Atomic structure and isotopes — OCR A-Level Chemistry

Isotopes are atoms of the same element, so they share the same atomic number and therefore the same number of protons and the same number of electrons in the neutral atom. What differs is the number of neutrons in the nucleus, and because neutrons carry mass but no charge, the mass number changes while the chemical identity stays the same. For example, chlorine-35 has 17 protons and 18 neutrons, while chlorine-37 has 17 protons and 20 neutrons; both are chlorine atoms. A question may give you a pair of nuclides and ask you to decide whether they are isotopes, so check the proton count first: equal protons means the same element, and unequal neutrons confirms isotopes.

Use the atomic number Z to give the proton count, which fixes the element. Use the mass number A to give the total of protons and neutrons, so neutrons equal A minus Z. For a neutral atom the electron count equals Z. For an ion, adjust the electron count by the charge: a positive charge means electrons have been lost, so electrons equal Z minus the charge; a negative charge means electrons have been gained, so electrons equal Z plus the magnitude of the charge. For example, an ion of magnesium with Z = 12, A = 25 and charge 2⁺ has 12 protons, 13 neutrons and 10 electrons. An oxide ion with Z = 8, A = 16 and charge 2⁻ has 8 protons, 8 neutrons and 10 electrons.

Relative mass compares a particle with a standard, not with a kilogram. OCR's standard is the mass of one ¹²C atom, defined as exactly 12 units, so one unit is one-twelfth of that mass. Relative isotopic mass compares a single isotope's atom with this standard; relative atomic mass compares the weighted mean mass of an element's atoms, allowing for isotope abundances, with the same standard. Because both are ratios to the same reference, they are dimensionless numbers. For chlorine, isotopes ³⁵Cl and ³⁷Cl in roughly 3:1 abundance give a weighted mean near 35.5, which is the relative atomic mass.

Both terms use one standard: the mass of a ¹²C atom, with one unit equal to one-twelfth of that mass. Relative isotopic mass compares one isotope's atom with this unit, so ³⁵Cl has a value near 35. Relative atomic mass compares the weighted mean mass of an element's atoms with the same unit, weighting each isotope by its fractional abundance. For chlorine, 0.75 × 35 plus 0.25 × 37 gives about 35.5. Because both are ratios to the same reference, they carry no units. A correct answer names the standard, distinguishes single isotope from weighted mean, and shows the weighting.

This row is a course-guidance pointer, not a standalone assessed statement: it opens the specification's list of the uses of mass spectrometry that OCR H432 requires you to know. Read it as a signpost to the sub-clauses that follow in 2.1.1, which set out the applications you must be able to describe and apply. Work through those sub-clauses in order, and for each one identify the species being detected and what the resulting data tell you. Note that the specification does not require you to explain how the mass spectrometer itself works internally, so direct your revision at interpreting and applying the data rather than at instrument operation. Build a short summary table of each use, the information it gives and a worked example, then check it against the specification wording.

A mass spectrometer separates ions by mass-to-charge ratio and records the relative intensity of each peak. For a sample of an element, each peak corresponds to an isotope: its position gives the relative isotopic mass and its height (relative intensity) gives the relative abundance of that isotope. To find relative atomic mass you take each isotope's relative isotopic mass, multiply by its relative abundance, add the products and divide by the total abundance. For example, if an element has isotopes of relative mass 63 and 65 in relative abundances 69 and 31, the relative atomic mass is (63 × 69 + 65 × 31) ÷ 100 = 63.62. In a multiple-choice question, check whether the data are given as relative abundances or percentages, and confirm which quantity (isotopic mass, abundance or relative atomic mass) is being asked for.

Relative atomic mass, Ar, is the weighted mean mass of the atoms of an element compared with 1/12 of the mass of an atom of carbon-12. When an element has several isotopes, each with its own abundance, the weighted mean lies between the lightest and heaviest isotope masses and is pulled towards the more abundant isotope. To calculate it, multiply each isotope mass by its relative abundance, add the products, then divide by the total abundance. For example, an element with isotopes of mass 63 (abundance 70) and 65 (abundance 30) gives (63 × 70 + 65 × 30) ÷ 100 = (4410 + 1950) ÷ 100 = 6360 ÷ 100 = 63.6. If abundances are given as fractions or percentages, divide by the total (1 or 100) accordingly. The answer is a weighted mean, not a simple average, so it need not equal any single isotope mass.

Relative molecular mass, Mr, applies to molecules and is the sum of the relative atomic masses of all atoms in one molecule. Relative formula mass applies to any formula unit, including ionic compounds such as NaCl, and is likewise the sum of the relative atomic masses of the atoms in the formula. For water, H₂O, Mr = (2 × 1.0) + 16.0 = 18.0. For sodium chloride, NaCl, the relative formula mass = 23.0 + 35.5 = 58.5. Use relative molecular mass only where discrete molecules exist; use relative formula mass for ionic or giant structures. Both are calculated the same way from relative atomic masses, so the arithmetic is identical and only the label differs.

This topic builds the model of the atom you need before bonding and periodicity. Learn the subatomic particles — protons, neutrons and electrons — with their relative masses and charges, and how they occupy the nucleus and shells. Understand atomic number (protons) and mass number (protons plus neutrons), and use these to work out numbers of each particle in atoms and ions. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so they share chemical properties but differ in mass. You should also meet relative isotopic and atomic masses and how a mass spectrometer produces them. Read the specification clauses in order and check each against your notes.