CHEM

Homogeneous mixture

A mixture that looks completely uniform throughout — you cannot see individual components. Also called a solution. Example: saltwater, air, vinegar.

Heterogeneous mixture

A mixture where you can see the different components — it is not uniform throughout. Example: sand + water, oil + water, soil.

Filtration

Separation technique used to separate an insoluble solid from a liquid. The solid is too large to pass through filter paper; the liquid (filtrate) passes through. Example: separating sand from water.

Evaporation

Separation technique used to recover a dissolved solid from a solution. The liquid is heated so it evaporates, leaving the solid behind. Example: recovering salt from saltwater.

Distillation

Separation technique used to recover a pure liquid from a solution. The solution is heated, the liquid boils and the vapour is collected and condensed back into liquid. Example: getting pure water from saltwater.

Fractional distillation

Separation technique for two or more liquids with different boiling points. Each liquid boils at its own temperature and is collected separately. Example: separating ethanol and water, separating crude oil fractions.

Separating funnel

Apparatus used to separate two immiscible (non-mixing) liquids. They form separate layers and the bottom layer is drained off through the tap. Example: separating oil and water.

Sieving

Separation technique for solids of different particle sizes. Smaller particles fall through the holes in the sieve; larger ones are retained. Example: separating sand from gravel or partly-crushed rock.

Decanting

Separation technique where a liquid is carefully poured off from a settled solid sediment, leaving the solid behind. Example: pouring off water after soil has settled.

Chromatography

Separation technique for dissolved substances within a mixture. Different substances travel different distances up the paper based on their solubility and attraction to the paper. Example: separating inks or dyes.

Crystallisation

Separation technique to recover a solid from solution by cooling the solution slowly so the dissolved solid forms crystals. Example: growing copper sulfate crystals from solution.

Gravimetric analysis

A quantitative technique that uses mass measurements to determine the amount or percentage composition of a component in a mixture or compound.

Percentage composition by mass (mixtures)

The mass of a component divided by the total mass of the mixture, multiplied by 100. Formula: % = (mass of component / total mass) × 100

Group (periodic table)

A vertical column of elements in the periodic table. Elements in the same group have the same number of outer (valence) electrons and therefore similar chemical properties.

Period (periodic table)

A horizontal row of elements in the periodic table. Going across a period, the atomic number increases by one each step, and elements transition from metals to metalloids to non-metals.

Atomic number

The number of protons in the nucleus of an atom. It uniquely identifies the element and determines its position in the periodic table.

Valence electrons

The electrons in the outermost shell of an atom. These determine the chemical properties of the element and how it bonds with other elements.

Metals — physical properties

Shiny lustre, good conductors of electricity and heat, malleable (can be beaten flat), ductile (can be drawn into wires), solid at room temperature (except mercury which is liquid).

Non-metals — physical properties

Dull appearance, poor conductors of electricity and heat, brittle if solid, many are gases at room temperature. Examples: oxygen, nitrogen, sulfur, chlorine.

Noble gases

Group 18 elements: He, Ne, Ar, Kr, Xe, Rn. They are monatomic (single atoms), unreactive, and have a full outer electron shell. They exist as single atoms, not molecules.

Diatomic elements

Elements that exist naturally as pairs of atoms bonded together. The 7 diatomic elements are: H2, O2, N2, F2, Cl2, Br2, I2. Memory trick: HOFBrINCl.

Monatomic elements

Elements that exist as single, unbonded atoms. All noble gases are monatomic: He, Ne, Ar, Kr, Xe, Rn.

Metalloids

Elements with properties between metals and non-metals. Found along the staircase line on the periodic table. Examples: Silicon (Si), Germanium (Ge), Arsenic (As).

Trend: metallic character across a period

Metallic character decreases from left to right across a period — elements go from metals on the left, through metalloids, to non-metals on the right.

Binary ionic compound naming rule

Name the metal (cation) first, then the non-metal with an '-ide' ending. Example: NaCl = sodium chloride; MgO = magnesium oxide; AlP = aluminium phosphide.

Transition metal compound naming rule

Use Roman numerals in brackets to indicate the charge of the transition metal. Example: FeCl2 = iron(II) chloride; FeCl3 = iron(III) chloride; CuO = copper(II) oxide.

Covalent compound naming rule

Use prefixes (mono, di, tri, tetra, penta…) to show the number of each atom. Do not use 'mono' for the first element. Example: CCl4 = carbon tetrachloride; N2O4 = dinitrogen tetroxide; CO2 = carbon dioxide.

Carbonate ion

CO3^2- (charge of 2−). Example compound: CaCO3 = calcium carbonate; Na2CO3 = sodium carbonate.

Sulfate ion

SO4^2- (charge of 2−). Example compound: Na2SO4 = sodium sulfate; CuSO4 = copper(II) sulfate.

Nitrate ion

NO3^- (charge of 1−). Example compound: NaNO3 = sodium nitrate; Ca(NO3)2 = calcium nitrate.

Hydroxide ion

OH^- (charge of 1−). Example compound: NaOH = sodium hydroxide; Mg(OH)2 = magnesium hydroxide.

Phosphate ion

PO4^3- (charge of 3−). Example compound: Ca3(PO4)2 = calcium phosphate; Na3PO4 = sodium phosphate.

Ammonium ion

NH4^+ (charge of 1+). Example compound: NH4Cl = ammonium chloride; (NH4)2SO4 = ammonium sulfate.

Cross-multiplication rule for ionic formulas

To write the formula of an ionic compound: swap the numbers of the charges and use them as subscripts. Example: Al^3+ and O^2- gives Al2O3. Ca^2+ and PO4^3- gives Ca3(PO4)2.

Molar mass of H2O

(2 × 1.0) + (1 × 16.0) = 18.0 g/mol

Molar mass of NaCl

23\ .0 + 35.5 = 58.5 g/mol

Molar mass of CO2

12\ .0 + (2 × 16.0) = 44.0 g/mol

Molar mass of NaHCO3

23 + 1 + 12 + (3 × 16) = 84.0 g/mol

Molar mass of Na2CO3

(2 × 23) + 12 + (3 × 16) = 106.0 g/mol

Molar mass of Mg(OH)2

24\ .3 + 2(16.0 + 1.0) = 58.3 g/mol

Synthesis (combination) reaction

Two or more reactants combine to form a single product. General form: A + B → AB. Example: 2Mg + O2 → 2MgO; 2NO + O2 → 2NO2.

Decomposition reaction

A single substance breaks down into two or more simpler products. General form: AB → A + B. Example: 2H2O2 → 2H2O + O2; 2HgO → 2Hg + O2.

Combustion reaction

A substance reacts with oxygen gas to release energy (heat and light). Complete combustion of hydrocarbons produces CO2 and H2O. Example: CH4 + 2O2 → CO2 + 2H2O.

Precipitation reaction

Two aqueous solutions are mixed and an insoluble solid (precipitate) forms. The precipitate is shown with (s) or a down-arrow symbol. Example: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq).

Acid-base (neutralisation) reaction

An acid reacts with a base to produce a salt and water. General form: acid + base → salt + water. Example: HCl + NaOH → NaCl + H2O.

Acid-carbonate reaction

An acid reacts with a carbonate to produce a salt, water, and carbon dioxide gas. Example: 2HCl + CaCO3 → CaCl2 + H2O + CO2.

Indicators of a chemical change

Colour change, gas produced (bubbles/fizzing), precipitate forms (solid appears in solution), temperature change (exo- or endothermic), light produced.

Exothermic reaction

A reaction that releases energy (heat) to the surroundings — the temperature of the surroundings increases. Example: combustion, neutralisation.

Endothermic reaction

A reaction that absorbs energy (heat) from the surroundings — the temperature of the surroundings decreases. Example: dissolving ammonium nitrate in water.

Law of conservation of mass

In a chemical reaction, matter is neither created nor destroyed. The total mass of reactants equals the total mass of products. This is why equations must be balanced.

State symbols in equations

(s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous (dissolved in water). Written in brackets after each formula in a balanced equation.

Coefficient (in a chemical equation)

A large number written in front of a formula showing how many moles of that substance are involved. Example: in 2H2O, the coefficient is 2, meaning 2 moles of water.

Subscript (in a chemical formula)

A small number written below and after an element symbol showing how many atoms of that element are in one molecule. Example: H2O has subscript 2 for H, meaning 2 hydrogen atoms per molecule.

Rule: what can you change when balancing an equation?

You may ONLY change coefficients (the big numbers in front of formulas). You must NEVER change subscripts inside a formula — that would change the substance entirely.

Steps to balance a chemical equation

1. Write the correct unbalanced equation with correct formulas. 2. Count atoms of each element on both sides. 3. Add/adjust coefficients to make both sides equal. 4. Check all atoms balance. Example: H2 + O2 → H2O becomes 2H2 + O2 → 2H2O.

Conservation of atoms in equations

Atoms are rearranged in a chemical reaction but never created or destroyed. The number of each type of atom must be the same on both sides of a balanced equation.

The mole (definition)

A unit used to count atoms, molecules, or ions. One mole = 6.022 × 10^23 particles. It is the amount of substance that contains as many particles as there are atoms in exactly 12g of carbon-12.

Avogadro's number

6\ .022 × 10^23 — the number of particles in one mole of any substance. Named after Amedeo Avogadro.

Molar mass (definition)

The mass of one mole of a substance, measured in g/mol. For elements, it equals the atomic mass from the periodic table. For compounds, it is the sum of all atomic masses in the formula.

Mole formula

n = m / MM where: n = amount in moles (mol), m = mass in grams (g), MM = molar mass (g/mol). Rearranged: m = n × MM and MM = m/n.

How many moles in 36g of H2O?

MM(H2O) = 18 g/mol. n = m/MM = 36/18 = 2.0 mol

What is the mass of 3 moles of CO2?

MM(CO2) = 44 g/mol. m = n × MM = 3 × 44 = 132 g

How many particles in 2 moles of NaCl?

Particles = n × Avogadro's number = 2 × 6.022×10^23 = 1.204×10^24 formula units.

Number of particles formula

N = n × Na where N = number of particles, n = moles, Na = 6.022×10^23 /mol. Rearranged: n = N / Na.

Percentage composition by mass (compound)

% of element X = (mass of X in formula / molar mass of compound) × 100. Example: % O in H2O = (16.0/18.0) × 100 = 88.9%

% Mg in Mg(OH)2

MM = 58.3 g/mol. %Mg = (24.3/58.3) × 100 = 41.7% ≈ 42%

Empirical formula (definition)

The simplest whole-number ratio of atoms of each element in a compound. Example: glucose C6H12O6 has empirical formula CH2O.

Steps to find an empirical formula from % composition

1. Assume 100g (so % = grams). 2. Divide each mass by its molar mass to get moles. 3. Divide all values by the smallest mole value. 4. Round to whole numbers (multiply through if needed to clear decimals).

Molecular formula vs empirical formula

Empirical formula = simplest ratio (e.g. CH2O). Molecular formula = actual number of atoms (e.g. C6H12O6). Molecular formula = n × empirical formula, where n is a whole number.

Stoichiometry (definition)

The calculation of quantities (masses, moles, volumes) of reactants and products in a chemical reaction, based on the balanced equation and mole ratios.

What the coefficients in a balanced equation tell you

The mole ratio of all reactants and products. Example: N2 + 3H2 → 2NH3 means 1 mol N2 reacts with 3 mol H2 to produce 2 mol NH3.

4-step method for stoichiometry calculations

Step 1: Write the balanced equation. Step 2: Find moles of the given substance using n = m/MM. Step 3: Use the mole ratio from the equation to find moles of the wanted substance. Step 4: Convert moles back to mass using m = n × MM.

Limiting reagent (definition)

The reactant that is completely used up first in a reaction, causing the reaction to stop. The amount of product formed is determined by the limiting reagent.

How to identify the limiting reagent

1. Find moles of each reactant. 2. Divide each by its coefficient in the balanced equation. 3. The reactant with the SMALLER result is the limiting reagent.

Excess reagent (definition)

The reactant that is not completely used up — some is left over after the limiting reagent runs out.

Theoretical yield

The maximum mass of product that could be obtained from a reaction, calculated from the moles of the limiting reagent using stoichiometry.

How does the mole ratio work in mass calculations?

Example: 2H2 + O2 → 2H2O. If you start with 4g H2 (n=2 mol), the ratio is 2:2 = 1:1 with water, so you produce 2 mol H2O = 36g. The ratio comes from the coefficients.

Concentration (definition)

The amount of solute dissolved per litre of solution. Measured in mol/L (also written as M or mol/L). A more concentrated solution has more solute per unit volume.

Concentration formula

c = n / V where: c = concentration (mol/L), n = moles of solute (mol), V = volume of solution in LITRES (L). Important: always convert mL to L by dividing by 1000.

Rearrangements of c = n/V

n = c × V (to find moles); V = n / c (to find volume); c = n / V (to find concentration). Always ensure V is in litres.

Solute

The substance that is dissolved in a solution. Example: in saltwater, salt (NaCl) is the solute.

Solvent

The liquid in which the solute is dissolved. Example: in saltwater, water is the solvent.

Solution

A homogeneous mixture formed when a solute dissolves in a solvent. Example: saltwater, copper sulfate solution.

Dilution formula

c1V1 = c2V2 where subscript 1 = before dilution, subscript 2 = after dilution. When you dilute a solution, moles of solute are conserved — only the volume changes.

What stays the same during a dilution?

The number of moles of solute stays constant. Adding more water (solvent) increases the volume, which decreases the concentration, but the total moles don't change.

Worked example: dilution

Take 20 mL of a 0.5 mol/L solution and dilute to 200 mL. Using c1V1 = c2V2: 0.5 × 20 = c2 × 200, so c2 = 0.05 mol/L.

How to find concentration of an ion in solution

First find concentration of the compound. Then multiply by the number of that ion per formula unit. Example: 0.1 mol/L Na2CO3 gives [Na+] = 0.1 × 2 = 0.2 mol/L (because Na2CO3 gives 2 Na+ ions per formula unit).

Standard solution (definition)

A solution whose concentration is known precisely and accurately. Used as a reference in volumetric analysis (titrations).

Primary standard (definition)

A substance used to make a standard solution directly by weighing. It must be: pure, stable (doesn't absorb moisture or react with air), high molar mass (reduces % error in weighing), and soluble in water.

Examples of primary standards

Sodium carbonate (Na2CO3), potassium hydrogen phthalate (KHP), anhydrous sodium carbonate. Sodium carbonate is commonly used in Year 11.

Steps to prepare a standard solution

1. Accurately weigh the required mass of primary standard. 2. Dissolve in a small volume of distilled water in a beaker. 3. Transfer quantitatively to a volumetric flask of the required volume. 4. Rinse the beaker several times and add rinsings to the flask. 5. Add distilled water to the calibration mark. 6. Invert and mix thoroughly.

Volumetric flask

Glassware used to prepare solutions of an exact, known volume. The calibration mark on the neck indicates the precise volume (e.g. 100 mL or 250 mL). The solution must be made up to this mark precisely.

Why use distilled water (not tap water) in standard solutions?

Tap water contains dissolved ions (Ca2+, Mg2+, Cl- etc.) that would contaminate the solution and change its concentration. Distilled water is pure H2O with no dissolved substances.

Quantitative transfer

The process of carefully rinsing a beaker multiple times with distilled water and adding all rinsings to the volumetric flask, to ensure no solute is left behind and the full amount is transferred.