Chemistry

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 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 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.

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: Sufficient Energy: Energy $\ge$ Activation Energy ($E_a$). 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 ...

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

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. ...

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