chem

Proton

Relative mass = 1, relative charge = +1

Neutron

Relative mass = 1, relative charge = 0

Electron

Relative mass = 1/1840, relative charge = -1

Isotope

Atoms of the same element with same number of protons, different number of neutrons, same chemical properties.

Relative atomic mass

Average mass of one atom of an element compared to 1/12 the mass of one atom of carbon-12.

First ionisation energy

Energy required to remove one mole of electrons from one mole of gaseous atoms, M(g) → M⁺(g) + e⁻.

Trend down a group

First ionisation energy decreases, atomic radius increases, more electron shielding, weaker nuclear attraction.

Trend across a period

First ionisation energy increases, nuclear charge increases, similar shielding, atomic radius decreases.

Small drops in ionisation energy

Between group 2 and 3 = new p subshell higher energy; between group 5 and 6 = paired electrons repel.

Mass spectrometry stages

Ionisation, acceleration, deflection, detection.

Moles

Moles = mass / molar mass.

Ideal gas equation

PV = nRT, where P = pressure in Pa, V = volume in m³, T = temperature in Kelvin, R = 8.31 J K⁻¹ mol⁻¹.

Percentage yield

(actual mass / theoretical mass) × 100.

Atom economy

(molar mass of desired product / total molar mass of all reactants) × 100.

Ionic equations

Show only reacting ions; cancel spectator ions.

Molar volume at rtp

24 dm³ mol⁻¹.

Ionic bonding

Strong electrostatic attraction between oppositely charged ions, forming a giant ionic lattice.

Properties of ionic compounds

High melting and boiling points, conductive only when molten or dissolved, soluble in polar solvents.

Covalent bonding

Shared pair of electrons between two atoms.

Dative covalent bond

Both electrons in the shared pair come from the same atom.

Simple molecular structure

Weak intermolecular forces, low melting and boiling points, does not conduct electricity.

Giant covalent structure

Strong covalent bonds throughout lattice, very high melting points.

Properties of metals

High melting points, conducts heat and electricity, malleable and ductile.

Electronegativity

Ability of an atom to attract the bonding pair of electrons in a covalent bond.

Intermolecular forces

Includes London forces, permanent dipole-dipole forces, and hydrogen bonding.

Dynamic equilibrium

Forward and reverse reactions occur at the same rate, concentrations remain constant in a closed system.

Oxidation

Loss of electrons, increase in oxidation number.

Reduction

Gain of electrons, decrease in oxidation number.

Oxidising agent

Causes oxidation, itself reduced.

Reducing agent

Causes reduction, itself oxidised.

Periodicity

Repeating pattern of physical and chemical properties across a period.

Group 2 reactions with oxygen

2M + O₂ → 2MO, forms white solid.

Uses of chlorine

Disinfect drinking water, make bleach, prevent disease.

Flame tests for calcium

Brick red flame.

Flame tests for strontium

Crimson red flame.

Flame tests for barium

Apple green flame.

Flame tests for magnesium

No visible flame colour.

Group 2 reactions with water

Group 2 metals react with water to form hydroxides and hydrogen gas. For example, $M + 2H<em>2O \rightarrow 2MOH + H</em>2$.

Group 2 hydroxides

Group 2 hydroxides, such as calcium hydroxide, are soluble in water and form alkaline solutions.

Group 2 sulfates

Group 2 sulfates (e.g., magnesium sulfate, calcium sulfate) generally decrease in solubility down the group, with barium sulfate being insoluble.

Flame tests for Group 2 metals

When subjected to flame tests, Group 2 metals produce different flame colors; e.g., calcium gives a brick red flame, strontium a crimson red, and barium an apple green.

Group 7 reactions with metals

Group 7 elements (halogens) can react with metals to form ionic compounds, e.g., $2M + nX<em>2 \rightarrow 2MX</em>n$ where X is a halogen.

Tests for halides

To test for halides, add dilute nitric acid followed by silver nitrate solution: AgCl forms a white precipitate, AgBr a cream precipitate, and AgI a yellow precipitate.

Reactivity of Group 7

Reactivity decreases down the group; fluorine is the most reactive, while iodine is the least.

Hydration of Group 2 hydroxides

The solubility of Group 2 hydroxides increases down the group, as seen with magnesium hydroxide (insoluble) to barium hydroxide (soluble).