1.1.1a Atoms, elements and compounds

Atoms and elements: every substance is built from atoms

1. All substances are made of atoms, whether they are solids, liquids or gases.
2. An atom is the smallest part of an element that can exist.
3. An element is a substance made from only one type of atom.
4. About 100 different elements are known, and they are set out in the periodic table.
5. Each element in the periodic table is shown by its name and its chemical symbol.

Chemical symbols: each element has its own abbreviation

1. Each element is represented by a chemical symbol made of one or two letters.
2. The first letter is always a capital; the second letter, when present, is always lower-case.
3. The symbol $O$ means one oxygen atom, while $Na$ means one sodium atom.
4. You are given a periodic table in the exam, so you can match element names to symbols.
5. The first 20 elements are $H,\; He,\; Li,\; Be,\; B,\; C,\; N,\; O,\; F,\; Ne$ then $Na,\; Mg,\; Al,\; Si,\; P,\; S,\; Cl,\; Ar,\; K,\; Ca$.
6. Other elements to recognise include the Group 1 metals $Rb,\; Cs$, the Group 7 elements $Br,\; I$, and transition metals such as $Fe,\; Cu,\; Zn,\; Ag,\; Au$.

Compounds: elements join in fixed proportions

1. A compound contains two or more different elements chemically combined in fixed proportions.
2. Fixed proportions means every sample of the compound has its elements in the same ratio.
3. A chemical formula uses symbols and numbers to show which atoms are present and in what ratio.
4. A subscript is the small lower number showing how many atoms of the element before it are present.
5. No subscript after a symbol means one atom of that element in the simplest particle.
6. $H_{2}O$ shows water has hydrogen and oxygen in a fixed ratio of $2:1$.
7. ${CO}_2$ shows carbon dioxide has carbon and oxygen in a fixed ratio of $1:2$.
8. Different proportions give different compounds, such as carbon monoxide $CO$ and carbon dioxide ${CO}_2$.

Chemical reactions: new substances form when atoms rearrange

1. A chemical reaction forms one or more new substances.
2. The starting substances are the reactants, and the substances formed are the products.
3. Elements join to form compounds through chemical reactions.
4. The products have different properties from the reactants because they are new substances.
5. Reactions often involve a detectable energy change, seen as a temperature change or the giving out of light.
6. A compound can be split into its elements only by chemical reactions, because its chemical bonds must be broken.
7. A physical method alone cannot separate a compound into its elements.

Names and formulae: read each symbol before naming the compound

1. Two-element compounds: the first element keeps its name and the second ends in -ide.
   1. $NaCl$ is sodium chloride, $MgO$ is magnesium oxide and ${CaCl}_2$ is calcium chloride.
2. Named groups: ${CaCO}_3$ is calcium carbonate, ${CuSO}_4$ is copper sulfate, ${KNO}_3$ is potassium nitrate and $NaOH$ is sodium hydroxide.
3. Common acids: $HCl$ is hydrochloric acid, $H_2{SO}_4$ is sulfuric acid and ${HNO}_3$ is nitric acid.
4. Diatomic elements exist as two-atom molecules: $H_2,N_2,O_2,F_2,{Cl}_2,{Br}_2,I_2$.
5. Other key formulae are water $H_{2}O$, carbon dioxide ${CO}_2$, methane ${CH}_4$ and ammonia ${NH}_3$.

Chemical equations: balancing keeps every atom accounted for

1. A word equation shows a reaction using the names of reactants and products.
2. The reactants go on the left of the arrow and the products on the right, separated by plus signs.
3. The word equation for magnesium burning in oxygen is:
   1. $$magnesium+oxygen\to magnesium\; oxide$$
4. A symbol equation shows the same reaction using formulae, which must be correct before balancing:
   1. $$Mg(s)+O_2(g)\to MgO(s)$$
5. A balanced equation has the same number of atoms of each element on both sides.
6. A coefficient is a large whole number before a formula that multiplies every atom in it.
7. Balance in steps: count the atoms on each side, add coefficients until they match, then check using the smallest whole numbers.
8. The balanced equation for magnesium and oxygen is:
   1. $$2Mg(s)+O_2(g)\to 2MgO(s)$$
9. This now has two magnesium atoms and two oxygen atoms on each side.
10. The state symbols $(s),(l),(g),(aq)$ mean solid, liquid, gas and dissolved in water.
11. Always balance by changing coefficients only, never by changing formulae.

1.1.1b Half equations and ionic equations

Balanced equations: atoms and total charge must match on both sides

1. Half equations show the change for one reactant or product and include any electrons transferred.
2. Ionic equations show only the particles that react or are formed.
3. Both types must have the same number of atoms of each element on both sides.
4. The total electrical charge must also be the same on both sides.
5. Balance using whole-number coefficients in front of formulae.
6. Never change a chemical formula or the charge on an ion to make an equation balance.

Half equations: electrons balance the change in charge

1. Write the correct formulae for the particle before and after the change.
2. Balance the atoms using coefficients before balancing the charge.
3. Add the charges on each side to find the total charge.
4. Add electrons to the more positive side until the total charge is equal on both sides.
5. Place electrons on the right when the particle becomes more positively charged (it loses electrons, so this is oxidation).
6. Place electrons on the left when the particle becomes less positive or more negative (it gains electrons, so this is reduction).
7. Include state symbols when they are given or can be worked out.

Ionic equations: remove particles that stay unchanged

1. Begin with a balanced full symbol equation that includes state symbols.
2. Separate aqueous ionic compounds into their individual ions.
3. Keep solids, liquids and gases together as complete formulae.
4. Cancel any spectator ions that appear unchanged on both sides.
5. Write the remaining particles as the balanced ionic equation.
6. Check the atoms and the total charge after cancelling ions.
7. For example, magnesium displaces copper from aqueous copper sulfate:
   1. $$Mg(s)+{CuSO}_4(aq)\to {MgSO}_4(aq)+Cu(s)$$
8. The aqueous ionic compounds can be separated into their ions:
   1. $$Mg(s)+{Cu}^{2+}(aq)+{SO}_4^{2-}(aq)\to {Mg}^{2+}(aq)+{SO}_4^{2-}(aq)+Cu(s)$$
9. The sulfate ions appear unchanged on both sides, so they are spectator ions and are cancelled:
   1. $$Mg(s)+{Cu}^{2+}(aq)\to {Mg}^{2+}(aq)+Cu(s)$$
10. An ionic equation can also be built by adding two half equations together.
11. If the half equations have different numbers of electrons, multiply one or both until the electron numbers match.
12. Cancel the electrons when the half equations are added, because they are transferred between the particles.

Final checks: change coefficients, never formulae or ion charges

1. A balanced half equation must contain electrons and must balance both atoms and total charge.
2. A balanced ionic equation must contain only the particles that react or are formed.
3. Use the lowest whole-number coefficients in the finished equation.
4. Keep the correct formula for every atom, molecule and ion.
5. Keep state symbols, because they show which substances can be separated into aqueous ions.
6. Work out charge by multiplying each ion's charge by its coefficient.
7. Recheck after every change, because balancing charge can affect the number of particles shown.