Grade 12 Chemistry Notes
Grade 12 Chemistry
Important Features of the Curriculum
The high school curriculum covers six main themes:
Particulate Nature of Substances
Periodicity
Chemical Calculations
Chemistry of Reactions
The Environment
Organic Chemistry
The Grade 12 Chemistry Textbook consists of eight chapters:
Chapter 1: CHEMICAL BONDING AND INTERMOLECULAR FORCES
Chapter 2: ENERGY CHANGES IN CHEMICAL REACTIONS
Chapter 3: CHEMICAL KINETICS: RATES OF REACTION
Chapter 4: CHEMICAL EQUILIBRIUM
Chapter 5: ACID-BASE REACTIONS
Chapter 6: TRANSITION ELEMENTS
Chapter 7: CHEMISTRY AND GREEN ENVIRONMENT
Chapter 8: ORGANIC COMPOUNDS AND MACROMOLECULES
After learning this course, students will be able to participate actively in all lessons through the 5Cs:
Collaboration: Working in groups to share ideas and find solutions together.
Communication: Developing verbal and non-verbal communication skills in group work.
Critical thinking & problem-solving: Solving interesting problems and looking for correction errors.
Creativity and innovations: Thinking 'outside the box' and exploring new ideas to solve problems in new ways.
Citizenship: Taking part in the school community and developing fairness and conflict resolution skills.
Chapter 1: CHEMICAL BONDING AND INTERMOLECULAR FORCES
1.1 Basic Concepts to Understand Chemical Bonding
Principles of filling electrons in atomic orbitals are fundamental:
Aufbau principle: Electrons fill lower energy orbitals first (1s 2s 2p 3s 3p 4s 3d 4p…).
Pauli’s exclusion principle: No more than two electrons can occupy the same orbital, and they must have opposite spins.
Hund’s rule: Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.
Valence electrons: Electrons in the outermost shells of an atom, equal to its group number in the Periodic Table.
Octet rule: Atoms lose, gain, or share electrons to achieve stable electronic structures of noble gases (eight electrons in the valence shell).
Exceptions to the octet rule can occur with elements of the second period, elements with an odd number of valence electrons, or too few or too many valence electrons
Lewis symbols: Representation of chemically important valence electrons in s and p orbitals using dots.
Chemical bond: Formed when atoms are held together by attractive forces after sharing or transferring valence electrons.
Electronegativity (EN): A measure of the ability of an atom to attract shared electrons to itself.
Electron affinity (EA): The amount of energy released when an electron is added to a gaseous atom.
Ionisation energy (IE): The amount of energy required to remove an electron from a gaseous atom.
Elements having high EA easily gain electrons, forming anions.
Elements having low IE easily lose electrons, forming cations.
Three main types of chemical bonding:
Ionic bonding
Covalent bonding
Metallic bonding
Intermolecular forces: Weak forces between molecules, including hydrogen bonding and van der Waals forces.
Hydrogen bond is the strongest intermolecular force.
Van der Waals forces include ion-dipole interaction, dipole-dipole interaction, and London dispersion forces.
London dispersion forces are the weakest intermolecular forces.
Most bonds have some degree of both ionic and covalent character.
1.2 Ionic Bonding
Ionic bonding (electrovalent bonding) results from the transfer of electrons and electrostatic attractions between cations and anions.
Takes place when the difference in electronegativity is greater than 1.8, generally between reactive metals and reactive non-metals.
Metal loses electrons to form cations, while the non-metal atom gains electrons to form anions.
Number of lost or gained electrons equals the valency of that element.
Ionic compounds are always neutral due to electron loss always equaling electron gain.
Ions arrange themselves into a crystal lattice.
The formula represents the ratio to achieve neutrality.
Greater the charge on the ions the stronger the ionic bond will be.
Solid ionic compounds do not conduct electricity because ions are held in fixed positions.
The crystalline structure is called a crystal lattice, formed by repeating unit cells, such as cubic unit cells.
Types of cubic unit cells: simple cubic (sc), face-centred cubic (fcc), and body-centred cubic (bcc).
Coordination number: The number of ions surrounding the central ion in the lattice.
1.3 Covalent Bonding
Covalent bonding occurs through electron sharing between similar or identical atoms, usually between non-metals.
Covalent bond is the electrostatic attraction between the nuclei and the shared electrons.
Three forces acting on the atoms at the same time:
repulsive force between the electrons of atoms (like charges repel)
attractive force between the nucleus of each atom and the electrons of another
repulsive force between the two positively charged nuclei
Optimum distance between nuclei is called the bond length where net attractive forces are maximised.
Outer orbitals of atoms overlap, sharing unpaired valence electrons.
Covalent bonds can be single, double, or triple bonds.
Single covalent bonds: One pair of electrons is shared.
Double bond: Two pairs of electrons are shared.
Triple bond: Three pairs of electrons are shared.
The higher the number of bonded electrons, the shorter the bond length and the stronger the bond.
Coordinate bond (dative bond): Both shared electrons are donated by one participating atom.
Coordinate bond is expressed as an arrow, →
A coordinate bond is formed when one of the participating atoms possesses an unshared pair (a lone pair) of electrons. This lone pair is donated to an atom needing them to build up, or complete, a stable electron octet or duplet.
Electron-deficient molecules (e.g., BeCl2, BF3, AlCl3) can accept an electron pair from a molecule bearing a lone pair, forming a coordinate compound.
Lewis structures: Two-dimensional representations of valence shell electron arrangement around atoms in a molecule or ion.
Covalent bond is shown as a pair of dots or a line between symbols of two atoms.
Steps to draw Lewis structures:
Calculate total valence electrons.
Write a skeleton arrangement, selecting the central atom.
Subtract shared electrons from total valence electrons and place remaining as unshared pairs.
Insert unshared paired electrons to complete the octet of every element (except H, which can share only 2 electrons).
Some molecules deviate from the octet rule (e.g., NO, BeCl2, BF3, PF5, SF6).
Electron-deficient molecules: Central atom has less than 8 valence electrons.
Expanded valence shell molecules: Central atom has more than 8 valence electrons.
Valence Shell Electron Pair Repulsion (VSEPR) theory: Predicts the three-dimensional geometry of a molecule based on minimizing electron pair repulsion around a central atom.
Steps to predict molecular shape:
Draw the Lewis structure.
Count the number of electron pairs around the central atom.
Determine the geometry of electron pairs (linear, triangular planar, tetrahedral).
Predict the shape of the molecule.
Diatomic molecules are linear.
Repulsive interaction increases in the order: bonding pair-bonding pair < lone pair-bonding pair < lone pair-lone pair.
Giant structures of covalent molecules: Atoms are held together by single covalent bonds, forming regular networks (e.g., diamond, graphite, silicon(IV) oxide).
1.4 Intermolecular Forces
Intermolecular forces: Attractive and repulsive forces between molecules.
Bond polarity: If the atoms are not identical, the electrons will not be equally shared.
Dipole moment: expressed numerically by taking into account the magnitude of the charge and the distance separating the charges.
Polar or non-polar molecules can be determined according to the total dipole moments of the molecules.
Types of intermolecular forces:
van der Waals forces:
dipole-dipole interaction: Partially negative portion of one molecule is attracted to the partially positive portion of another polar molecule.
ion-dipole interaction: Interaction between an ion and a polar molecule.
London dispersion forces: Temporary dipoles in atoms induce electrostatic attraction.
Hydrogen bonding:
First molecule has hydrogen attached to a high electronegative atom (F, O, or N).
Second molecule possesses a lone pair of electrons on a small electronegative atom (F, O, or N).
Hydrogen bond is a special type of dipole- dipole interaction that occurs between polar molecules and is stronger than van der Waals forces
1.5 Metallic Bonding
Metallic bonding: Metal atoms are held together by a 'sea of electrons'.
Valence electrons are loosely held, forming a common electron cloud.
Positively charged metallic nuclei are surrounded by free-moving electrons.
Strength of metallic bonding increases with:
Increasing positive charge on the ions in the metal lattice
Decreasing size of metal ions in the lattice
Increasing number of mobile electrons (valence electrons) per atom.
Metallic bond is lower than that of an ionic bond.
Chapter 2: ENERGY CHANGES IN CHEMICAL REACTIONS
2.1 Energy Changes
Energy can neither be created nor destroyed (Law of Conservation of Energy).
Chemical energy: Energy stored in the bonds of chemical compounds.
Temperature changes, electricity are manifestations of chemical energy changes.
Spontaneous reactions: Occur without the addition of energy.
Non-spontaneous processes need the energy to take place.
Exothermic reactions: Give out energy so products have lower energy than reactants and ΔH is negative (reactants → products + energy).
Endothermic reactions: Take in energy so products have higher energy than reactants and ΔH is positive (reactants + energy → products).
Energy is absorbed to break bonds (endothermic), and energy is released when bonds form (exothermic).
2.2 Enthalpy Changes in Chemical Reactions
Enthalpy (H): The total chemical energy of a substance.
Enthalpy change (ΔH): Heat transferred at constant pressure during a chemical reaction, expressed in kJ mol-1.
Standard enthalpy change (ΔH°): Heat absorbed or released in a process occurring at standard conditions (25 °C and 1 atm).
Thermochemical equation: Balanced chemical equation including physical states and standard enthalpy change (ΔH°).
Types of standard enthalpy changes:
Standard enthalpy change of reaction (ΔH𝑟°): Enthalpy change when molar quantities of reactants react to give products under standard conditions.
Standard enthalpy change of formation (ΔH𝑓°): Enthalpy change when one mole of a compound is formed from its elements under standard conditions.
Standard enthalpy change of combustion (ΔH𝑐°): Enthalpy change when one mole of a substance is burnt completely in oxygen under standard conditions (always negative).
Standard enthalpy change of neutralisation (ΔH𝑛°): Enthalpy change when solutions of an acid and an alkali react to produce one mole of water under standard conditions (always negative).
Enthalpy change (heat changes) can be measured using a calorimeter.
q = mcΔT: Heat change (q) calculation using mass (m), specific heat capacity (c), and temperature change (ΔT).
2.3 Hess’s Law
Hess’s law: Enthalpy change for a chemical reaction is the same, whatever route is taken from reactants to products.
Enthalpy cycle: Diagram illustrating Hess’s law.
Steps to construct enthalpy cycle
ΔHr = ΔH1 + ΔH2 .
Write down the equation for the enthalpy change of the reaction.
Based on the data given, draw the enthalpy cycle with the correct arrow direction.
Write the corresponding symbols for enthalpy changes over the arrows.
Find two routes (clockwise and anticlockwise) and apply Hess\'s law to calculate the required enthalpy change (i.e., Route I is equal to Route II).
Enthalpy change of reaction from enthalpy changes of formation: ΔH1 + ΔHr = ΔH2 .
Enthalpy change of formation from enthalpy changes of combustion: ΔH1 = ΔHf + ΔH2 .
Bond energy (bond enthalpy, E): Energy needed to break or make bonds. Bond breaking absorbs energy (ΔH is positive), and bond making releases energy (ΔH is negative).
Hess’s law with bond enthalpies: estimate approximate enthalpy change, ΔHr θ = ΔH1 + ΔH2
Chapter 3: CHEMICAL KINETICS: RATES OF REACTION
3.1 Reaction Rates
Chemical kinetics: Study of reaction rates.
Rate of reaction: Change in the amount (mass, volume, or concentration) of reactants or products per unit of time. Units: pressure, volume, or concentration.
Rate = \frac{change\ in\ amount\ (mass,volume,or\ concentration)}{time\ taken\ for \ this \ change}
Rate of reactant used up = - \frac{∆[reactant]}{∆t}
Rate of product formed = +\frac{∆[product]}{∆t}
For a general chemical reaction aA + bB \longrightarrow cC + dD the rate for the reaction is:
rate = -\frac{1}{a} \frac{∆[A]}{∆t} = -\frac{1}{b} \frac{∆[B]}{∆t} = \frac{1}{c} \frac{∆[C]}{∆t} = \frac{1}{d} \frac{∆[D]}{∆t}
3.2 Collision Theory and Activation Energy
Collision theory: Reacting particles must collide in the correct orientation with sufficient energy to overcome the energy barrier to react.
frequency of collision should be increased
proportion of particles with energy that is greater than the activation energy increases.
Activation energy (Ea): Minimum energy reactant particles must possess for a successful collision to take place.
H2 (g) + I2 (g) ⇌ [H2I2] → 2HI (g)
Activated complex is unstable
High Ea means relatively few collisions, and so the reaction rate is slow.
3.3 Factors Affecting Reaction Rates
Concentration: Increasing concentration increases the rate of reaction.
Particle size: Smaller particles (larger surface area) result in faster reactions.
Pressure: Increasing pressure increases the rate of reactions involving gases.
Temperature: Increasing temperature generally increases the rate of reaction.
Boltzmann distribution: Distribution of molecular energies at a given temperature.
Catalyst: Lowers the activation energy.
Positive catalysts speed reactions.
Negative catalysts slow reactions.
Homogeneous catalysts are in the same phase as reactants.
Heterogeneous catalysts are in a different phase from reactants.
Enzymes: Biocatalysts for specific biochemical reactions.
Photochemical reactions: Reactions initiated by light, where brighter light means a faster reaction.
Chapter 4: CHEMICAL EQUILIBRIUM
4.1 State of Dynamic Equilibrium
Dynamic equilibrium: Reactants are continuously being changed to products and products are continuously being changed back to reactants.
Reversible reaction: A reaction that can occur in both directions.
Reversible reactions in a closed system eventually reach equilibrium.
In dynamic equilibrium, the forward and reverse reactions occur at the same rate, leading to constant concentrations of the reactants and products.
Classification of chemical equilibria:
Homogeneous chemical equilibrium: Reactants and products are all in a single phase.
Heterogeneous chemical equilibrium: Reactants and products are present in more than one phase.
4.2 Le Chatelier’s Principle and Factors Affecting the Chemical Equilibrium
Le Chatelier’s principle: If a system is in equilibrium and a factor affecting the chemical equilibrium (temperature, pressure, or concentration) is changed, the equilibrium shifts to cancel out the effect of the change.
Effect of concentration: Increasing reactant concentration shifts equilibrium to products (right), and increasing product concentration shifts equilibrium to reactants (left).
Effect of temperature: In exothermic reactions (A ⇌ B ΔH = −), decreasing temperature favors products.
In endothermic reactions (A ⇌ B ΔH = +), increasing temperature favors products.Effect of pressure: Increasing pressure shifts equilibrium to the side with fewer moles of gas. If moles of gas are equal on both sides, pressure has no effect.
4.3 Equilibrium Constants
Law of mass action: For a reversible reaction at equilibrium and a constant temperature, the ratio between the concentration of reactants and products is constant.
Equilibrium constant (Keq): Ratio of product concentrations to reactant concentrations, each raised to the power of its coefficient in the balanced chemical equation. The value is constant only at a specified temperature.
Keq =\frac{{\left[ C \right]^c \left[ D \right]^d}}{{\left[ A \right]^a \left[ B \right]^b}}
Equilibrium lies to the right and there are more products than reactants when {Keq > > 1}.
Equilibrium lies to the left and there are more reactants than products when {Keq < < 1}.
Reactants and products are equally favored and concentration of reactants and products will be equal when {Keq = 1}.
The units of the equilibrium constant depend on the form of the equilibrium expression.
Pure solids or liquids are excluded from the equilibrium expression because their concentrations remain constant.
Equilibrium constant Kc: Molar are used for the equilibrium.
Equilibrium constant Kp: Pressure can be used for equilibriums that involves gases. Then, the equilibrium constant can be given as the ratio of the partial pressures of products and reactants
Chapter 5: ACID-BASE REACTIONS
5.1 Theories of Acids and Bases
Arrhenius theory:
Acids produce hydrogen ions (H+) in water.
Bases produce hydroxide ions (OH–) in water.
Brønsted-Lowry theory:
Acids are proton (H+) donors.
Bases are proton (H+) acceptors.
Conjugate acid-base pair: Acid and its conjugate base or base and its conjugate acid that relates to the acid or base by the gain or loss of a single hydrogen ion.
Lewis theory:
Acids accept an electron pair.
Bases donate an electron pair.
5.2 Ionic Dissociation of Water and pH
Autoionisation of water: water can show both acidic and basic behaviours.
H2O (l) ⇌ H+ (aq) + OH– (aq)
*ionic product The ionic product of water (Kw) is Kw ={[H^+][OH^-]}=1.0 x 10^{-14} {mol^2dm^{-6}} at 25 °C (298 K).
pH measures the acidity or alkalinity of a solution.
pH = -log[H^+] ; [H^+] = 10^{-pH}
pOH = -log[OH^-] ; [OH^-] = 10^{-pOH}
pH + pOH = 14
5.3 Ionisation of Acids and Bases
Strong acids: Completely ionise in water.
Weak acids: Partially ionise in water.
The basicity of an acidity, what shows the number of H+ ions that one molecule of acid can produce? is the molecule will be one, two, or three:
acids such can Hydrochloric acid (HCl), ethanoic acid (CH3COOH) can be monoprotic
acids such can H2SO4, H2CO3 can be diprotic
acids such can H3PO4 can be triprotic.
Acid dissociation constant (Ka): Measure of acid strength (HA (aq) \rightleftharpoons H^+ (aq) + A^– (aq); Ka =\frac{{\left[ H^+ \right] \left[ A^- \right]}}{{\left[ HA \right]}}).
Strong bases completely ionise in water. (hydroxides of group I and II except those with Be and Mg).
Weak bases partially ionise in water.
Base dissociation constant (Kb): Measure of base strength (B (aq) + H2O (l) \rightleftharpoons HB^+ (aq) + OH^– (aq); Kb = \frac{{\left[ HB^+ \right] \left[ OH^- \right]}}{{\left[ B \right]}}).
5.4 Salt Hydrolysis
Salts of strong acids and strong bases are neutral.
Salts of strong bases and weak acids are basic.
Anion (conjugate base of weak acid) reacts with water to produce OH–.
Salts of weak bases and strong acids are acidic.
Cation (conjugate acid of weak base) reacts with water to produce H+.
Salts of weak acids and weak bases depend on the relative strengths of Ka and Kb.
If Kb > Ka then the solution must be basic.
If Kb < Ka then the solution will be acidic.
If Kb ≈ Ka then the solution will be nearly neutral