CSEC Chemistry: Atomic Structure, Bonding, and Reactions Review

Atomic Structure, Bonding, and Periodicity

Comparison of Sodium Chloride (NaCl) and Calcium Oxide (CaO) * Both substances possess giant ionic crystalline structures with high melting points. * The melting point of Calcium Oxide is significantly higher than that of Sodium Chloride because Calcium has a charge of $+2$ ( $Ca^{2+}$ ) while Sodium has a charge of $+1$ ( $Na^+$ ). Due to the higher charge, Calcium is more strongly bonded to Oxygen in the lattice, requiring greater energy to break the ionic bonds and melt the solid.
Atomic Arrangement of Chlorine ( $Cl$ ) * Atomic Number: $17$ * Proton count: $17$ ; Neutron count: $18$ * Electronic Configuration: $2, 8, 7$
Bonding in Iodine Monochloride ( $ICl$ ) * Iodine ( $I$ ) has an electronic configuration of $2, 8, 18, 18, 7$ with $7$ valence electrons. * Chlorine ( $Cl$ ) has $7$ valence electrons. * Type of bonding: Covalent bond formed by the sharing of one pair of electrons. * Reasons for low melting point: Covalent compounds are typically held together by weak Van der Waals forces, and little energy is required to separate the molecules. * Isotopic Variation: Samples of $ICl$ may have different molar masses due to different isotopes of Chlorine. Isotopes are forms of the same element with the same proton number but different neutron (and mass) numbers.
Properties of Oxides of Magnesium and Sulfur * Magnesium Oxide ( $MgO$ ): Metal oxide with ionic bonding ( $EC: 2, 8, 2$ for $Mg$ ). It exists as a solid at room temperature with a melting point of $2852\,^{\circ}C$ . Strong electrostatic forces in the ionic lattice require high energy to break. * Sulfur Oxide ( $SO_2$ ): Non-metal oxide with covalent bonding ( $EC: 2, 8, 6$ for $S$ ). It is a gas at room temperature with a melting point of $-72\,^{\circ}C$ . Weak Van der Waals interactions require minimal energy to break.
Periodic Trends and Ionization * Ease of Ionization: Increases down a metallic group. As atomic radius increases, the attractive pull of the positive nucleus on valence electrons decreases, allowing atoms to lose electrons more easily to form cations. Example: Element $X$ (below Magnesium) has a greater ease of ionization than Magnesium ( $Mg$ ). * Oxidizing Power: Increases up a non-metallic group. As atomic radius decreases, the attractive pull of the nucleus on electrons to be gained increases, making the atoms gain electrons more readily to form anions. Example: Chlorine ( $Cl$ ) has a greater oxidizing power than Element $Y$ (below Chlorine).

States of Matter and Particle Theory

Primary States: Solid, Liquid, and Gas.
Forces of Attraction: Solids have strong forces, liquids have moderate forces, and gases have weak forces of attraction between particles.
Phase Changes: * Melting Point: The constant temperature at which a solid changes into a liquid. * Boiling Point: The temperature at which a liquid changes into a gas. * Condensation: Process of conversion from water vapor (gas) to liquid water. * Sublimation: The process where a substance (e.g., Iodine) changes directly from a solid state to a gaseous state upon heating.
Cooling Curve Analysis: * During cooling, the temperature remains constant during a phase change (e.g., at $65\,^{\circ}C$ for unknown solid A) as kinetic energy changes into potential energy. * A horizontal plateau on the graph identifies the melting point.

Radioactivity, Isotopy, and Allotropy

Isotopes: Different atoms of a single element with the same number of protons but a different number of neutrons (e.g., $^{35}Cl$ and $^{37}Cl$ ).
Radioisotopes and Their Uses: * Carbon-14: Used for dating archaeological specimens. * Uranium-235: Used for electricity generation in nuclear power stations. * Iodine-131: Used for imaging and treating thyroid disorders; ideal due to its short half-life of $8$ days. * Americium-241: Used in smoke detectors. * Plutonium-238: Powers spacecraft (e.g., NASA missions) by emitting steady heat during radioactive decay and used for long-lasting pacemaker batteries ( $T_{1/2} \approx 87\,years$ ).
Allotropes of Carbon: * Diamond: Giant covalent cubic lattice where every carbon is bonded to $4$ others. It does not conduct electricity because it has no mobile electrons. * Graphite: Made of layers of carbon where each carbon is bonded to $3$ others. The presence of delocalized (mobile) electrons allows it to conduct electricity.

Separation of Mixtures and Solubility

Fermentation: A chemical reaction where carbohydrates (e.g., glucose) are converted into ethanol and carbon dioxide by yeast (using the enzyme zymase) under anaerobic conditions. * Equation: $C_6H_{12}O_6(aq) \xrightarrow{\text{zymase}} 2C_2H_5OH(aq) + 2CO_2(g)$
Distillation: Necessary to separate ethanol from the fermentation mixture. Fractional distillation uses a fractionating column packed with glass beads to achieve better separation based on boiling point differences.
Sugar Manufacturing Process: * Mill Train: Crushes sugar cane to produce juice and bagasse (used as fuel in boiler furnaces). * Clarification (Process P): Purifies the juice. * Boiling/Crystallization (Process Q): Produces sugar crystals and molasses. * Centrifugation: Separates sugar crystals from product X (molasses).
Factors Affecting Dissolving Rate: Temperature, concentration, and surface area (particle size).
Solution vs. Suspension: A solution is a homogeneous mixture; a suspension is a heterogeneous mixture.

Metals: Reactivity, Extraction, and Applications

Metals in Biology: * Magnesium ( $Mg$ ): Essential for chlorophyll production; deficiency leads to yellowing of leaves (chlorosis) and low crop yield. * Iron ( $Fe$ ): Important for plant health; deficiency also causes chlorosis. In humans, it is found in hemoglobin to carry oxygen; deficiency leads to anemia.
Environmental Hazards: * Lead ( $Pb$ ): Found in exhaust from leaded petrol; damages tissues/organs, reduces IQ in children, and causes anemia. * Carbon Monoxide ( $CO$ ): From incomplete combustion; binds to hemoglobin more readily than oxygen, causing dizziness, headaches, and death.
Reactivity Series: * A more reactive metal displaces a less reactive metal from its compound. * Order: K > Na > Ca > Mg > Al > Zn > Fe > Pb > (H) > Cu > Hg > Ag > Au.
Metal Extraction: * Electrolysis: Required for highly reactive metals (e.g., $Al$ from Bauxite/Alumina) because they form very stable ions. * Reduction with Carbon: Suitable for lower reactivity metals (e.g., $Fe$ using coke/carbon monoxide) as their ions are less stable.
Alloys: Mixtures of metals (and sometimes non-metals) designed to improve properties like hardness and corrosion resistance. * Stainless Steel: $70\%\,Fe$ , $20\%\,Cr$ , $10\%\,Ni$ . Used in cutlery and surgical equipment. * Magnalium: $95\%\,Al$ , $5\%\,Mg$ . Used in aircraft construction. * Duralumin: $94\%\,Al$ , $4\%\,Cu$ , trace $Mn$ and $Mg$ . Harder and more resistant to corrosion than pure aluminium.
Corrosion: * Aluminium: Beneficial; forms an unreactive oxide layer ( $Al_2O_3$ ) that protects the metal below. * Iron: Harmful; forms rust ( $Fe_2O_3 \cdot xH_2O$ ) which flakes off, exposing fresh iron to further decay.

Mole Concept and Quantitative Analysis

Definitions: * Mole: Amount of substance containing $6.02 \times 10^{23}$ particles. * Molar Mass ( $M$ ): Mass in grams of one mole of a chemical substance.
Standard Calculations: * Molar Volume at RTP: $1\,mole = 24\,000\,cm^3$ or $24\,dm^3$ . * Calculation for mass from moles: $mass = moles \times Molar\,Mass$ . * Heat transfer: $Q = m \times c \times \Delta T$ , where $c = 4.2\,J\,g^{-1}\,^{\circ}C^{-1}$ for water.

Rates of Reaction and Energetics

Rate of Reaction: Measured change in the concentration of a reactant or product per unit time.
Influencing Factors: Concentration, temperature, surface area (particle size), and catalysts.
Temperature Effect: Increases kinetic energy; particles move faster and collide more frequently with energy exceeding the activation energy.
Surface Area Effect: Smaller particles (powder) offer more surface area for collisions, increasing the reaction rate.
Catalysts: Increase rate by providing an alternative pathway with a lower activation energy ( $E_a$ ).
Exothermic vs. Endothermic: * Exothermic: Releases energy ( $\Delta H = negative$ ). * Endothermic: Absorbs energy ( $\Delta H = positive$ ).

Electrolysis

Definitions: * Electrolysis: Chemical decomposition by passing an electric current through an electrolyte. * Electrolyte: A compound that forms ions when molten or in aqueous solution. * Anode: Positive electrode (oxidation occurs). * Cathode: Negative electrode (reduction occurs).
Faraday's Calculations: * Charge ( $Q$ ) in Coulombs: $Q = I \times t$ , where $I$ is current in Amperes and $t$ is time in seconds. * $1\,Faraday = 96\,500\,C\,mol^{-1}$ .
Preferential Discharge Rules: Ions are discharged based on their concentration (for halides) and their position in the electrochemical series.

Redox Reactions

Oxidation: Loss of electrons; increase in oxidation state.
Reduction: Gain of electrons; decrease in oxidation state.
Oxidising Agent: Causes another substance to be oxidized; it is itself reduced. Example: Acidified Potassium Manganate ( $VII$ ) changes from purple to colorless ( $Mn^{+7} \rightarrow Mn^{+2}$ ).
Reducing Agent: Causes another substance to be reduced; it is itself oxidized.

Acids, Bases, and Salts

Definitions: * Salt: Compound formed when hydrogen ions in an acid are replaced by metal or ammonium ions. * Strong Acid: Fully ionizes in solution (e.g., $HCl$ ). * Weak Acid: Partially ionizes in solution (e.g., Vinegar/Ethanoic acid).
Salt Preparation: * Insoluble Salts (e.g., $CaCO_3$ ): Prepared via ionic precipitation (mixing two soluble solutions and filtering). * Soluble Salts (e.g., $NaNO_3$ , $KNO_3$ ): Prepared via titration or reaction of acid with metal/base/carbonate followed by crystallization.
Hard Water: Caused by $Ca^{2+}$ and $Mg^{2+}$ ions that react with soap to form scum. Permanent hardness is caused by sulfates ( $CaSO_4$ , $MgSO_4$ ) and cannot be removed by boiling.

Organic Chemistry and Hydrocarbons

Homologous Series: Groups of organic compounds with the same functional group and a general formula (e.g., Alkanes: $C_nH_{2n+2}$ , Alkenes: $C_nH_{2n}$ , Alcohols: $C_nH_{2n+1}OH$ ).
Structural Isomerism: Existence of compounds with the same molecular formula but different connectivity/arrangement of atoms.
Reactions: * Alkanes: Undergo substitution (halogenation) in the presence of UV light. * Alkenes: Undergo addition reactions (hydrogenation, hydration to form alcohols, halogenation). * Esterification: Reaction between an alcohol and a carboxylic acid (e.g., ethanol + ethanoic acid $\rightarrow$ ethyl ethanoate + water).
Polymerization: * Addition: Monomer double bonds ( $C=C$ ) break and join (e.g., Polyethene). * Condensation: Monomers join by eliminating a small molecule like water (e.g., Polyesters, Polyamides, Starch).

Green Chemistry

Definition: A set of principles used to reduce or eliminate the generation of hazardous substances in the design and manufacture of chemicals.
Automotive Fuel: Development of fuels from renewable resources (sugar/cellulose from plants) to reduce petroleum dependency.
Atmospheric Protection: Focus on reducing Ozone-depleting substances ( $CFCs$ ) and Greenhouse gases ( $CO_2$ ) that cause global warming.

Qualitative Analysis (Ion Identification)

Gas Tests: * $CO_2$ : Turns lime water ( $Ca(OH)_2$ ) cloudy (white precipitate). * $NH_3$ : Pungent smell; turns moist red litmus blue. * $NO_2$ : Brown gas with pungent odor; acidic.
Cation precipitate tests (using NaOH or $NH_3$ ): * $Pb^{2+}$ : White precipitate with NaOH (soluble in excess), white with $NH_3$ (insoluble), yellow precipitate with Potassium Iodide ( $KI$ ). * $Al^{3+}$ : White precipitate (soluble in excess NaOH, insoluble in excess $NH_3$ ). * $Zn^{2+}$ : White precipitate (soluble in both excess NaOH and excess $NH_3$ ). * $Fe^{2+}$ : Green precipitate. * $Fe^{3+}$ : Red-brown precipitate.
Anion tests: * $Cl^-$ : White precipitate with Silver Nitrate ( $AgNO_3$ ) that dissolves in aqueous ammonia. * $I^-$ : Yellow precipitate with Silver Nitrate or Lead Nitrate.