Chemistry

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.

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. ...