Formulas — Combined Science 0653 - Chemistry Revision Notes & Topical Questions

Section 1 of 24Chemical symbols and formulae

Symbol

• Symbol each element has a chemical symbol from the Periodic Table, such as Li, Ca, Au and S.

Molecular formula

• = the number and type of atoms in one molecule of a compound.
• The symbol shows the type of atom.
• The subscript number after a symbol shows how many atoms there are.

Section 2 of 24The seven diatomic elements

Diatomic molecules

• Diatomic molecules these elements exist as molecules of two atoms: H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂.

Simple molecules

• Simple molecules the seven diatomic elements are also simple molecules.

Section 3 of 24Formulae of common simple molecules

Substance	Formula	Atoms present
Water	H2O	2 H, 1 O
Methane	CH4	1 C, 4 H
Ammonia	NH3	1 N, 3 H
Sulfuric acid	H2SO4	2 H, 1 S, 4 O

Section 4 of 24Deducing a formula from a model

Method

• Method count the atoms of each element in the diagram and use the key for the colours.

Arrangement

• Arrangement a model also shows how atoms are arranged, such as the two hydrogen atoms on either side of the oxygen atom in water.

Section 5 of 24Empirical formula (supplement)

Empirical formula

• = simplest whole-number ratio of atoms of each element in a molecule or formula unit.

Molecular and empirical

• Molecular and empirical for organic compounds the two are often different.
• Ethanoic acid is C₂H₄O₂ with empirical formula CH₂O.

Ionic compounds

• Ionic compounds the formula is always the empirical formula, such as NaCl.

Section 6 of 24Ions needed for formulae (supplement)

Type	Ions
Simple positive ions	H+, Li+, Na+, K+, Cu2+, Fe2+, Fe3+
Simple negative ions	F-, Cl-, Br-
Polyatomic ions	CO32-, SO42-, OH-, NO3-, NH4+

Section 7 of 24Formula of an ionic compound (supplement)

Rule

• Rule an ionic compound has no overall charge, so positive and negative charges must cancel.

Direct comparison

• Direct comparison balance the charges by choosing numbers of each ion.
• Fe²⁺ and SO₄^2- give FeSO₄.

Swap and drop

• Swap and drop swap the charge numbers between the ions, drop them as subscripts and drop any 1.
• Cu²⁺ and Cl^- give CuCl₂.
• Fe³⁺ and Cl^- give FeCl₃.
• If both charges are equal, give the simplest ratio.

Brackets

• Brackets brackets are used round polyatomic ions when more than one is needed, as in Mg(NO₃)₂.

Section 8 of 24More from the specification

1Empirical formula determination

• Empirical formula determination calculating simplest integer mole ratios of constituent elements provides an analytical means to establish chemical formulas.

2Balanced chemical equation

• A symbolic representation of a reaction with equal numbers of each type of atom on both reactant and product sides.

3Ionic formula deduction

• Balancing total positive and negative ionic charges so that compounds like copper(II) chloride (CuCl₂) or sodium sulfate (Na₂SO₄) remain electrically neutral.

4Relative formula mass (Mr)

• The sum of all relative atomic masses of elements indicated in a given chemical formula.

5Word equation

• An introductory expression naming the reacting starting materials and final products before quantitative balancing.

6Symbol equation

• Reaction written with formulae.

7Carbon dioxide

• A gas that turns limewater cloudy.

8Sodium chloride

• Compound of sodium and chlorine; common salt.

9Chloride

• Salt of hydrochloric acid.

10Sulfate

• Salt of sulfuric acid.

11Hydroxide

• A compound containing the OH⁻ ion, formed by Group I metals with water.

12Product

• A substance formed in a reaction.

13Nitrate

• Salt of nitric acid.

14Reactants

• The substances that react together at the start of a reaction.

Section 9 of 24Word equations

Word equation

• Word equation shows the reactants and products of a reaction by their full names.

Reactants

• Reactants the substances on the left of the arrow, the chemical ingredients.

Products

• Products the new substances on the right of the arrow.

Arrow

• Arrow means "to form" or "produces" and conditions or a catalyst can be written above it.

Section 10 of 24Writing a word equation from a description

1

Pick out the reactants and the products from the description.

2

Check which side is which, because products may be given first.

3

Add reactants that are implied, since burning in air means reacting with oxygen.

4

Write reactants + reactants → products.

• Example form
  • sodium hydroxide + hydrochloric acid → sodium chloride + water.

Section 11 of 24Symbol equations

Symbol equation

• Symbol equation uses formulae of reactants and products.

Rules

• Rules reactants on the left and products on the right.
• Write H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂ as molecules.
• Include state symbols.

Section 12 of 24State symbols

Symbol	Meaning
(s)	Solid
(l)	Liquid
(g)	Gas
(aq)	Aqueous (dissolved in water)

Clues

• Clues metal compounds are solids unless in solution and ionic compounds are usually solids.

Non-metal compounds

• Non-metal compounds can be solids, liquids or gases, so use the information given.

Precipitate

• Precipitate a precipitate formed in solution counts as a solid.

Solutions

• Solutions ethanol mixed with water is (aq); pure ethanol is (l).

Section 13 of 24Balancing equations

Conservation of mass

• Conservation of mass the number of atoms of each element must be the same on both sides.

Method

• Method work from left to right, changing the numbers in front of formulae one at a time.
• Balance elements that appear on their own last.
• Count an unchanged group such as NO₃ as one unit.

Never

• Never change a subscript, because that changes the substance.

Section 14 of 24Balanced equations as models

Neutralisation

• NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l).

Combustion of sulfur

• S(s) + O₂(g) → SO₂(g) shows one atom of sulfur reacting with one oxygen molecule.

Metal oxide + acid

• MgO(s) + 2HNO₃(aq) → Mg(NO₃)₂(aq) + H₂O(l).

Section 15 of 24Checking a balanced equation

Checking a balanced equation

Count each element on both sides, including atoms inside brackets.

Multiply bracket subscripts by the number outside.

Check state symbols last.

Section 16 of 24More from the specification

1Relative Atomic Mass (Ar) and Formula Mass (Mr)

• Relative atomic mass is the weighted mean mass of an atom of an element compared to 1/12th the mass of a carbon-12 atom;
  • relative formula mass (Mr) is the sum of Ar values in a formula unit.

2Determining Empirical Formulas

• The empirical formula is the simplest whole-number ratio of atoms of each element in a compound, deduced by dividing the mass or percentage of each element by its Ar and simplifying the molar ratio.

3Balancing Ionic Compound Charges

• Chemical formulas of neutral ionic compounds are established by balancing positive and negative charges (e.g. Al³⁺ and NO₃⁻ require three nitrate ions, yielding aluminium nitrate Al(NO₃)₃).

4Hydrated Copper(II) Sulfate Formula

• Copper(II) sulfate combines with five molecules of water of crystallisation to form blue hydrated crystals with the chemical formula CuSO₄·5H₂O.

5Empirical formula

• The simplest whole-number ratio of atoms or ions.

6Carbon dioxide

• A gas that turns limewater cloudy.

7Simple molecule

• Small group of atoms joined by covalent bonds, with weak forces between molecules.

8Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

9Sulfate

• Salt of sulfuric acid.

10Nitrate

• Salt of nitric acid.

Section 17 of 24Deducing a symbol equation (supplement)

Deducing a symbol equation (supplement)

Work out the formulae and state symbols of every reactant and product.

Write the unbalanced equation.

Balance by changing coefficients only.

Metals

• Metals a pure metal is an element, so its formula is its symbol.

Non-metal gases

• Non-metal gases hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine and iodine are diatomic.

Ionic compounds

• Ionic compounds the formula comes from the charges on the ions so that the compound is neutral.

Section 18 of 24Balancing tips for deduced equations

Even numbers

• Even numbers make the number of atoms of an element even on one side by changing coefficients, then match the other side.

Order

• Order balance one element at a time and recount at the end.

Section 19 of 24Dissociation of ionic compounds

Dissociation

• Dissociation in aqueous solution ionic compounds separate into their component ions.
• HCl(aq) → H⁺(aq) + Cl^-(aq).
• KOH(aq) → K⁺(aq) + OH^-(aq).

Compounds to recognise

• Compounds to recognise acids such as HCl and H₂SO₄, Group I and II hydroxides and soluble salts.

Section 20 of 24Writing an ionic equation

Writing an ionic equation

Write the full balanced symbol equation with state symbols.

Replace each aqueous ionic compound with its ions.

Cancel the ions that appear unchanged on both sides.

State symbols

• State symbols ionic equations should always include state symbols.

Spectator ions

• Spectator ions ions that appear on both sides take no part in the reaction.

Section 21 of 24Halogen displacement as an ionic equation

• Reaction
  • aqueous chlorine displaces iodide ions: 2I^-(aq) + Cl₂(aq) → 2Cl^-(aq) + I₂(aq).
  • The K⁺ ions are spectator ions.

Section 22 of 24More from the specification

1State symbols and ionic representation

• State symbols and ionic representation indicating physical states and spectator ion cancellation provides an accurate means to formulate net ionic equations.

2Chemical formulas and valency rules

• Chemical formulas reflect the ratio of constituent elements, determined by balancing positive oxidation states with negative oxidation states.

3Relative molecular mass vs formula mass

• Relative molecular mass (Mr) applies to discrete covalent molecules (e.g. CO₂, H₂O),
• while relative formula mass applies to giant ionic or network lattices.

4Determination of empirical formula from combustion data

• Combustion of organic compounds provides masses of CO₂ and H₂O,
• from which moles of carbon and hydrogen are calculated to deduce empirical formulas.

5Empirical and molecular formula determination

• Empirical and molecular formula determination empirical formulas express the simplest whole-number atomic ratio,
  • whereas molecular formulas show actual atom counts in a molecule.

6Precipitation of insoluble copper compounds

• Precipitation of insoluble copper compounds mixing copper sulfate solution with aqueous sodium hydroxide precipitates light blue copper(II) hydroxide, leaving soluble sodium sulfate.

7Decomposition of metal nitrate salts

• Decomposition of metal nitrate salts heating hydrated copper nitrate releases brown nitrogen dioxide gas, oxygen, and black copper(II) oxide residue.

8Empirical formula

• The simplest whole-number ratio of atoms or ions.

9Carbon dioxide

• A gas that turns limewater cloudy.

10Sodium chloride

• Compound of sodium and chlorine; common salt.

11Simple molecule

• Small group of atoms joined by covalent bonds, with weak forces between molecules.

12Sulfate

• Salt of sulfuric acid.

13Molecule

• Two or more atoms held together by covalent bonds.

14Nitrate

• Salt of nitric acid.

Section 23 of 24Key terms

Term	Meaning
Molecular formula	Number and type of atoms in one molecule.
Empirical formula	Simplest whole-number ratio of atoms.
Diatomic	Molecule made of two atoms.
Polyatomic ion	Ion containing more than one atom.
Subscript	Small number showing how many atoms.
Formula unit	Simplest ratio of ions in an ionic compound.

Term	Meaning
Reactant	Substance that reacts.
Product	Substance formed.
Word equation	Reaction written with names.
Symbol equation	Reaction written with formulae.
State symbol	Letters (s), (l), (g), (aq) showing physical state.
Coefficient	Number in front of a formula in an equation.
Catalyst	Substance that speeds up a reaction and is not used up.

Term	Meaning
Dissociation	Separation of an ionic compound into ions in solution.
Ionic equation	Equation showing only the ions that react.
Spectator ion	Ion unchanged by the reaction.
Aqueous	Dissolved in water.
Diatomic	Molecule of two atoms.
Coefficient	Number in front of a formula.

Section 24 of 24Exam technique

1

Writing a diatomic element as a single atom — chlorine gas is Cl_{2}.

2

Changing a subscript to balance charges or atoms — subscripts are fixed by the compound.

3

Treating ionic compounds as molecules — they are giant lattices written with the simplest ratio.

4

Forgetting that the ionic compound has no overall charge.

1

Changing subscripts to balance an equation.

2

Missing the products when they are written first in the description.

3

Forgetting that a precipitate is a solid (s).

4

Writing chlorine or oxygen gas as single atoms.

1

Leaving spectator ions in the ionic equation.

2

Splitting a solid, liquid or gas into ions — only aqueous ionic compounds dissociate.

3

Forgetting charges when balancing ionic equations.

4

Using Cl instead of Cl_{2} for chlorine.