Topic 2
6 subtopics · Cambridge IGCSE Chemistry 0620
Key Idea
Atoms are made of protons, neutrons, and electrons. The proton number defines the element. In a neutral atom, protons = electrons.
Explanation
Subatomic particles: - Proton: in nucleus, relative mass = 1, charge = +1 - Neutron: in nucleus, relative mass = 1, charge = 0 - Electron: in shells around nucleus, relative mass negligible, charge = -1 Proton number (atomic number) = number of protons = number of electrons (in neutral atom). Nucleon number (mass number) = protons + neutrons. Number of neutrons = mass number - proton number.
💡 Analogy
Think of the atom like an orange: the nucleus (protons + neutrons) is the small, dense pip in the centre. The electrons are like the peel and flesh — they take up most of the space but have almost no mass. The atom is mostly empty space!
Worked Examples
Practice Questions
An atom has a proton number of 17 and a mass number of 35. How many neutrons does it have?
Key Idea
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. Relative atomic mass (Ar) is the weighted average mass of all isotopes.
Explanation
Isotopes have identical chemical properties (same electron configuration) but different physical properties. RELATIVE ATOMIC MASS (Ar) — Extended: Ar is the weighted average mass of all the isotopes of an element, compared to 1/12th of the mass of a carbon-12 atom. Formula: Ar = sum of (mass of isotope x % abundance) / 100
💡 Analogy
Isotopes are like identical twins — they look the same (same chemical behaviour) but one is heavier than the other (different number of neutrons).
Worked Examples
Practice Questions
Bromine has two isotopes: Br-79 (50.5% abundance) and Br-81 (49.5% abundance). Calculate the relative atomic mass of bromine. Show your working. [2 marks]
Key Idea
Electrons occupy shells around the nucleus. Shell 1: max 2; Shell 2: max 8; Shell 3: max 8 (for elements 1-20). Group number = outer electrons. Period number = number of shells.
Explanation
Electron shell rules: - Shell 1 (closest to nucleus): maximum 2 electrons - Shell 2: maximum 8 electrons - Shell 3: maximum 8 electrons (for elements 1-20) - Electrons always fill the lowest energy shell first. Key elements: - H (1): 1 | C (6): 2,4 | N (7): 2,5 | O (8): 2,6 - Na (11): 2,8,1 | Mg (12): 2,8,2 | Cl (17): 2,8,7 | Ar (18): 2,8,8 | Ca (20): 2,8,8,2 Group number = number of electrons in the outer shell. Period number = number of electron shells occupied.
Practice Questions
Which electron configuration belongs to an element in Group II, Period 3?
Key Idea
Ionic bonding is the electrostatic attraction between oppositely charged ions. A metal transfers electrons to a non-metal. Both achieve full outer shells.
Explanation
How ionic bonding forms: 1. Draw the electron configuration of each atom. 2. Show the transfer of electrons from the metal to the non-metal. 3. Draw the resulting ions with their charges in square brackets. Properties of ionic compounds: - High melting and boiling points: strong electrostatic forces in giant ionic lattice. - Conduct electricity when molten or dissolved in water: ions are free to move. - Do NOT conduct electricity when solid: ions are in fixed positions. - Often soluble in water. Examples: NaCl (Na+ and Cl-), MgO (Mg2+ and O2-), CaCl2 (Ca2+ and 2Cl-).
💡 Analogy
In ionic bonding, the metal gives away its outer electrons (like giving a gift) to the non-metal, which accepts them. Both atoms end up with a full outer shell (like a noble gas). The opposite charges attract each other strongly.
Worked Examples
Practice Questions
Explain why ionic compounds have high melting points. [2 marks]
Key Idea
Covalent bonding is the sharing of a pair of electrons between two non-metal atoms. Each shared pair forms one covalent bond.
Explanation
Simple molecular covalent substances: - Low melting and boiling points: weak intermolecular forces between molecules. - Do NOT conduct electricity: no free charged particles. - Often gases or liquids at room temperature. Giant covalent structures: - Very high melting points: thousands of strong covalent bonds must be broken. - Do NOT conduct electricity (except graphite). - Examples: diamond (C), graphite (C), silicon dioxide (SiO2). Diamond vs Graphite: - Diamond: each C bonded to 4 others in tetrahedral arrangement. Very hard, does not conduct. - Graphite: each C bonded to 3 others in layers. Layers can slide (lubricant). Conducts electricity (delocalised electrons between layers). Common molecules: H2, O2 (double bond), N2 (triple bond), H2O (bent), CO2 (linear), CH4 (tetrahedral), NH3 (pyramidal).
Practice Questions
Which of the following substances has a giant covalent structure?
Key Idea
Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a sea of delocalised (free) electrons.
Explanation
Properties of metals explained by metallic bonding: - Good conductors of electricity: delocalised electrons are free to move and carry charge. - Good conductors of heat: delocalised electrons transfer kinetic energy quickly. - High melting points: strong electrostatic forces between ions and electron sea. - Malleable and ductile: layers of ions can slide over each other without breaking the metallic bond. - Shiny appearance: delocalised electrons reflect light.
Practice Questions
Explain why metals are good conductors of electricity. [2 marks]