Comprehensive Namibia Grades 10-11 Chemistry Study Guide

Scientific Processes: Mathematical Requirements and Measurements

• Fundamental Operations:

  • Addition: Used to calculate total volumes or masses. Example: Three cylinders contain $50\,\text{cm}^3$, $30\,\text{cm}^3$, and $25\,\text{cm}^3$ of acid. Total volume: $50 + 30 + 25 = 105\,\text{cm}^3$.

  • Subtraction: Used to find differences, such as the mass of a substance in a container. Example: Mass of dish and magnesium oxide is $15.7\,\text{g}$, empty dish is $10.3\,\text{g}$. Substance mass: $15.7 - 10.3 = 5.4\,\text{g}$.

  • Multiplication: Essential for calculating relative molecular mass ($M_r$). Example: Oxygen molecule ($O_2$) mass is $16 \times 2 = 32\,\text{g}$.

  • Division: Used to determine portions. Example: A $12\,\text{cm}$ ribbon divided into four pieces is $12 \div 4 = 3\,\text{cm}$ per piece.

• Averages, Decimals, and Fractions:

  • Average: Calculated by summing items and dividing by the count ($n$). Example: Titrations of $24.8\,\text{cm}^3$, $25.1\,\text{cm}^3$, and $24.6\,\text{cm}^3$ average to $24.6\,\text{cm}^3$.

  • Decimals: Standard precision is often one decimal place. Example: $\frac{98.7 + 98.6 + 98.9}{3} = 98.7333 \approx 98.7\,\text{cm}^3$.

  • Fractions: Written as $\frac{a}{b}$. Calculation examples: $\frac{1}{2} + \frac{1}{3} = \frac{5}{6}$; $\frac{1}{2} \div \frac{3}{4} = \frac{2}{3}$.

  • Percentages: Example: $40\%$ carbon in a $180\,\text{g}$ sample is $\frac{40}{100} \times 180 = 72\,\text{g}$.

  • Ratios: Relationship showing containment. Example: $MgO$ with Relative Atomic Masses of $24$ (Mg) and $16$ (O). Ratio: $24:16 = 3:2$.

• Proportionality and Indices:

  • Direct Proportion: $Y \propto X$ or $Y = kX$. Both variables increase/decrease in the same ratio. Graph is a straight line through the origin.

  • Inverse Proportion: $Y \propto \frac{1}{x}$ or $Y = \frac{k}{X}$. As one variable increases, the other decreases.

  • Indices rules:

  • Multiplying: $a^m \times a^n = a^{m+n}$.

  • Dividing: $a^m \div a^n = a^{m-n}$.

  • Negative index: $a^{-m} = \frac{1}{a^m}$. Example: $3^{-3} = \frac{1}{3^3} = \frac{1}{27}$.

• Scientific Notation and Units:

  • Formatted as $A \times 10^n$ where 1 \le A < 10. Example: $340000 = 3.4 \times 10^5$.

  • Metric Prefixes:

  • Tera ($T = 10^{12}$), Giga ($G = 10^9$), Mega ($M = 10^6$), Kilo ($K = 10^3$), Milli ($m = 10^{-3}$), Micro ($\mu = 10^{-6}$), Nano ($n = 10^{-9}$), Pico ($p = 10^{-12}$).

  • Temperature Scales: Absolute zero is the point where particles stop moving ($0\,\text{K}$ or $-273.15\,^\circ\text{C}$). Conversion: $\text{T/}^\circ\text{C} + 273.15 = \text{T/K}$.

Scientific Skills: Investigations and Data Handling

• Planning an Investigation:

  1. Develop a Hypothesis: A testable statement.

  2. Determine Aim: What the goal is.

  3. Design the Investigation: Ensure validity and reliability.

  4. Variable Determination:

  • Independent Variable: The factor you change.

  • Dependent Variable: The factor you measure.

  • Control Variable: Factors kept constant for a fair test.

  1. Conduct and Record: Use appropriate apparatus and safety gear (gloves, goggles).

  2. Analyse Results: Graphing and tables.

  3. Conclusion: Proving or disproving the hypothesis.

• Recording Data:

  • Tables must have headings with physical quantities and units (e.g., $\text{Time/s}$, $\text{Mass/kg}$).

  • Graphing: Independent variable on the x-axis, Dependent on the y-axis. Use lines of best fit or smooth curves.

• Error, Accuracy, and Precision:

  • Accuracy: How close a measurement is to the true value.

  • Precision: Variations between repeated measurements.

  • Uncertainty: The range within which the true value lies. Example: $11.0 \pm 0.2\,\text{in}$.

  • Anomalous Results: Data points that do not fit the trend. These must be excluded from calculations.

  • Systematic Errors: Caused by faulty instruments or zero errors. Cannot be reduced by averaging.

  • Random Errors: Unpredictable variations. Reduced by repeated readings and averaging.

Experimental Techniques and Purification

• Measurement Apparatus:

  • Time: Stop watch.

  • Temperature: Thermometer.

  • Volume: Measuring cylinder, burette, or pipette.

  • Mass: Beam or electronic balance.

• Purity and Separation Methods:

  • Purity Importance: Essential in pharmaceuticals (safety) and food industry (health).

  • Filtration: Separating an insoluble solid from a liquid (e.g., sand and water).

  • Crystallisation: Separating a soluble solid by evaporating the solvent until crystals form.

  • Re-crystallisation: Purifying an impure solid by dissolving it in hot solvent and cooling slowly.

  • Simple Distillation: Separating liquids with boiling point differences > 50\,^\circ\text{C}.

  • Fractional Distillation: For liquids with boiling point differences < 50\,^\circ\text{C} (e.g., ethanol at $78\,^\circ\text{C}$ and water at $100\,^\circ\text{C}$).

  • Paper Chromatography: Separates substances in solution based on solubility.

  • $R_f\,\text{value} = \frac{\text{distance moved by component}}{\text{distance moved by solvent front}}$.

  • Locating agents are used to visualize colourless spots.

The Particle Nature of Matter

• States of Matter:

  • Solids: Fixed volume/shape, regular pattern, vibrate in fixed positions, strong forces, low kinetic energy.

  • Liquids: Fixed volume, no fixed shape, slide past each other, moderate forces/kinetic energy.

  • Gases: No fixed volume/shape, far apart, fast random motion, weak forces, high kinetic energy.

• Kinetic Particle Theory: Matter is made of tiny, moving particles with spaces and forces between them.

• Changes of State:

  • Melting: Solid to liquid (heat added).

  • Evaporation/Boiling: Liquid to gas (heat added).

  • Sublimation: Solid to gas.

  • Freezing: Liquid to solid (heat removed).

  • Condensation: Gas to liquid (heat removed).

  • Deposition: Gas to solid.

• Diffusion: Movement of particles from high to low concentration. Rate increases with higher temperature and decreases with higher molecular mass.

Atomic Structure and the Periodic Table

• Development of the Atomic Model:

  • John Dalton (1803): Atoms as indestructible spheres.

  • J.J. Thomson (1898): Discovered protons and electrons (Plum Pudding model).

  • Ernest Rutherford (1910): Gold-foil experiment; discovered the positive nucleus.

  • Niels Bohr (1913): Electrons orbit in fixed energy levels (shells).

  • James Chadwick (1932): Discovered the neutron.

• Subatomic Particles:

  • Proton: Charge $+1$, Mass $1\,\text{amu}$, located in nucleus.

  • Neutron: Charge $0$, Mass $1\,\text{amu}$, located in nucleus.

  • Electron: Charge $-1$, Mass $\frac{1}{1840}\,\text{amu}$, located in shells.

• Atomic Notation: ${^A_Z}X$

  • Proton/Atomic Number (Z): Number of protons in the nucleus.

  • Nucleon/Mass Number (A): Sum of protons and neutrons.

• Isotopes: Atoms of the same element with the same proton number but different nucleon numbers (e.g., Carbon-12, Carbon-13, Carbon-14). Radioactive isotopes emit particles to become stable.

  • Uses: $C-14$ (Dating), $Co-60$ (Radiotherapy), $U-235$ (Power generation).

Periodic Table Groups and Trends

• Structure:

  • Groups: Vertical columns (I to VIII); number of valence electrons.

  • Periods: Horizontal rows (1 to 7); number of electron shells.

• Group Properties:

  • Group I (Alkali Metals): Lithium, Sodium, Potassium. Soft, low density, highly reactive with water (forms metal hydroxide + hydrogen). Flame colors: Li (Red), Na (Orange), K (Lilac).

  • Group VII (Halogens): Diatomic non-metals ($F_2, Cl_2, Br_2, I_2$). Reactivity decreases down the group. Darker color down the group. Displace less reactive halides.

  • Group VIII (Noble Gases): Unreactive, full outer shells. Argon used in lamps; Helium in balloons.

  • Transition Elements: High density, high melting points, form coloured compounds, act as catalysts.

• Period 3 Trends: Gradual change from metallic (Na, Mg, Al) to metalloid (Si) to non-metallic (P, S, Cl, Ar). Atomic radius decreases across the period.

Chemical Bonding and Lattices

• Ionic Bonding: Transfer of electrons from metal (cation) to non-metal (anion). Held by electrostatic attraction. Form regular lattices (e.g., $NaCl$). High melting points, non-volatile.

• Covalent Bonding: Sharing of electron pairs between non-metals. Forms molecules (homonuclear like $O_2$ or heteronuclear like $CH_4$).

  • Shapes:

  • Water ($H_2O$): V-shaped.

  • Methane ($CH_4$): Tetrahedral.

  • Ammonia ($NH_3$): Pyramidal.

  • Carbon dioxide ($CO_2$): Linear.

• Giant Covalent Lattices:

  • Diamond: Carbon bonded to four others; extremely hard, non-conductor.

  • Graphite: Carbon bonded to three others in layers; soft lubricant, conducts electricity.

  • Silicon Dioxide ($SiO_2$): Structure similar to diamond; very hard, high melting point.

• Metallic Bonding: Lattice of positive ions in a "sea" of delocalised electrons. Explains high conductivity, malleability, and ductility.

• Alloys: Mixtures of metal with other metals/carbon (e.g., Brass = $Cu + Zn$; Steel = $Fe + C$).

Stoichiometry

• Relative Masses:

  • Relative Atomic Mass ($A_r$): Ratio of average mass of atoms to $\frac{1}{12}$ the mass of Carbon-12.

  • Relative Formula Mass ($M_r$): Sum of $A_r$ of all atoms in a formula.

• The Mole Concept:

  • $1\,\text{mole} = 6.022 \times 10^{23}\,\text{particles}$ (Avogadro Constant).

  • $n = \frac{m}{M_r}$ where $n$ is moles and $m$ is mass in grams.

  • Molar Gas Volume: $24\,\text{dm}^3$ at RTP ($24000\,\text{cm}^3$).

• Concentration:

  • $C = \frac{n}{V}$ (units: $\text{mol/dm}^3$).

  • $1\,\text{dm}^3 = 1000\,\text{cm}^3$.

• Analysis and Formulas:

  • Empirical Formula: Simplest whole-number ratio of elements.

  • Percentage Yield: $\frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$.

  • Percentage Purity: $\frac{\text{mass of pure substance}}{\text{mass of impure substance}} \times 100\%$.

Electrochemistry

• Terminology:

  • Electrolyte: Substance that conducts and decomposes via electricity.

  • Anode: Positive electrode; oxidation takes place.

  • Cathode: Negative electrode; reduction takes place.

• Electrolysis Examples:

  • Molten Lead(II) Bromide ($PbBr_2$): Pb at cathode, $Br_2$ at anode.

  • Concentrated Aqueous NaCl: $H_2$ at cathode, $Cl_2$ at anode, $NaOH$ remains in solution.

• Electroplating: Using electricity to coat a substrate with a metal (e.g., copper).

  • Object to be plated is the Cathode.

  • Coating metal is the Anode.

Chemical Reactions and Energetics

• Energetics:

  • Exothermic: Heat released to surroundings ($-\Delta H$). Bond forming is exothermic.

  • Endothermic: Heat absorbed ($+\Delta H$). Bond breaking is endothermic.

  • Enthalpy Change Calculation: $q = mc\Delta T$.

• Rate of Reaction:

  • Increases with: Concentration, Pressure, Temperature, Surface Area, and Catalysts.

  • Catalyst: Lowers activation energy without being consumed.

  • Inhibitor: Reduces the rate of reaction.

• Equilibrium: Occurs in closed systems for reversible reactions when Forward Rate = Backward Rate.

  • Le Chatelier's Principle: Systems shift to counteract changes in temperature, pressure, or concentration.

• Redox:

  • Oxidation: Gain of oxygen, loss of electrons (OIL).

  • Reduction: Loss of oxygen, gain of electrons (RIG).

Metals: Properties and Extraction

• Reactivity Series: Order: K, Na, Ca, Mg, Al, (C), Zn, Fe, Sn, Pb, (H), Cu, Ag, Au, Pt.

• Iron Extraction (Blast Furnace):

  • Raw materials: Haematite ($Fe_2O_3$), Coke (C), Limestone ($CaCO_3$), Hot air.

  • Reduction: $Fe_2O_3(s) + 3CO(s) \rightarrow 2Fe(l) + 3CO_2(g)$.

  • Slag formation: $CaO(s) + SiO_2(s) \rightarrow CaSiO_3(l)$.

• Aluminium Extraction: By electrolysis of Bauxite dissolved in molten cryolite (to lower melting point).

• Zinc Extraction: Roasting Zinc Blende ($ZnS$) then reduction with carbon.

• Rust Prevention: Painting, galvanizing (Zinc), tin-plating, and sacrificial protection.

Organic Chemistry

• Naming: Meth- ($1C$), Eth- ($2C$), Prop- ($3C$), But- ($4C$).

• Functional Groups:

  • Alkanes: $-ane$; saturated; single bonds; $C_nH_{2n+2}$.

  • Alkenes: $-ene$; unsaturated; double bonds; $C_nH_{2n}$; test with bromine water (decolorizes).

  • Alcohols: $-ol$; hydroxyl group ($-OH$); $C_nH_{2n+1}OH$.

  • Carboxylic Acids: $-oic \, \text{acid}$; group ($-COOH$).

• Fractional Distillation of Petroleum: Order: Refinery Gas, Gasoline (Petrol), Naphtha, Kerosene (Jet fuel), Diesel, Fuel oil, Lubricants, Bitumen.

• Polymers:

  • Addition: Breaking double bonds (e.g., Polyethene).

  • Condensation: Linking units with loss of water (e.g., Nylon, Terylene).

  • Natural Macromolecules: Proteins (amide Link), Fats (ester Link), Carbohydrates (sugar units).

Environmental Chemistry

• Water:

  • Chemical Test: Anhydrous $CuSO_4$ (white to blue) or $CoCl_2$ (blue to pink).

  • Hardness: Caused by $Ca^{2+}$ and $Mg^{2+}$.

  • Temporary: Hydrogen carbonates (remove by boiling).

  • Permanent: Sulfates (remove by washing soda or ion exchange).

  • Purification: Filtration, Sedimentation, Chlorination.

• Air:

  • Composition: $78\% \, N_2, 21\% \, O_2, 1\% \, \text{other gases}$.

  • Pollutants:

  • $CO$: Toxic, blocks haemoglobin.

  • $SO_2$ and $NO_x$: Cause acid rain.

  • $CH_4$ and $CO_2$: Greenhouse gases; global warming.

• Industrial Processes:

  • Haber Process: Ammonia production; $450\,^\circ\text{C}$, $200\,\text{atm}$, Iron catalyst.

  • Contact Process: Sulfuric acid; $V_2O_5$ catalyst, $450\,^\circ\text{C}$, $2\,\text{atm}$.