Guide on SOH-CAH-TOA

C8 - Periodic Table

Sat, 27 Jun 2026 20:07:18 +0700

The Periodic Table

1. General Structure

Arrangement

Proton Number: Elements are arranged in order of increasing proton number.
Periods: Horizontal rows.
Groups: Vertical columns.
Periodicity: Elements in the same group have similar chemical properties because they have the same number of valence electrons.

Trends Across a Period

Character: Transition from metallic character (left) to non-metallic character (right).
Ion Charge: Group number relates to the charge of the ion formed to achieve a noble gas configuration.

2. Group I: Alkali Metals

Elements: Lithium (Li), Sodium (Na), Potassium (K), etc.
Physical Properties: Relatively soft, low density.
Trends (Down the Group):
- Melting Point: Decreases.
- Density: Increases.
- Reactivity: Increases (outer electron is further from nucleus, easier to lose).

3. Group VII: Halogens

Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I).
General Nature: Diatomic non-metals ($\text{F}_2, \text{Cl}_2, \text{Br}_2, \text{I}_2$).
Appearance:
- $\text{Cl}_2$: Pale yellow-green gas.
- $\text{Br}_2$: Red-brown liquid.
- $\text{I}_2$: Grey-black solid.
Trends (Down the Group):
- Density: Increases.
- Reactivity: Decreases (harder to attract/gain an electron as shell number increases).
Displacement Reactions: A more reactive halogen will displace a less reactive halogen from its salt.
- Example: $\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$

4. Transition Elements

Location: Middle block of the Periodic Table.
Properties:
- High density.
- High melting points.
- Form coloured compounds.
- Often act as catalysts (e.g., Fe in Haber process).
- Often exhibit variable oxidation numbers (e.g., $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$).

5. Group VIII: Noble Gases

Elements: Helium (He), Neon (Ne), Argon (Ar), etc.
General Nature: Monatomic gases.
Reactivity: Very unreactive (inert) because they have full outer electron shells.

6. Group Trends Analysis

To identify group trends from provided data, look for consistent increases or decreases in physical or chemical properties (e.g., boiling point, atomic radius, ionization energy) as the atomic number increases within a group.

C7 - Acids, Bases and Salts

Sat, 27 Jun 2026 20:07:11 +0700

Acids, Bases and Salts

1. Acids

Properties and Indicators

General Properties: Sour taste, corrosive, conduct electricity in aqueous solution.
Chemical Reactions:
- $\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen}$
- $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$
- $\text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}$
Indicator Effects:

Indicator Acidic Condition

Litmus Red

Thymolphthalein Colourless

Methyl Orange Red

Indicator	Acidic Condition
Litmus	Red
Thymolphthalein	Colourless
Methyl Orange	Red

Definitions and Strength

Proton Donor: Acids are substances that donate protons ($\text{H}^+$ ions) to other substances.
Strong Acids: Completely dissociate in aqueous solution (e.g., $\text{HCl}$, $\text{H}_2\text{SO}_4$).
Weak Acids: Partially dissociate in aqueous solution (e.g., $\text{CH}_3\text{COOH}$).

2. Bases and Alkalis

Definitions and Properties

Base: A substance (usually a metal oxide or hydroxide) that neutralises an acid.
Alkali: A soluble base.
Chemical Reactions:
- $\text{Base} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$
- $\text{Base} + \text{Ammonium Salt} \rightarrow \text{Salt} + \text{Water} + \text{Ammonia}$
Indicator Effects:

Indicator Alkaline Condition

Litmus Blue

Thymolphthalein Blue

Methyl Orange Yellow

Indicator	Alkaline Condition
Litmus	Blue
Thymolphthalein	Blue
Methyl Orange	Yellow

Oxides

Basic Oxides: Metal oxides that react with acids (e.g., $\text{CuO}$, $\text{CaO}$).
Acidic Oxides: Non-metal oxides that react with bases (e.g., $\text{SO}_2$, $\text{CO}_2$).
Amphoteric Oxides: Oxides that react with both acids and bases (e.g., $\text{Al}_2\text{O}_3$, $\text{ZnO}$).

3. pH and Neutralisation

The pH Scale

$\text{H}^+$ Concentration: Acids contain $\text{H}^+$ ions; alkalis contain $\text{OH}^-$ ions.
Universal Indicator: Used to compare acidity/alkalinity across the pH scale (0-14).
pH 7: Neutral.

Neutralisation

Ionic Equation: $\text{H}^+(\text{aq}) + \text{OH}^-(\text{aq}) \rightarrow \text{H}_2\text{O}(\text{l})$
This reaction occurs when an acid and a base react to form a salt and water.

4. Salts

Solubility Rules

Soluble	Insoluble
All $\text{Na}^+, \text{K}^+, \text{NH}_4^+, \text{NO}_3^-$ salts	All $\text{CO}_3^{2-}$ except $\text{Na}, \text{K}, \text{NH}_4$
All $\text{Cl}^-$ except $\text{Pb}^{2+}, \text{Ag}^+$	$\text{OH}^-$ except $\text{Na}, \text{K}, \text{NH}_4, \text{Ca}$
All $\text{SO}_4^{2-}$ except $\text{Ba}^{2+}, \text{Ca}^{2+}, \text{Pb}^{2+}$

Preparation of Salts

Insoluble Salts: Prepared by precipitation (mixing two soluble salts).
Soluble Salts:
- Titration: Used when both reactants are solutions (Acid + Alkali).
- Excess Solid Method: Used when one reactant is an insoluble base or metal.
  - $\text{Acid} + \text{Excess Metal/Base/Carbonate} \rightarrow \text{Salt} + \text{H}_2/\text{H}_2\text{O}/(\text{H}_2\text{O} + \text{CO}_2)$.
  - Excess is filtered off, then the solution is evaporated to crystallisation.

Hydration

Hydrated Salt: A salt that contains chemically combined water (e.g., $\text{CuSO}_4\cdot 5\text{H}_2\text{O}$).
Anhydrous Salt: A salt that contains no water.
Water of Crystallisation: The fixed amount of water molecules associated with each formula unit of a salt.

C6 - Chemical Reactions

Sat, 27 Jun 2026 20:07:02 +0700

Chemical Reactions

1. Physical and Chemical Changes

Physical Change

No new substance formed.
Often reversible (e.g., melting ice).
Change in state or shape.

Chemical Change

New substance(s) formed.
Often irreversible.
Accompanied by energy change, colour change, or gas evolution.

2. Rates of Reaction

Collision Theory

For a reaction to occur, particles must collide with:
1. Sufficient Energy: Energy $\ge$ Activation Energy ($E_a$).
2. Correct Orientation.
Effect of Factors:
- Temp/Concentration/Pressure: Increase collision frequency and/or proportion of particles with energy $\ge E_a$.
- Catalysts: Provide an alternative pathway with a lower activation energy ($E_a$).

Factors Affecting Rate

Concentration: Higher concentration $\rightarrow$ more particles per unit volume $\rightarrow$ higher rate.
Pressure (Gases): Higher pressure $\rightarrow$ particles closer together $\rightarrow$ higher rate.
Surface Area: Smaller particles (powder) $\rightarrow$ more exposed surface $\rightarrow$ higher rate.
Temperature: Higher temperature $\rightarrow$ particles move faster $\rightarrow$ higher rate.
Catalyst: Substance that increases rate without being consumed. Enzymes are biological catalysts.

Investigating Rates

Methods:
- Measuring mass loss (if gas escapes).
- Measuring volume of gas produced (using gas syringe).
- Measuring time for a colour change or precipitate to form.
Evaluation: Consider accuracy of apparatus (e.g., gas syringe vs. measuring cylinder) and precision of timing.

3. Reversible Reactions and Equilibrium

Reversible Reactions

Reactions that can proceed in both forward and reverse directions.
Symbol: $\rightleftharpoons$
Hydrated vs Anhydrous:
- $\text{CuSO}_4\cdot 5\text{H}_2\text{O}$ (Blue) $\rightleftharpoons \text{CuSO}_4$ (White) + $5\text{H}_2\text{O}$
- $\text{CoCl}_2\cdot 6\text{H}_2\text{O}$ (Pink) $\rightleftharpoons \text{CoCl}_2$ (Blue) + $6\text{H}_2\text{O}$

Dynamic Equilibrium

Occurs in a closed system.
Condition: Rate of forward reaction = Rate of reverse reaction.
Observation: Concentrations of reactants and products remain constant.

Equilibrium Shifts

Le Chatelier’s Principle: If a system at equilibrium is stressed, it shifts to oppose the change.
Temperature: Increase in temp shifts equilibrium in the endothermic direction.
Pressure: Increase in pressure shifts equilibrium towards the side with fewer gas molecules.
Concentration: Increasing a reactant shifts equilibrium towards the products.

4. Industrial Processes

Haber Process (Ammonia Synthesis)

Equation: $\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})$
Sources: $\text{N}_2$ from air, $\text{H}_2$ from methane (natural gas).
Conditions: 450 °C, 20,000 kPa, Iron catalyst.
Optimization: Balance between rate (high temp) and yield (low temp for exothermic reaction).

Contact Process (Sulfuric Acid Synthesis)

Equation: $2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g})$
Sources: $\text{SO}_2$ from sulfur burning/roasting, $\text{O}_2$ from air.
Conditions: 450 °C, 200 kPa, Vanadium(V) oxide catalyst.

5. Redox Reactions

Basic Definitions

Redox: A reaction where oxidation and reduction happen simultaneously.
Oxidation: Gain of oxygen.
Reduction: Loss of oxygen.
Identification: Look for gain/loss of oxygen in equations.

Advanced Redox

C5 - Chemical Energetics

Sat, 27 Jun 2026 20:06:31 +0700

Chemical Energetics

1. Exothermic and Endothermic Reactions

Exothermic Reactions

Definition: Reactions that transfer heat energy to the surroundings.
Observation: Temperature of the surroundings increases.
Example: Combustion of fuels, neutralisation reactions.

Endothermic Reactions

Definition: Reactions that take in heat energy from the surroundings.
Observation: Temperature of the surroundings decreases.
Example: Thermal decomposition, photosynthesis.

Reaction Pathway Diagrams

Exothermic: Reactants have more energy than products. The difference is released as heat.
Endothermic: Products have more energy than reactants. The difference is absorbed from surroundings.

2. Enthalpy and Activation Energy

Enthalpy Change ($\Delta H$)

Definition: The heat energy change during a chemical reaction.
Exothermic: $\Delta H$ is negative (e.g., $\Delta H = -100 \text{ kJ/mol}$).
Endothermic: $\Delta H$ is positive (e.g., $\Delta H = +100 \text{ kJ/mol}$).

Activation Energy ($E_a$)

Definition: The minimum energy that colliding particles must possess for a reaction to occur.
Diagramming: Represented as a “hump” or energy barrier on a reaction pathway diagram.
Labelling: In a diagram, $E_a$ is measured from the energy level of the reactants to the peak of the curve.

3. Bond Energies

Bond Breaking and Making

Bond Breaking: Requires energy (Endothermic).
Bond Making: Releases energy (Exothermic).

Calculating Enthalpy Change ($\Delta H$)

Formula: $$\Delta H = \sum (\text{bond energies of reactants}) - \sum (\text{bond energies of products})$$

C4 - Electrochemistry

Sat, 27 Jun 2026 20:06:22 +0700

Electrochemistry

1. Fundamentals of Electrolysis

Definitions

Electrolysis: Decomposition of an ionic compound (molten or aqueous) by an electric current.
Electrolyte: An ionic compound that conducts electricity when molten or in aqueous solution.
Electrodes:
- Anode: Positive electrode (+).
- Cathode: Negative electrode (-).

Charge Transfer and Ion Movement

External Circuit: Electrons flow from anode to cathode.
Electrodes:
- Anode: Oxidation occurs (loss of electrons).
- Cathode: Reduction occurs (gain of electrons).
Electrolyte: Cations move to the cathode; anions move to the anode.

C3 - Stoichiometry

Sat, 27 Jun 2026 20:06:15 +0700

Stoichiometry

1. Formulae and Equations

Chemical Formulae

Formula of Elements: Represents the simplest ratio of atoms (e.g., $\text{O}_2$, $\text{P}_4$).
Molecular Formula: Shows the actual number of atoms of each element in a molecule of a compound.
Deducing Formulae: Determine formulae from molecular models or diagrams.

Chemical Equations

Word Equations: Use names of reactants and products.
Symbol Equations: Use chemical symbols and formulae.
State Symbols:
- $(s)$ solid
- $(l)$ liquid
- $(g)$ gas
- $(aq)$ aqueous solution
Ionic Equations: Show only the species that change during the reaction; spectator ions are omitted.

2. Relative Masses

Relative Atomic Mass ($A_r$)

The weighted average mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12.

Relative Molecular Mass ($M_r$) and Formula Mass

Relative Molecular Mass ($M_r$): Sum of the relative atomic masses of all atoms in a molecule.
Relative Formula Mass: Sum of relative atomic masses for ionic compounds (where molecular formula is not applicable).

3. The Mole Concept

The Mole and Avogadro Constant

The Mole ($\text{mol}$): The unit for amount of substance.
Avogadro Constant: $6.02 \times 10^{23}$ particles per mole.
Calculation: $\text{Number of particles} = \text{moles} \times (6.02 \times 10^{23})$

Molar Mass Calculations

Formula: $$\text{mass (g)} = \text{moles (mol)} \times \text{molar mass (g/mol)}$$
$\text{Molar mass}$ is numerically equal to $M_r$ or $A_r$.

Empirical and Molecular Formulae

Empirical Formula: The simplest whole-number ratio of atoms of each element in a compound.
Molecular Formula: A multiple of the empirical formula.
Calculation: Determine empirical formula from percentage composition or mass, then use $M_r$ to find the molecular formula.

4. Stoichiometry and Reacting Masses

Simple Proportions

Calculate reacting masses using the ratio of $M_r$ without the mole concept.

Stoichiometric Calculations

Use balanced symbol equations to determine the molar ratio between reactants and products.
Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

5. Gas and Solution Stoichiometry

Molar Gas Volume

At room temperature and pressure (r.t.p.), 1 mole of any gas occupies $24\text{ dm}^3$.
Formula: $\text{volume (dm}^3) = \text{moles} \times 24$

Solution Concentrations

Mass Concentration: $\text{g/dm}^3$
Molar Concentration: $\text{mol/dm}^3$
Conversion: $\text{mol/dm}^3 = \frac{\text{g/dm}^3}{\text{molar mass}}$

Titrations

Use titration data (volume and concentration of a known solution) to calculate the unknown concentration or volume of another solution.

6. Yield and Purity

$$\text{Percentage Yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100%$$

C1 - States of Matter

Sat, 27 Jun 2026 20:06:05 +0700

States of Matter

1. Solids, Liquids, and Gases

Distinguishing Properties

Solids: Fixed shape and volume. Incompressible. High density.
Liquids: Fixed volume but no fixed shape (take shape of container). Slightly compressible.
Gases: No fixed shape or volume. Highly compressible. Low density.

Particle Structure

State	Arrangement	Separation	Motion
Solid	Regular lattice	Very close	Vibrate about fixed positions
Liquid	Random/Irregular	Close	Slide over each other
Gas	Random	Far apart	Rapid and random in all directions

2. Changes of State

Definitions

C2 - Atoms Elements Compounds

Sat, 27 Jun 2026 20:05:54 +0700

Atoms, Elements and Compounds

1. Basic Definitions

Elements, Compounds and Mixtures

Element: Pure substance consisting of only one type of atom.
Compound: Pure substance consisting of two or more different elements chemically bonded together.
Mixture: Two or more substances physically blended but not chemically bonded.

2. Atomic Structure

The Atom

Structure: Nucleus (containing protons and neutrons) surrounded by electrons in shells.

Subatomic Particles:

Particle	Relative Charge	Relative Mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	1/1840	Shells

Atomic and Mass Numbers

Proton Number (Atomic Number): Number of protons in the nucleus of an atom.
Mass Number (Nucleon Number): Total number of protons and neutrons in the nucleus.

Electronic Configuration

Electrons occupy shells around the nucleus.
Configuration for proton numbers 1-20: 2, 8, 8, 2.
Periodic Table Relation:
- Group Number: Equals the number of electrons in the outer shell (Group I-VII).
- Period Number: Equals the number of occupied shells.
- Group VIII (Noble Gases): Have full outer shells, making them unreactive.

Isotopes

Definition: Atoms of the same element with the same number of protons but different numbers of neutrons.
Chemical Properties: Isotopes have identical chemical properties because they have the same electronic configuration.
Relative Atomic Mass: Calculated using the weighted average of isotope abundances: $\text{Relative Atomic Mass} = \frac{\sum (\text{isotope mass} \times \text{abundance})}{\text{total abundance}}$

3. Ions and Bonding

Ion Formation

Cations: Positive ions formed when an atom loses electrons.
Anions: Negative ions formed when an atom gains electrons.

Ionic Bonding

Definition: Strong electrostatic attraction between oppositely charged ions.
Formation: Typically between Group I (metal) and Group VII (non-metal).
Formation: Occurs between any metallic and non-metallic elements.
Representation: Use dot-and-cross diagrams to show electron transfer.
Structure: Exists as a giant ionic lattice.
Properties:
- High melting and boiling points due to strong electrostatic forces throughout the lattice.
- Conduct electricity when molten or aqueous (ions are free to move).
- Poor conductors when solid (ions fixed in position).

Covalent Bonding

Definition: A pair of shared electrons between two atoms to achieve noble gas configurations.
Simple Molecules:
- Examples: $\text{H}_2$, $\text{Cl}_2$, $\text{H}_2\text{O}$, $\text{CH}_4$, $\text{NH}_3$, $\text{HCl}$.
- Representation: Use dot-and-cross diagrams.
Simple Molecules:
- Examples: $\text{CH}_3\text{OH}$, $\text{C}_2\text{H}_4$, $\text{O}_2$, $\text{CO}_2$, $\text{N}_2$.
Properties:
- Low melting and boiling points due to weak intermolecular forces, despite strong covalent bonds within the molecule.
- Poor electrical conductivity (no free ions or electrons).

Giant Covalent Structures

Diamond: Each carbon atom bonded to four others in a tetrahedral lattice. Extremely hard, used in cutting tools.
Graphite: Each carbon atom bonded to three others in hexagonal layers. Layers slide (lubricant) and delocalised electrons allow conductivity (electrode).
Silicon(IV) Oxide ($\text{SiO}_2$): Structure similar to diamond (Si bonded to 4 O atoms). High melting point and hard.

Metallic Bonding

Definition: Electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons.
Properties:
- Electrical Conductivity: Delocalised electrons are free to move and carry charge.
- Malleability and Ductility: Layers of ions can slide over each other without breaking the metallic bond.

C9 - Metals

Tue, 23 Jun 2026 22:30:00 +0700

Comprehensive guide covering the chemical and physical properties of metals, reactivity, extraction, and corrosion.