Namibia NSSCO Chemistry Syllabus (2018) - Ordinary Level (Grade 10-11)

Introduction

• Namibia Senior Secondary Certificate Ordinary (NSSCO) level syllabus designed as a two-year course leading to examination after the Junior Secondary phase.
• Syllabus aligns with the National Curriculum for Basic Education (NCBE) and approved by the National Examination, Assessment and Certification Board (NEACB).
• Namibia National Curriculum Guidelines emphasize learner-centred education, recognising learner diversity and multilingual/multicultural contexts. Core aims include developing values, attitudes, knowledge and skills; promoting self-awareness and understanding of others; respecting human rights and free speech; engaging with global issues (AIDS pandemic, global warming, environmental concerns, wealth distribution, conflicts, technology and connectivity).
• Skills highlighted by the guidelines (with some marked as relevant to this syllabus):- communication skills*, numeracy skills*, information skills*, problem-solving skills*, self-management and competitive skills*, social and cooperative skills, physical skills, work and study skills*, critical and creative thinking*.
• NSSCO Chemistry intended to develop essential skills across disciplines; skills cannot be developed in isolation.

Rationale

• The syllabus describes intended learning and assessment for Chemistry in the NSSCO phase; Chemistry is within the natural sciences with thematic links to other subjects.
• Emphasis on understanding the physical and biological world locally, regionally and internationally, including resource use and ecologically sustainable changes.
• Focus on the application of scientific knowledge and attitudes to health, environment, and resource sustainability.
• Critical thinking, investigation, data interpretation, and application of knowledge to practical (experimental/investigative) skills are central.
• Use of modern technology and equipment to solve problems through planning, design, realization and evaluation of activities.

Aims

• The aims are the same for all learners and describe educational purposes of Chemistry for the NSSCO examination (not listed in priority):
  1) Provide a worthwhile educational experience through theoretical and practical science to enable learners to become confident citizens in a technological world, take an informed interest in scientific matters, recognise usefulness and limitations of the scientific method, and be prepared for further study in pure, applied sciences or science-related vocational courses.
  2) Develop abilities and skills relevant to Chemistry and useful in daily life; promote safe practice and effective communication.
  3) Develop attitudes such as accuracy, objectivity, integrity, enquiry, initiative and inventiveness relevant to Chemistry.
  4) Stimulate interest in and care for the environment.
  5) Promote awareness that scientific theories/methods evolve through cooperative activity; science is influenced by societal, economic, technological, ethical and cultural factors; science applications can be beneficial or detrimental; science language is universal across borders.

Additional information

• Guided learning hours: approximately 130 hours per subject over two years (guidance). NCBE indicates 8 periods of 40 minutes per 7-day cycle or 6 periods per 5-day cycle.
• Prior learning: recommended that learners have studied Physical Science at Junior Secondary level.
• progression: NSSCO grades (A–G); grades C to A* prepare for NSCCAS level Chemistry; (Ungraded indicates fail to meet minimum pass standard).
• Support materials: syllabuses, question papers, examiner reports, and approved learning support materials are distributed; Senior Secondary Textbook Catalogue available on the NIED website.

Learning content

The course content is organized into topics and sub-topics:
1) Topic 1: Scientific processes

• 1.1 Mathematical requirements
• 1.2 Scientific skills
  • 1.2.1 Planning and conducting investigations
  • 1.2.2 Recording data
• 1.3 Basic units and derived units
• 1.4 Error, accuracy and uncertainty
• 1.5 Experimental techniques
  Practical activities (Topic 1): crystallisation of a salt solution; distillation of fermented fruits/juice; filtration examples (muddy water, insoluble salt); investigate components of coloured mixtures (food colouring, inks).

2) Topic 2: Matter

2.1 Atoms, Elements, Molecules and Compounds

2.1.1 Particle nature of matter

• Matter: Anything that has mass and occupies space.
• States of Matter:
  • Solid: Particles are closely packed in fixed positions, vibrate but do not move past each other. Has a definite shape and volume.
  • Example: Ice, rock.
  • Liquid: Particles are closely packed but can slide past each other. Has a definite volume but takes the shape of its container.
  • Example: Water, oil.
  • Gas: Particles are far apart and move randomly and rapidly. No definite shape or volume, fills container.
  • Example: Oxygen, steam.
• Kinetic Particle Theory: States that all matter consists of tiny particles that are in constant, random motion. The energy of these particles is related to temperature.
  • Example: Heating water increases the kinetic energy of its particles, leading to evaporation.

2.1.2 Atomic structure

• Atom: The smallest unit of an element that retains the chemical identity of that element. Atoms consist of a central nucleus and electrons orbiting the nucleus.

• Subatomic Particles:

  • Proton (p^+): Positively charged particle, found in the nucleus. Relative mass \approx 1 .
  • Neutron (n_0): Neutrally charged particle, found in the nucleus. Relative mass \approx 1 .
  • Electron (e^-): Negatively charged particle, found orbiting the nucleus in electron shells. Relative mass \approx \frac{1}{1836} (negligible).

• Nucleus: The dense, central part of an atom containing protons and neutrons.

• Electron Shells/Energy Levels: Regions around the nucleus where electrons are found. Electrons fill these shells starting from the innermost shell (2, 8, 8 rule for first 3 shells).

• Atomic Number (Z): The number of protons in an atom's nucleus. It determines the element's identity.

• Mass Number (A): The total number of protons and neutrons in an atom's nucleus.

  • Formula: Mass Number = Number of protons + Number of neutrons.
  • Representation: ^A_Z X, where X is the element symbol.
  • Example: ^{12}_6\mathrm{C} has 6 protons and 12-6=6 neutrons.

• Element: A pure substance consisting of only one type of atom (i.e., all atoms have the same atomic number).

  • Example: Oxygen (O), Hydrogen (H), Gold (Au).

• Molecule: A group of two or more atoms chemically bonded together.

  • Example: Water (\mathrm{H2O}), Oxygen gas (\mathrm{O2}), Carbon dioxide (\mathrm{CO_2}).

• Compound: A substance formed when two or more different elements are chemically bonded together in a fixed ratio.

  • Example: Water (\mathrm{H_2O} is a compound as it contains H and O atoms), Sodium chloride (NaCl).
  • Distinction: All compounds are molecules, but not all molecules are compounds (e.g., \mathrm{O_2} is a molecule but an element, not a compound).

2.2 Isotopes

• Isotopes: Atoms of the same element (same number of protons/atomic number) but with different numbers of neutrons (and thus different mass numbers).
  • Example:
  • Carbon-12 (^{12}_6\mathrm{C}): 6 protons, 6 neutrons.
  • Carbon-14 (^{14}_6\mathrm{C}): 6 protons, 8 neutrons. Carbon-14 is radioactive and used in carbon dating.
  • Hydrogen: Protium (^11\mathrm{H}), Deuterium (^21\mathrm{H}), Tritium (^3_1\mathrm{H}).

2.3 Periodic Table

• Periodic Table: A table that arranges elements by increasing atomic number, displaying their recurring chemical properties.
• Groups (Vertical Columns): Elements in the same group have the same number of valence electrons (outermost electrons) and thus exhibit similar chemical properties.
  • Example: Group 1 (Alkali Metals) are highly reactive metals with 1 valence electron. Group 17 (Halogens) are very reactive non-metals with 7 valence electrons.
• Periods (Horizontal Rows): Elements in the same period have the same number of electron shells. As you move across a period, atomic number increases, and properties change gradually from metallic to non-metallic.
• Periodicity: The recurrence of similar chemical properties at regular intervals when elements are arranged by increasing atomic number. This is explained by the periodic recurrence of electron configurations.

2.4 Bonding: the structure of matter

2.4.1 Building blocks of matter

• Ions: Atoms that have gained or lost electrons, resulting in a net electrical charge.
  • Cation: Positively charged ion (lost electrons). Example: Na^+
  • Anion: Negatively charged ion (gained electrons). Example: Cl^-

2.4.2 Ionic bonding / electrovalent bonds

• Ionic Bond: A strong electrostatic force of attraction between oppositely charged ions, formed by the transfer of electrons from a metal atom to a non-metal atom.
  • Formation: Metal loses electrons to form a positive ion (cation), while a non-metal gains electrons to form a negative ion (anion).
  • Properties: Typically form crystalline solids with high melting and boiling points, are soluble in water, and conduct electricity when molten or in aqueous solution (ions are free to move).
  • Example: Sodium chloride (NaCl). Sodium (Group 1) loses 1 electron to become Na^+; Chlorine (Group 17) gains 1 electron to become Cl^- formed by Na^+ + Cl^- \rightarrow NaCl.

2.4.3 Molecules and covalent bonds

• Covalent Bond: A chemical bond formed by the sharing of one or more pairs of electrons between two non-metal atoms.
  • Formation: Each atom contributes an electron to the shared pair, allowing both atoms to achieve a stable electron configuration (e.g., a full outer shell).
  • Types: Single, double, or triple covalent bonds.
  • Properties of Simple Molecular Structures: Low melting and boiling points (weak intermolecular forces), do not conduct electricity.
  • Example: Water (\mathrm{H2O}), Methane (\mathrm{CH4}), Oxygen gas (\mathrm{O2}), Carbon dioxide (\mathrm{CO2}).

2.4.4 Giant covalent structures

• Giant Covalent Structure (Macromolecular Structure): A substance where all atoms are held together by strong covalent bonds in a vast, continuous network.
  • Properties: Very high melting and boiling points (strong covalent bonds must be broken), usually insoluble in common solvents, and typically do not conduct electricity (except graphite).
  • Examples:
  • Diamond: Each carbon atom is bonded to four other carbon atoms in a tetrahedral arrangement. Extremely hard.
  • Graphite: Each carbon atom is bonded to three other carbon atoms in hexagonal layers. Layers are held by weak intermolecular forces, allowing them to slide (soft, slippery). Conducts electricity due to delocalized electrons within layers.
  • Silicon Dioxide (\mathrm{SiO_2}): Similar structure to diamond, each silicon atom is bonded to four oxygen atoms, and each oxygen to two silicon atoms.

2.4.5 Metallic bonding

• Metallic Bond: The electrostatic force of attraction between positively charged metal ions (cations) and a "sea" of delocalized electrons.
  • Formation: Metal atoms lose their outermost electrons, which become delocalized (free to move throughout the structure) forming a lattice of positive ions.
  • Properties: High melting and boiling points (strong metallic bonds), malleable (can be hammered into shape) and ductile (can be drawn into wires) because the layers of ions can slide past each other without breaking the bond, good conductors of electricity and heat (due to mobile electrons), lustrous (shiny).
  • Example: Copper, Iron, Aluminium.

2.4.6 Writing and balancing equations

• Chemical Equation: A symbolic representation of a chemical reaction, showing the reactants on the left and products on the right, separated by an arrow.
  • Law of Conservation of Mass: In a chemical reaction, mass is neither created nor destroyed. The number of atoms of each element must be the same on both sides of the equation.
• Balancing Equations: Adjusting the stoichiometric coefficients (numbers in front of chemical formulas) to ensure that the number of atoms of each element is equal on both sides of the equation.
  • Steps:
  1. Write the unbalanced equation with correct chemical formulas.
  2. Count the number of atoms of each element on both sides.
  3. Add coefficients to balance one element at a time (metals first, then non-metals, then hydrogen, then oxygen often last). Do not change subscripts.
  4. Verify that all atoms are balanced.
  • Example 1 (Formation of water):
  • Unbalanced: \mathrm{H2(g) + O2(g) \rightarrow H_2O(l)}
  • Balanced: \mathrm{2H2(g) + O2(g) \rightarrow 2H_2O(l)}
  • Example 2 (Combustion of Methane):
  • Unbalanced: \mathrm{CH4(g) + O2(g) \rightarrow CO2(g) + H2O(l)}
  • Balanced: \mathrm{CH4(g) + 2O2(g) \rightarrow CO2(g) + 2H2O(l)}

Exercise for Topic 2: Matter

1. Classify the following as an element, compound, or mixture:
   • a) Air
   • b) Water (\mathrm{H_2O})
   • c) Gold (Au)
   • d) Carbon dioxide (\mathrm{CO_2})
   • e) Oxygen gas (\mathrm{O_2})
2. An atom has 17 protons, 18 neutrons, and 17 electrons.
   • a) What is its atomic number?
   • b) What is its mass number?
   • c) What element is this atom?
   • d) Is this a neutral atom or an ion? Explain.
3. Name the type of bonding present in:
   • a) Potassium iodide (KI)
   • b) Methane (\mathrm{CH_4})
   • c) Copper (Cu)
   • d) Diamond (C)
4. Describe one key difference in properties between simple molecular structures (like water) and giant covalent structures (like diamond) and explain why using bonding principles.
5. Balance the following chemical equations:
   • a) \mathrm{N2(g) + H2(g) \rightarrow NH_3(g)}
   • b) \mathrm{Al(s) + O2(g) \rightarrow Al2O_3(s)}
   • c) \mathrm{C3H8(g) + 5O2(g) \rightarrow CO2(g) + H_2O(l)}

Answers for Topic 2: Matter Exercise

1. a) Air - Mixture (of gases like N2, O2, Ar, CO2)
   b) Water (\mathrm{H2O}) - Compound c) Gold (Au) - Element d) Carbon dioxide (\mathrm{CO2}) - Compound
   e) Oxygen gas (\mathrm{O_2}) - Element (contains only oxygen atoms, but it is a molecule)
2. a) Atomic number = 17 (number of protons)
   b) Mass number = 17 (protons) + 18 (neutrons) = 35
   c) The element is Chlorine (Cl), as its atomic number is 17.
   d) This is a neutral atom because the number of protons (17) is equal to the number of electrons (17), resulting in no net charge.
3. a) Potassium iodide (KI) - Ionic bonding (metal K and non-metal I)
   b) Methane (\mathrm{CH_4}) - Covalent bonding (non-metals C and H)
   c) Copper (Cu) - Metallic bonding
   d) Diamond (C) - Covalent bonding (specifically, a Giant Covalent Structure)
4. One key difference is their melting and boiling points. Simple molecular structures, like water, have low melting and boiling points because the forces between molecules (intermolecular forces) are weak and require little energy to overcome. Inside the molecule, the covalent bonds are strong, but these are not broken during melting or boiling. Giant covalent structures, like diamond, have very high melting and boiling points because all atoms are held together by a vast network of strong covalent bonds which require a large amount of energy to break throughout the entire structure.
5. a) \mathrm{N2(g) + 3H2(g) \rightarrow 2NH3(g)} b) \mathrm{4Al(s) + 3O2(g) \rightarrow 2Al2O3(s)}
   c) \mathrm{C3H8(g) + 5O2(g) \rightarrow 3CO2(g) + 4H_2O(l)}

3) Topic 3: Materials

• 3.1 Types of materials
• 3.2 Building materials
• 3.3 Cleaning materials
• 3.4 Nano-materials
  Practical activities: building material strength tests; surveys; insulating materials experiments; emulsifying effects of soaps/detergents; soap synthesis via hydrolysis.

4) Topic 4: Stoichiometry

• 4.1 The mole concept
• 4.2 Mole calculations
• Includes empirical formula, relative formula mass, Avogadro constant, and concentration units (g/dm^3 and mol/dm^3).

5) Topic 5: Electrochemistry

• 5.1 Electrodes and electrolytes
• Includes electrolysis principles, electroplating, electrode definitions, and electrode reactions in molten/aqueous systems; products of electrolysis in specified cells (e.g., molten PbBr2, sulfuric acid, HCl, NaCl).
  Practical: conductivity of salt solution and lemon juice; electroplating zinc with copper sulfate.

6) Topic 6: Chemical reactions

• 6.1 Chemical and physical changes
• 6.2 Energetics of a reaction (enthalpy; exothermic/endothermic; q = mcΔT; bond energies)
• 6.3 Production of energy (heat from fuels; nuclear isotopes as energy sources; fuel cells and solar cells)
• 6.4 Rate of reaction (factors: concentration, temperature, surface area, catalysts; activation energy; enzyme catalysis; experimental investigation of rate; safety considerations in powder handling)
• 6.5 Reversible reactions (dynamic equilibrium; effect of condition changes; closed systems)
• 6.6 Redox (oxidation/reduction; identifying redox reactions; reducing/oxidising agents)
  Practical activities: energy from burning methylated spirits; rate experiments; decomposition reactions; light-driven decomposition.

7) Topic 7: Acids, bases and salts

• 7.1 Properties and definitions; proton transfer in aqueous solutions; pH concepts; differences between alkalis, bases; weak vs strong acids/bases; neutralisation; H+ + OH− → H2O
• 7.2 Types of oxides (basic, acidic, amphoteric)
• 7.3 Neutralisation: reactions with oxides, hydroxides, carbonates; test for evolving gas; preparation of soluble salts via neutralisation
• 7.4 Preparation of salts: solubility rules; methods to prepare soluble/insoluble salts; common salts solubility rules (Na+, K+, NH4+ soluble; nitrates soluble; lead compounds often insoluble; hydroxides generally insoluble; carbonates insoluble except Na+, K+, NH4+; chlorides/ sulfates generally soluble with exceptions)
  Practical: pH testing; salt crystallisation; acid-base reactions; titration; salt preparation experiments.

8) Topic 8: Qualitative analysis

• 8.1 Identification of ions in solution (cations: NH4+, Al3+, Zn2+, Ca2+, Cu2+, Fe2+/Fe3+; anions: carbonate, chloride, bromide, iodide, nitrate, sulfate)
• 8.2 Identification of gases (NH3, CO2, Cl2, H2, O2)
  Practical: ion and gas tests; use of annexe B (qualitative analyses notes not supplied in Paper 3).

9) Topic 9: Metals

• 9.1 Properties of metals
• 9.2 Reactivity series
• 9.3 Extraction of metals (Al, Cu, Zn, Fe) from ores; Namibia-specific ores and uses; methods of extraction including electrolysis and reduction with carbon; native metals; rust prevention; steel types (basic oxygen steelmaking, electric arc furnace) and ore resources; recycling considerations
  Practical: metal reactivity tests; displacement reactions; corrosion tests.

10) Topic 10: Organic chemistry

• 10.1 Names of compounds (IUPAC; functional groups) up to butane family; identification of functional groups by name
• 10.2 Hydrocarbons: Fractional distillation of petroleum; Fractions and uses; Homologous series and isomerism; structural isomers for C4–C5; isomer examples
• 10.3 Alkanes: general formula CnH2n+2; properties; substitution reactions with Cl2
• 10.4 Alkenes: general formula CnH2n; addition reactions; catalytic cracking; tests to distinguish saturated vs unsaturated hydrocarbons
• 10.5 Alcohols: OH group; fermentation and steam addition to ethene; uses of ethanol
• 10.6 Carboxylic acids: COOH group; oxidation of ethanol; fermentation
• 10.7 Esters: formation from carboxylic acids and alcohols; uses in flavours/fragrances
• 10.8 Polymers: synthetic polymers; addition vs condensation polymerisation; natural macromolecules; materials like nylon, Terylene; uses of plastics
  Practical: distillation of wine/juice; soap preparation; tests for saturation; fibre properties.

11) Topic 11: Environmental and industrial chemistry

• 11.1 Water: tests for water; tests for purity; hard/soft water; temporary vs permanent hardness; softening methods; water availability in Namibia; fertilisers impact; sewage and landfill effects; purification of water and safe supply considerations
• 11.2 The air around us: composition of dry air; CO2 formation from combustion/respiration/fermentation; commercial preparation/use of O2, N2, CO2
• 11.3 Gas preparation and use: oxygen, nitrogen, carbon dioxide uses
• 11.4 Pollution of the air: carbon monoxide dangers; acidic oxides; acid rain; vehicle exhausts; greenhouse gases (CO2, methane); lead pollution; role of catalysts; air pollutants as global concerns
• 11.5 Chemical industrial plants: lime/lime stone; ammonia (Haber process) conditions; sulfuric acid (Contact process); other industrial plants; salt extraction; pollution control and environmental considerations

Assessment objectives (A, B, C)

• Three assessment objectives in Chemistry:-
  • A: Knowledge with understanding
  • A1 scientific phenomena, facts, laws, definitions, concepts and theories
  • A2 scientific vocabulary, terminology and conventions (symbols, quantities, units)
  • A3 scientific instruments and apparatus, techniques, safety aspects
  • A4 scientific quantities and their determination
  • A5 scientific and technological applications with social, economic and environmental implications
  • B: Handling information, application and solving problems
  • B1 locate, select, organize information from sources
  • B2 translate information from one form to another
  • B3 manipulate numerical and other data
  • B4 identify patterns/trends and report inferences
  • B5 present explanations for phenomena, patterns and relationships
  • B6 form predictions/hypotheses
  • B7 solve quantitative/qualitative problems related to everyday life
  • C: Practical (experimental and investigative) skills and abilities
  • C1 safe use of techniques/apparatus/materials
  • C2 plan experiments/investigations
  • C3 make/record observations and measurements
  • C4 interpret/evaluate observations and data, handle anomalous results
  • C5 evaluate methods and suggest improvements

Scheme of assessment

• Papers are compulsory: Papers 1, 2 and 3.
  • Paper 1: Theory – Multiple choice questions; 40 items; 45 minutes; 40 marks; assesses A and B.
  • Paper 2: Theory – Structured questions; 1 hour 30 minutes; 80 marks; assesses A and B.
  • Paper 3: Alternative to practical – Assessment of practical skills; 1 hour 15 minutes; 40 marks; assesses C; includes questions on practical techniques; notes for qualitative analysis (Annexe B) not supplied for Paper 3.
• Total: 160 marks across papers 1–3.
• All learners sit Papers 1–3.

Weighting and specification grid

• Overall weighting across papers:
  • A (Knowledge with understanding): 50%
  • B (Handling information, application and solving problems): 30%
  • C (Practical skills): 20%
• Paper-specific weightings:
  • Paper 1: A 25 marks; B 50 marks; C 0 marks
  • Paper 2: A 50 marks; B 30 marks; C 0 marks
  • Paper 3: A 0 marks; B 0 marks; C 40 marks
• Total across all papers: A 50% (80 marks); B 30% (48 marks); C 40% (40 marks) – note total sums to 160 marks with distribution per paper as above.

Grade descriptions

• A grade: Recall wide range of knowledge; apply detailed scientific knowledge in diverse contexts; use advanced scientific vocabulary; explain how theories may change with new evidence; synthesize information from multiple sources; solve multi-variable problems; analyze data and identify patterns; generate and test hypotheses.
• C grade: Recall a range of scientific information; apply knowledge in general contexts; use appropriate terminology; explain how theories may be modified by new evidence; select and present information clearly; identify patterns; solve multi-step problems with a range of variables; generate hypotheses.
• F grade: Recall limited knowledge; apply limited facts to familiar phenomena; communicate simple scientific ideas with limited terminology; select a single piece of information; solve multi-step problems only with guidance; analyze data for patterns with help; evaluate basic techniques.

Glossary of terms (selected definitions)

• Define: a formal statement or equivalent paraphrase; provide a definition and significance/context.
• State: concise answer, often numerical or exact; minimal justification.
• List: provide a number of points without elaboration.
• Explain: provide reasoning or theory; relate to phenomena or experiments.
• Describe: state main points with explanatory detail; reference observations.
• Describe and explain: combine description with explanation.
• Discuss: present a critical account, including considerations and evaluation.
• Outline: give essentials succinctly.
• Deduce/Predict: infer from given information; apply logic rather than recall.
• Construct: write a balanced equation or construct a model/diagram as required.
• Compare: identify similarities and differences.
• Classify: group things by shared characteristics.

Annexes

• Annexe A: Assessment criteria for Paper 3 (Alternative to Practical)-
  • Experimental contexts may include: simple quantitative experiments; reaction rates; temperature measurements; practical investigatory problems; filtration; electrolysis; ion/gas identification.
  • Experimental skills/investigations requirements include: taking readings with proper accuracy/precision; describing experimental arrangements; completing data tables; drawing/applying apparatus diagrams; drawing conclusions linked to data; interpreting observations; plotting/interpreting graphs; identifying errors and suggesting improvements; planning experiments with predictions and apparatus choices; selecting appropriate apparatus; commenting on precautions.
• Annexe B: Notes for use in qualitative analysis (tests for ions/gases)-
  • Tests for ions in solution: cations (NH4+, Al3+, Zn2+, Ca2+, Cu2+, Fe2+, Fe3+) with NaOH/NaNH3; anions (CO3^2−, Cl−, Br−, I−, NO3−, SO4^2−) with indicated procedures; flame tests for Li+, Na+, K+, Ca2+, Ba2+.
  • Gases: NH3, CO2, Cl2, H2, O2 with corresponding test observations.
• Annexe C: Data sheet – the Periodic Table of the elements-
  • A simplified Periodic Table with Groups I–VIII and blocks; example entries (H, He, Li, Be, etc.) and key notes such as the volume of one mole of a gas is 24 dm^3 at room temperature and pressure (rtp).
• Annexe D: Units of physical quantities-
  • Lists common units and symbols plus instructions on how to present units (e.g., metres per second as m/s or m s^-1); examples include length (km, m, cm, mm), mass (tonne Mg, kg, g, mg), weight (N), time (y, d, h, min, s), amount of substance (mol), area (m^2, dm^2, cm^2), volume (km^3, m^3, dm^3), density (kg/m^3, g/cm^3).

Key equations and numerical references (LaTeX)

• Avogadro constant and mole concept:
  N*A = 6.022 \times 10^{23} 1 \text{ mole} = N*A \text{ particles}
  \text{Mr} = \sum (\text{relative atomic masses})
  \text{Ar} = \frac{\text{average mass of atoms of element}}{\tfrac{1}{12} \text{ mass of } ^{12}\mathrm{C}}
• Concentration and solution quantities:
  c = \frac{n}{V} \quad(\text{mol/dm}^3)
  \text{concentration (g/dm}^3) = \frac{\text{mass (g)}}{\text{volume (dm}^3)}
• Molecular formula relationships:
  \text{Empirical formula} = \text{simplest whole-number ratio}
  \text{Molecular formula} = n \times \text{Empirical formula (when } n = \frac{\text{Mr}}{\text{Empirical formula mass}}\text{)}
• Energy and calorimetry (simple):
  $$q =