogen? (M=24, H=1, molar volume of a gas at s.t.p = 22.4dm3).
ASSIGNMENT:
Express the formula for the relationship between number of moles, reacting mass and molar mass.
CONTENT: EMPIRICAL FORMULA AND MOLECULAR FORMULA
EMPIRICAL FORMULA of a compound is its simplest formula which tells about the component elements in a molecule of the compound and the ratio in which these elements are combined together, i.e. the mole ratio. Mole ratio of each component element in a compound
= % by mass of the element in the compound
Relative atomic mass of the element
OR
= Reacting mass of the element in the compound
Molar mass of the element
The MOLECULAR FORMULA of a compound gives the exact number of moles of atoms of the component elements in one mole of the compound.
In most cases, the empirical and molecular formulae of a compound are the same. In some cases, the molecular formula is a simple multiple of the empirical formula which can be calculated if the empirical formula and the formula mass/the relative molecular mass are known.
1. Find the empirical formula of the compound given from the percentage composition by mass.
Ag = 63.53%; N = 8.23%; O = 28.24% (Ag=108, N=14, O=16)
SOLUTION
Ag N O
% by mass 63.53% 8.23% 28.24%
Relative atomic mass 108 14 16
Mole ratio = 63.53% 8.23% 28.24%
108 14 16
= 0.59 0.59 1.765
Dividing by the smallest,
0.59 =1 0.59 = 1 1.765 = 3
0.59 0.59 0.59
Ag = 1; N=1; O=3
Empirical formula = AgNO3
2. In a certain reaction. 0.20g of hydrogen gas combines with 3.20g of oxygen gas. If the relative molecular mass of the compound is 34.0, find its molecular formula. (H=1, O=16).
SOLUTION
H O
Reacting mass 0.20g 3.20g
Relative atomic mass 1 16
Mole ratio 0.20 3.20
1 16
= 0.2 0.2
Dividing by the smallest,
0.2 =1 0.2 = 1
0.2 0.2
Empirical formula = HO
But relative molecular mass = 34.0
Therefore (HO)n = 34.0
(I + 16)n = 34.0
17n = 34.0
n = 34.0 = 2
17
n = 2
(HO)2 = H2O2
Therefore: molecular formula of the compound = H2O2
EVALUATION:
1. Define empirical formula and establish the formula for calculating it.
2. Define molecular formula and establish how it is mathematically related to empirical formula.
ASSIGNMENT:
Find the empirical formula of the compound below from the percentage composition by mass
N= 26.17%; H= 7.48%; Cl = 66.35%
CONTENT: LAWS OF CHEMICAL COMBINATION
The following are the laws of chemical combination:
1. The law of conservation of mass
2. The law of definite proportion
3. The law of multiple proportion
1. THE LAW OF CONSERVATION OF MASS states that matter can neither be created nor destroyed during a chemical reaction but changes from one form to another.
2. THE LAW OF DEFINITE PROPORTION states that all pure samples of a particular chemical compound contains the same elements combined in the same proportion by mass.
3. THE LAW OF MULTIPLE PROPORTIONS states that if two elements, A and B, combine to form more than one chemical compound, then the various masses of A which separately combine with a fixed mass of B are in a simple multiple ration, e.g.
a. Black copper (II) oxide, CuO, and red copper (I) oxide, Cu2O.
b. Brown iron (III) oxide, Fe2O3, and black iron (II) oxide, FeO.
c. Brownish-yellow iron (III) chloride, and green iron (II) chloride, FeCl2
d. Black lead (II) sulphide, PbS, and lead (IV) sulphide, PbS2
CALCULATION
1) A metal X forms two different chlorides. If 12.7g of chloride A and 16.3g of chloride B contain 7.1g and 10.7g of chlorine respectively, show that the figures agree with the Law of Multiple proportions. Write their formulae.
SOLUTION
A B
X: (12.7 - 7.1)g (16.3 - 10.7)g
= 5.6g 5.6g
Cl: 7.1g 10.7g
From the data above, a fixed mass of metal X combined separately with different masses of chlorine and this is in accordance with the law of multiple proportions.
A B
X 5.6g 5.6g
Cl 7.1g 10.7g
5.6g of X combine with 7.1g of Cl in a
1g of X combine with 7.1g X 1g = 1.3g of Cl
5.6g
5.6g of X combine with 10.7g of Cl in B
1g of X will combine with 10.7g X 1g = 2.0g of Cl
5.6g
A has formula of XCl and B has formula of XCl2
EVALUATION:
1. State the law of conservation of matter
2. State the law of definite proportions
3. State the law of multiple proportions
ASSIGNMENT:
Give two examples of compounds of paired elements that conform with the law of multiple proportions.
BEHAVIOURAL OBJECTIVES: BY THE END OF THE LESSON, THE STUDENTS SHOULD BE ABLE TO:
a. Express the relationship between chemical quantities such as mole ratios and mass relationships.
b. Balance any given chemical equation stoichiometrically.
c. Determine the empirical formula and the molecular formula of chemical compounds.
d. Express the laws of chemical combination
REFERENCE: NEW CERTIFICATE CHEMISTRY FOR SSCE by OSEI YAW ABABIO (NEW EDITION)
CONTENT: STOICHIOMETRY OF REACTIONS
Stoichiometry of reaction is the mole ratio in which reactants combine and products are formed. Since a mole represents the molar mass of a substance in grams, we can use this to calculate the reacting mass of reactants and products formed from a balanced equation,
A MOLE is the amount of substance which contains as many elementary entities (atoms, molecules, ions) as there are carbon atoms in 12g of carbon -12.
CALCULATIONS
1. A metal X with relative atomic mass 56 forms an oxide with formula X2O3. How many grams of the metal will combine with 10g of oxygen?
SOLUTION:
4 X +3O2 2X2O3
4 : 3 : 2
(4 x 56g) (3 X 16 X 2)
224g 96g
Therefore: 96g of oxygen combined with 224g of X
Therefore: 10g of oxygen will combine with 224g X 10g
96g
= 23.33g of X
Therefore: 10g of oxygen will combine with 23.33g of X.
2. 12.5g of zinc trioxocarbonate (iv), ZnCO3, were heated very strongly to a constant mass and the residue treated with excess hydrochloric acid, HCl. Calculate the mass of Zinc chloride, ZnCl2, that would be obtained. (Zn= 65, C = 12, O = 16, H = 1, Cl = 35.5)
SOLUTION:
ZnCO3(s) ZnO(s) + CO2(g)
(65 + 12 + 48)g (65 + 16)g
= 125g 81g
ZnO (s) + 2HCl(aq) ZnCl2(aq) + H2O(l)
81g (65 + 71)g
81g 136g
ZnCO3(s) ZnO(s) + CO2(g)
125g 81g
125g of ZnZO3 produced 81g of ZnO
Therefore: 12.5g of ZnCO3 will produce 81g X 12.5g of ZnO
125g
ZnO(s) + 2HCl(aq) ZnCl2(aq) + H2O(l)
81g 136g
81g of ZnO produced 136g of ZnCl2
Therefore: 8.1g of ZnO will produce 136g X 8.1g of ZnCl2
81g
= 13.6g of ZnCl2
Therefore: 13.6g of ZnCl2 would be obtained.
EVALUATION:
1. Calculate the mass concentration in gdm-3 of the following solutions:
a. 1moldm-3 potassium hydroxide solution
b. 0.1moldm-3 trioxonitrate (V) acid
c. 0.05moldm-3 solution of sodium trioxocarbonate (IV)- decahydrate (K = 39, H=1, O= 16, N=14, Na=23, C=12)
ASSIGNMENT:
Explain the significance of stoichiometry in calculations involving chemical reactions.
CONTENT: CALCULATIONS INVOLVING MASS-VOLUME IN CHEMICAL REACTIONS
1. Find the volume of oxygen produced by 1mole of potassium trioxochlorate (V) at s.t.p in the following reaction. (K=39, Cl=35.5, O=16. Molar volume at s.t.p = 22.4dm3)
2KClO3(s) 2KCl(s) + 3O2(g)
SOLUTION:
2KClO3(s) 2KCl(s) + 3O2(g)
2moles 3moles
2 (39 + 35.5 + 48)gmol-1 3(22.4dm3)
2(122.5)gmol-1
From the equation,
2moles of KClO3 produced 3 moles of O2
Therefore: 2moles of KClO3 produced 3 X 22.4dm3 of O2
Therefore 1 mole of KClO3 produced 3 X 22.4dm3 X 1 of O2
2
= 67.2dm3 of O2 = 33.6dm3
2
Therefore : 33.6dm3 of O2 will be produced.
2. Ethane burns completely in oxygen according to the equation below
C2H6(g) + 7 O2(g) 2CO2(g) + 3H2O(l)
2
What is the amount in moles of carbon (IV) oxide that will be produced when 6.0g of ethane are completely burnt in oxygen? (O=16, C=12, H=1)
SOLUTION:
C2H6(g) + 7 O2(g) 2CO2(g) + 3H2O(l)
2
2C2H6(g) + 7 O2(g) 4CO2(g) + 3H2O(l)
2moles 4moles
Number of moles = 6g = 1 mol
30gmol-1 5
2 moles of C2H6 produced 4 moles of CO2
1/5 mol of C2H6 will produce 4 moles X 1 mol
2 moles 3
= 0.4 moles of O2
Therefore: 0.4 moles of CO2 will be produced.
EVALUATION:
a. Find the mass of sodium trioxocarbonate (IV) needed to give 22.4dm3 of carbon (IV) oxide at s.t.p in the reaction.
Na2CO3(s) + 2HCl(aq) 2NaCl(AQ) + CO2(g) + H2O(l)
(Na=23, O=16, C=12, molar volume of a gas at s.t.p =22.4dm3)
b. A excess of a divalent metal M was dissolved in a limited volume of hydrochloric acid. If 576cm3 of hydrogen were liberated at s.t.p. what was the mass of the metal that produced this volume of hydrogen? (M=24, H=1, molar volume of a gas at s.t.p = 22.4dm3).
ASSIGNMENT:
Express the formula for the relationship between number of moles, reacting mass and molar mass.
CONTENT: EMPIRICAL FORMULA AND MOLECULAR FORMULA
EMPIRICAL FORMULA of a compound is its simplest formula which tells about the component elements in a molecule of the compound and the ratio in which these elements are combined together, i.e. the mole ratio. Mole ratio of each component element in a compound
= % by mass of the element in the compound
Relative atomic mass of the element
OR
= Reacting mass of the element in the compound
Molar mass of the element
The MOLECULAR FORMULA of a compound gives the exact number of moles of atoms of the component elements in one mole of the compound.
In most cases, the empirical and molecular formulae of a compound are the same. In some cases, the molecular formula is a simple multiple of the empirical formula which can be calculated if the empirical formula and the formula mass/the relative molecular mass are known.
1. Find the empirical formula of the compound given from the percentage composition by mass.
Ag = 63.53%; N = 8.23%; O = 28.24% (Ag=108, N=14, O=16)
SOLUTION
Ag N O
% by mass 63.53% 8.23% 28.24%
Relative atomic mass 108 14 16
Mole ratio = 63.53% 8.23% 28.24%
108 14 16
= 0.59 0.59 1.765
Dividing by the smallest,
0.59 =1 0.59 = 1 1.765 = 3
0.59 0.59 0.59
Ag = 1; N=1; O=3
Empirical formula = AgNO3
2. In a certain reaction. 0.20g of hydrogen gas combines with 3.20g of oxygen gas. If the relative molecular mass of the compound is 34.0, find its molecular formula. (H=1, O=16).
SOLUTION
H O
Reacting mass 0.20g 3.20g
Relative atomic mass 1 16
Mole ratio 0.20 3.20
1 16
= 0.2 0.2
Dividing by the smallest,
0.2 =1 0.2 = 1
0.2 0.2
Empirical formula = HO
But relative molecular mass = 34.0
Therefore (HO)n = 34.0
(I + 16)n = 34.0
17n = 34.0
n = 34.0 = 2
17
n = 2
(HO)2 = H2O2
Therefore: molecular formula of the compound = H2O2
EVALUATION:
1. Define empirical formula and establish the formula for calculating it.
2. Define molecular formula and establish how it is mathematically related to empirical formula.
ASSIGNMENT:
Find the empirical formula of the compound below from the percentage composition by mass
N= 26.17%; H= 7.48%; Cl = 66.35%
CONTENT: LAWS OF CHEMICAL COMBINATION
The following are the laws of chemical combination:
1. The law of conservation of mass
2. The law of definite proportion
3. The law of multiple proportion
1. THE LAW OF CONSERVATION OF MASS states that matter can neither be created nor destroyed during a chemical reaction but changes from one form to another.
2. THE LAW OF DEFINITE PROPORTION states that all pure samples of a particular chemical compound contains the same elements combined in the same proportion by mass.
3. THE LAW OF MULTIPLE PROPORTIONS states that if two elements, A and B, combine to form more than one chemical compound, then the various masses of A which separately combine with a fixed mass of B are in a simple multiple ration, e.g.
a. Black copper (II) oxide, CuO, and red copper (I) oxide, Cu2O.
b. Brown iron (III) oxide, Fe2O3, and black iron (II) oxide, FeO.
c. Brownish-yellow iron (III) chloride, and green iron (II) chloride, FeCl2
d. Black lead (II) sulphide, PbS, and lead (IV) sulphide, PbS2
CALCULATION
1) A metal X forms two different chlorides. If 12.7g of chloride A and 16.3g of chloride B contain 7.1g and 10.7g of chlorine respectively, show that the figures agree with the Law of Multiple proportions. Write their formulae.
SOLUTION
A B
X: (12.7 - 7.1)g (16.3 - 10.7)g
= 5.6g 5.6g
Cl: 7.1g 10.7g
From the data above, a fixed mass of metal X combined separately with different masses of chlorine and this is in accordance with the law of multiple proportions.
A B
X 5.6g 5.6g
Cl 7.1g 10.7g
5.6g of X combine with 7.1g of Cl in a
1g of X combine with 7.1g X 1g = 1.3g of Cl
5.6g
5.6g of X combine with 10.7g of Cl in B
1g of X will combine with 10.7g X 1g = 2.0g of Cl
5.6g
A has formula of XCl and B has formula of XCl2
EVALUATION:
1. State the law of conservation of matter
2. State the law of definite proportions
3. State the law of multiple proportions
ASSIGNMENT:
Give two examples of compounds of paired elements that conform with the law of multiple proportions.
