What you'll learn
This topic forms the foundation of Edexcel GCSC Chemistry, covering atomic structure, the periodic table, chemical bonding, and quantitative chemistry. Understanding these principles is essential for every other topic you'll encounter, and questions on these concepts appear across all three exam papers. Mastering this content ensures you can write formulae, balance equations, and understand why substances behave the way they do.
Key terms and definitions
Atom — the smallest particle of an element that retains its chemical properties, consisting of a nucleus containing protons and neutrons surrounded by electrons in shells.
Element — a substance made of only one type of atom, which cannot be broken down into simpler substances by chemical means.
Compound — a substance formed when two or more elements chemically combine in fixed proportions, with properties different from the constituent elements.
Ion — an atom or group of atoms that has gained or lost electrons, resulting in a positive or negative electrical charge.
Isotope — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.
Relative atomic mass (Ar) — the average mass of atoms of an element compared to 1/12th the mass of a carbon-12 atom, taking into account the abundance of isotopes.
Covalent bond — a chemical bond formed when two non-metal atoms share a pair of electrons.
Ionic bond — a chemical bond formed by the electrostatic attraction between oppositely charged ions, typically between a metal and non-metal.
Core concepts
Atomic structure and the periodic table
Atoms consist of a central nucleus containing protons (positive charge, relative mass 1) and neutrons (no charge, relative mass 1), surrounded by electrons (negative charge, relative mass 1/1836) arranged in shells. The atomic number equals the number of protons and determines the element's identity. The mass number equals the total number of protons plus neutrons.
Electron configuration follows specific rules:
- First shell holds maximum 2 electrons
- Second shell holds maximum 8 electrons
- Third shell holds maximum 8 electrons (at GCSE level)
- Electrons fill shells from the innermost outward
For example, sodium (atomic number 11) has electron configuration 2,8,1.
The periodic table arranges elements by atomic number. Elements in the same group (vertical columns) have the same number of outer shell electrons, giving them similar chemical properties. The period number (horizontal rows) indicates the number of electron shells an atom possesses.
Group 1 elements (alkali metals) have one outer electron and become increasingly reactive down the group. Group 7 elements (halogens) have seven outer electrons and become less reactive down the group. Group 0 elements (noble gases) have full outer shells, making them unreactive.
Chemical bonding
Ionic bonding occurs when metals transfer electrons to non-metals. Metal atoms lose electrons to form positive cations, while non-metal atoms gain electrons to form negative anions. The resulting ions have full outer shells and are held together by strong electrostatic forces.
For example, when sodium reacts with chlorine:
- Sodium (2,8,1) loses one electron to form Na⁺ (2,8)
- Chlorine (2,8,7) gains one electron to form Cl⁻ (2,8,8)
- These ions form sodium chloride with the formula NaCl
Ionic compounds form giant ionic lattices with:
- High melting and boiling points due to strong electrostatic forces
- Conductivity when molten or dissolved (ions free to move)
- No conductivity when solid (ions fixed in position)
Covalent bonding occurs when non-metal atoms share pairs of electrons to achieve full outer shells. Each shared pair constitutes one covalent bond.
Simple molecular substances contain covalent bonds within molecules but weak intermolecular forces between molecules, resulting in:
- Low melting and boiling points
- No electrical conductivity (no free electrons or ions)
Examples include water (H₂O), carbon dioxide (CO₂), and methane (CH₄).
Giant covalent structures have every atom bonded covalently to others in a continuous network:
- Diamond: each carbon atom bonded to four others in a tetrahedral structure, making it extremely hard with a very high melting point
- Graphite: each carbon atom bonded to three others in layers, with delocalised electrons between layers allowing electrical conductivity; layers can slide over each other making it soft and slippery
- Silicon dioxide: each silicon atom bonded to four oxygen atoms, creating a hard substance with a very high melting point
Metallic bonding consists of positive metal ions arranged in a regular lattice surrounded by delocalised electrons. This structure explains metallic properties:
- Good electrical and thermal conductivity (delocalised electrons can move)
- Malleability and ductility (layers can slide over each other while maintaining bonding)
- High melting and boiling points (strong metallic bonds)
Chemical formulae and equations
Chemical formulae show the number and type of atoms in a substance. For ionic compounds, the formula reflects the ratio needed to balance charges:
- Magnesium oxide: Mg²⁺ and O²⁻ combine in a 1:1 ratio → MgO
- Calcium chloride: Ca²⁺ and Cl⁻ combine in a 1:2 ratio → CaCl₂
- Aluminium oxide: Al³⁺ and O²⁻ combine in a 2:3 ratio → Al₂O₃
Balancing equations ensures the same number of each type of atom appears on both sides. The process involves:
- Write the correct formulae for all reactants and products
- Count atoms of each element on both sides
- Add numbers in front of formulae (coefficients) to balance
- Never change subscript numbers within formulae
For example: magnesium + oxygen → magnesium oxide
- Unbalanced: Mg + O₂ → MgO
- Balanced: 2Mg + O₂ → 2MgO
State symbols indicate physical states: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution.
Relative formula mass and moles
Relative formula mass (Mr) is calculated by adding the relative atomic masses of all atoms in a formula.
For calcium carbonate (CaCO₃):
- Ca: 40
- C: 12
- O: 16 × 3 = 48
- Mr = 40 + 12 + 48 = 100
The mole is the unit for amount of substance. One mole contains 6.02 × 10²³ particles (Avogadro's constant). The mass of one mole equals the relative formula mass in grams.
Key calculations:
Number of moles = mass (g) ÷ Mr
Mass (g) = number of moles × Mr
Conservation of mass states that mass cannot be created or destroyed in chemical reactions. If mass appears to change:
- Mass increase: a gas from the air has reacted with the substance (e.g., magnesium burning in oxygen)
- Mass decrease: a gas has escaped from the reaction mixture (e.g., thermal decomposition of carbonates releasing CO₂)
Concentration and yields
Concentration measures the amount of solute dissolved in a given volume of solution.
Concentration (g/dm³) = mass of solute (g) ÷ volume (dm³)
To convert cm³ to dm³, divide by 1000.
For molar concentration: Concentration (mol/dm³) = moles of solute ÷ volume (dm³)
Percentage yield compares actual yield to theoretical yield:
Percentage yield = (actual yield ÷ theoretical yield) × 100
Yields less than 100% occur due to:
- Incomplete reactions
- Side reactions producing different products
- Loss of product during separation and purification
- Reversible reactions not going to completion
Worked examples
Example 1: Electron configuration and bonding
Question: Calcium has atomic number 20. Fluorine has atomic number 9. Describe what happens when calcium reacts with fluorine and explain the bonding in the compound formed. [4 marks]
Answer:
- Calcium has electron configuration 2,8,8,2 and loses two electrons to form Ca²⁺ [1 mark]
- Fluorine has electron configuration 2,7 and gains one electron to form F⁻ [1 mark]
- Two fluorine atoms react with each calcium atom to form CaF₂ [1 mark]
- Ionic bonding occurs through electrostatic attraction between oppositely charged ions [1 mark]
Example 2: Balancing equations and relative formula mass
Question: Aluminium reacts with iron(III) oxide: Al + Fe₂O₃ → Al₂O₃ + Fe
(a) Balance the equation [1 mark]
(b) Calculate the relative formula mass of iron(III) oxide (Ar values: Fe = 56, O = 16) [2 marks]
Answer:
(a) 2Al + Fe₂O₃ → Al₂O₃ + 2Fe [1 mark]
(b) Mr = (56 × 2) + (16 × 3) [1 mark] = 112 + 48 = 160 [1 mark]
Example 3: Mole calculations
Question: Calculate the mass of carbon dioxide produced when 5.0 g of calcium carbonate thermally decomposes completely. (Ar values: Ca = 40, C = 12, O = 16) [4 marks]
CaCO₃ → CaO + CO₂
Answer:
- Mr of CaCO₃ = 40 + 12 + (16 × 3) = 100 [1 mark]
- Moles of CaCO₃ = 5.0 ÷ 100 = 0.05 mol [1 mark]
- From equation, 1 mol CaCO₃ produces 1 mol CO₂, so 0.05 mol CO₂ produced [1 mark]
- Mr of CO₂ = 12 + (16 × 2) = 44
- Mass = 0.05 × 44 = 2.2 g [1 mark]
Common mistakes and how to avoid them
Mistake: Confusing mass number with atomic number or assuming they're the same value. Correction: Atomic number = protons only; mass number = protons + neutrons. For most elements these differ because neutrons are present.
Mistake: Changing subscript numbers when balancing equations (e.g., changing H₂O to H₂O₂). Correction: Only add coefficients in front of formulae. Changing subscripts creates different substances with different chemical properties.
Mistake: Stating that ionic compounds conduct electricity because "ions can move" without specifying when. Correction: Always state that ionic compounds conduct when molten or dissolved in water, not when solid, because ions must be free to move.
Mistake: Calculating Mr by counting each element once rather than multiplying by subscript numbers. Correction: For Ca(OH)₂, there are two oxygen atoms and two hydrogen atoms. Count every atom: 40 + (16 × 2) + (1 × 2) = 74.
Mistake: Forgetting to convert cm³ to dm³ in concentration calculations. Correction: Always divide volume in cm³ by 1000 to convert to dm³ before calculating concentration in g/dm³ or mol/dm³.
Mistake: Describing covalent bonding as "atoms giving or taking electrons". Correction: Covalent bonding involves sharing pairs of electrons between non-metal atoms, not transferring them. Use precise terminology.
Exam technique for Key Concepts in Chemistry
Command word recognition: "Describe" requires you to state features or characteristics (2-3 marks typically); "Explain" requires reasons or mechanisms (often worth more marks). For bonding questions, "describe" might ask what happens to electrons, while "explain" requires linking structure to properties.
Calculation questions: Always show working clearly across multiple steps. Even if your final answer is incorrect, you can gain method marks. Include units in your final answer. For mole calculations, write out the formula you're using first (e.g., moles = mass ÷ Mr) before substituting values.
Structure and properties questions: When explaining why a substance has particular properties, always link structure explicitly to the property. State the type of bonding, describe the structure, then explain how this causes the property. For example: "Graphite conducts electricity because it has delocalised electrons between layers that are free to move and carry charge."
Equation questions: Check your balanced equation has the same number of each atom type on both sides before moving on. If asked for state symbols, ensure you include all four correctly. Read the question carefully—some ask for word equations, others for symbol equations.
Quick revision summary
Atoms contain protons, neutrons, and electrons arranged in shells. Elements in the same group have similar properties due to identical outer electron numbers. Ionic bonding involves electron transfer between metals and non-metals, forming charged ions in giant lattices. Covalent bonding involves electron sharing between non-metals. Metallic bonding has positive ions in a sea of delocalised electrons. Chemical formulae show atom ratios; equations must be balanced. Moles link mass to Mr: moles = mass ÷ Mr. Concentration = mass ÷ volume. Structure determines properties for all substance types.