Atoms, ions, formulae and relative massesEdexcel International A Level Chemistry: Revision notes
Section 1
Atoms, elements, ions and molecules
An atom is the smallest part of an element that can exist and still show the chemical properties of that element. An element contains atoms with the same number of protons. An ion is an atom or group of atoms that has gained or lost electrons, so it carries an overall charge: Mg becomes Mg2+ when it loses two electrons, and Cl becomes Cl- when it gains one.
A molecule is a group of two or more atoms joined by covalent bonds, such as O2, H2O and CH4. A compound contains two or more different elements chemically bonded in a fixed ratio. Compounds may be molecular (H2O) or ionic (NaCl).
A compound is not a mixture: its elements are chemically bonded and cannot be separated by physical methods.
Section 2
Empirical and molecular formulae
The empirical formula is the simplest whole-number ratio of atoms of each element in a compound. The molecular formula is the actual number of atoms of each element in one molecule. The molecular formula is a whole-number multiple of the empirical formula: glucose has empirical formula CH2O and molecular formula C6H12O6.
Ionic compounds and giant covalent structures are written only as an empirical (formula-unit) ratio, for example NaCl or SiO2, because they have no discrete molecules.
Ethene (C2H4) and C6H12 share the empirical formula CH2. The empirical formula does not tell you the molecular formula without the relative molecular mass.
Section 3
Relative atomic mass
Relative atomic mass, Ar, is the weighted mean mass of an atom of an element compared with one twelfth of the mass of an atom of carbon-12. The word 'weighted' matters: elements have isotopes, and Ar takes account of their relative abundances. This is why chlorine has Ar = 35.5 although no chlorine atom has that mass number.
Because it is a ratio, relative atomic mass has no units. The carbon-12 isotope is the standard, so it is exactly 12 on this scale.
Do not leave out 'weighted mean' or 'one twelfth' in the definition: both are needed for the mark.
Section 4
Relative molecular mass, relative formula mass and molar mass
Relative molecular mass, Mr, is the weighted mean mass of a molecule compared with one twelfth of the mass of an atom of carbon-12. Add the relative atomic masses of all atoms in the molecular formula. For compounds with giant structures (ionic lattices, giant covalent) there are no molecules, so the term relative formula mass is used for the formula unit. Neither has units.
Molar mass, M, is the mass per mole of a substance, in g mol-1. It is numerically equal to the relative mass.
Worked example: (NH4)2SO4 = (2 x 14.0) + (8 x 1.0) + 32.1 + (4 x 16.0) = 132.1. The molar mass is 132.1 g mol-1. The percentage of N is 28.0 / 132.1 x 100 = 21.2%.
Multiply every atom inside a bracket by the subscript outside it: in (NH4)2SO4 there are 2 N and 8 H.
Section 5
Parts per million (ppm)
Parts per million expresses very small concentrations: 1 ppm is one part in 10^6 parts. To convert a percentage to ppm, multiply by 10 000: 0.93% = 9300 ppm. For gases in the atmosphere, ppm is normally by volume, so volume of gas = total volume x ppm / 10^6.
Example: 420 ppm CO2 in 100 m3 of air is 100 x 420 / 10^6 = 0.042 m3 of CO2. Remember 1 m3 = 1000 dm3 = 10^6 cm3.
Must Know
- Atom, element, ion, molecule, compound, empirical and molecular formula: know each definition
- Ar is the weighted mean mass of an atom relative to 1/12 of the mass of a carbon-12 atom, with no units
- Mr is for molecules, relative formula mass for giant structures: add the Ar values
- Molar mass is in g mol-1 and numerically equal to Mr
- ppm = parts per 10^6; % x 10 000 = ppm
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Atoms, ions, formulae and relative masses
- A garden-centre chemist checks the label on a fertiliser that contains ammonium sulfate, (NH₄)₂SO₄. Relative atomic masses: H = 1.0, N = 14.0, O = 16.0, S = 32.1.Calculate the percentage by mass of nitrogen in ammonium sulfate. Give your answer to 3 significant figures.2 marks
- Chemists measure relative atomic masses on the carbon-12 scale. A data book gives the relative atomic mass of magnesium as 24.3, although naturally occurring magnesium contains only atoms with mass numbers 24, 25 and 26.Explain why the relative atomic mass of magnesium has no units, whereas the molar mass of magnesium is 24.3 g mol⁻¹.2 marks
- A school air-quality sensor reports the carbon dioxide concentration in a sealed classroom of volume 180 m³ as 1250 ppm by volume. Dry air contains 0.93% argon by volume. 1 m³ = 1000 dm³.Calculate the volume, in dm³, of carbon dioxide in the classroom.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).