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1.1Physical Chemistry

Atomic Structure — AQA A Level Chemistry Revision Notes & Topical Questions

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Atomic Structure: Fundamental Particles, Mass Number and Isotopes — AQA A LEVEL CHEMISTRY — COMPLETE NOTES notes: State the relative mass and charge of protons, neutrons and electrons, work out particle numbers from atomic and mass numbers, and explain isotopes.Atoms consist of a…

  • atomic number
  • relative atomic mass
  • Worked example 1
  • Worked example 2
  • Worked example 3
  • number of protons = Z;

Section 1 of 18The fundamental particles

Atoms consist of a tiny, dense, positively charged nucleus containing protons and neutrons, surrounded by electrons in shells (energy levels). Almost all of the mass of an atom is in the nucleus, but the atom's volume is mostly empty space occupied by electrons.

ParticleRelative massRelative charge
Proton1+1
Neutron10
Electron1/1836 (negligible)-1

Section 2 of 18Atomic number, mass number and ions

The atomic number (Z) is the number of protons in the nucleus; it defines the element. The mass number (A) is the total number of protons plus neutrons. In the notation AZX:

Atomic number, mass number and ions
number of protons = Z;
number of neutrons = A - Z;
number of electrons = Z in a neutral atom.

Ions form when atoms gain or lose electrons. A positive ion (cation) has fewer electrons than protons; a negative ion (anion) has more. For example, 2713Al3+ has 13 protons, 14 neutrons and 10 electrons.

Section 3 of 18Isotopes

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so they have different mass numbers. For example, 35Cl has 18 neutrons and 37Cl has 20 neutrons.

Isotopes of an element have the same chemical properties because chemical behaviour depends on the electron arrangement, which is determined by the number of protons. They differ in physical properties that depend on mass, such as density and rate of diffusion.

The relative atomic mass (Ar) of an element is the weighted mean mass of its atoms compared with 1/12 of the mass of a carbon-12 atom, and is calculated from the relative abundance of the isotopes: Ar = Σ(isotopic mass × abundance) ÷ 100 (if abundances are percentages).

Section 4 of 18Worked examples

  1. State the number of protons, neutrons and electrons in 2412Mg2+. [3]3 marks

  2. Chlorine consists of 75% 35Cl and 25% 37Cl. Calculate its relative atomic mass. [2]2 marks

  3. Explain why isotopes of an element have the same chemical properties but different physical properties. [2]2 marks

Section 5 of 18What mass spectrometry does

Mass spectrometry is a powerful analytical technique.

What mass spectrometry does
It is the most useful instrument for the accurate determination of relative atomic and molecular masses,
  • and it also gives information on the abundance of each isotope in an element.
A sample is turned into positive ions, which are separated according to their mass-to-charge ratio (m/z).
For ions with a charge of +1, m/z equals the relative mass of the ion.

Section 6 of 18The stages in a TOF spectrometer

1
Ionisation.
The sample is vaporised.
  • In electrospray ionisation the sample is dissolved in a volatile solvent and pushed through a fine needle at high voltage,
    • so each molecule gains a proton: X + H+ → XH+.
  • In electron impact ionisation the vaporised sample is bombarded with high-energy electrons which knock off an electron: X(g) + e- → X+(g) + 2e-.
2
Acceleration.
An electric field accelerates the positive ions so that all have the same kinetic energy.
3
Ion drift.
  • The ions pass through a flight tube with no electric field.
  • Lighter ions travel faster (same kinetic energy, KE = ½mv2) and heavier ions more slowly.
    What the symbols mean: Ek = kinetic energy (J) · m = mass (kg) · v = speed (m s−1)
4
Detection.
  • Ions reach a detector, which records the time of flight; the time is used to calculate m/z.
  • The signal height is proportional to the abundance of the ion.

The flight time is t = d ÷ v. The sample is kept under a vacuum so ions do not collide with air particles.

Section 7 of 18Using the spectrum

A mass spectrum plots relative abundance against m/z.

For an element, each peak is an isotope, and the relative atomic mass is the weighted mean: Ar = Σ(m/z × abundance) ÷ Σabundance.
Molecular ion
For a molecule, the peak with the highest m/z is the molecular ion (M+), and its m/z gives the relative molecular mass.
Note that in electrospray ionisation the peak is M + 1, since a proton has been added.

Section 8 of 18Worked examples

  1. A sample of an element shows peaks at m/z 24 (79%), 25 (10%) and 26 (11%). Calculate Ar to 1 decimal place. [2]2 marks

  2. Explain why a TOF mass spectrometer is kept under vacuum. [1]1 mark

  3. Explain why lighter ions reach the detector first. [2]2 marks

Section 9 of 18Shells, subshells and orbitals

Electrons occupy shells (principal quantum number n = 1, 2, 3 ...). Each shell is divided into subshells (s, p, d, f), each of which contains orbitals: a region of space where there is a high probability of finding an electron. Each orbital holds a maximum of two electrons, with opposite spins.

SubshellOrbitalsMax electrons
s12
p36
d510
f714

The maximum number of electrons in shell n is 2n2. An s orbital is spherical; p orbitals are dumbbell-shaped, pointing along the x, y and z axes.

Section 10 of 18Filling the orbitals

Electrons fill the lowest-energy subshell first. The order is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. Notice that 4s fills before 3d. Within a subshell, electrons occupy separate orbitals singly before pairing (Hund's rule), because electrons repel each other.

Examples

  • Na = 1s22s22p63s1;
    • Fe = 1s22s22p63s23p63d64s2.
  • Two exceptions to learn: chromium is [Ar]3d54s1 and copper is [Ar]3d104s1, because half-full and full d subshells are particularly stable.

When forming positive ions, the 4s electrons are lost first: Fe2+ = [Ar]3d6. Negative ions add electrons to the outer subshell: Cl- = 1s22s22p63s23p6.

Section 11 of 18Blocks of the periodic table

The Periodic Table is divided into four main blocks based on electron configuration. Elements are classified as s-, p-, d- or f-block according to the type of orbital that contains their outermost (valence) electron:

s-block
s-block valence electron(s) in an s orbital: Groups 1 and 2 (and helium);
p-block
p-block valence electrons in a p orbital: Groups 3 to 0 (13-18);
d-block
d-block the transition metals, filling the d subshell;
f-block
f-block the lanthanides and actinides.

The group number of an s- or p-block element equals the number of electrons in its outer shell.

Section 12 of 18Worked examples

  1. Write the full electron configuration of a sulfide ion, S2-. [1]1 mark

  2. Write the electron configuration of Fe3+. [2]2 marks

  3. State which block contains vanadium and explain why. [2]2 marks

Section 13 of 18Definitions

The first ionisation energy (IE1) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions: X(g) → X+(g) + e-.

Second ionisation energy
The second ionisation energy (IE2) is the energy required to remove the second mole of electrons from each +1 ion in a mole of gaseous +1 ions, to form one mole of +2 ions: X+(g) → X2+(g) + e-.
Third ionisation energy
The third ionisation energy (IE3) is the energy required to remove the third mole of electrons from each +2 ion in a mole of gaseous +2 ions, to form one mole of +3 ions.

Electrons can be removed from an atom until only the nucleus is left. Successive ionisation energies always increase, because each electron is removed from an increasingly positive ion, so the attraction is greater and there is less electron repulsion.

Section 14 of 18Factors affecting ionisation energy

Nuclear charge
Nuclear charge more protons gives stronger attraction;
Atomic radius
Atomic radius the further the electron from the nucleus, the weaker the attraction;
Shielding
Shielding inner electrons repel outer electrons and reduce the attraction to the nucleus.

Section 15 of 18Trends as evidence

Down a group, IE1 decreases: atomic radius increases and shielding increases, outweighing the increased nuclear charge, so the outer electron is held less strongly.

Across a period, IE1 generally increases: nuclear charge increases, the electron is added to the same shell with similar shielding, and radius decreases. There are two dips:

Trends as evidence
from Be to B (and Mg to Al): the outer electron in B is in a 2p subshell, which has higher energy and is further from the nucleus than the 2s, and is slightly shielded by the 2s electrons;
from N to O (and P to S): in O two electrons are paired in one p orbital, and electron-electron repulsion makes one easier to remove.

Successive ionisation energies give evidence for shells. A large jump in successive values shows that the next electron is being removed from an inner shell, closer to the nucleus. For example, a big jump between IE2 and IE3 shows the element is in Group 2, with two outer electrons.

Section 16 of 18Worked examples

  1. Write the equation for the second ionisation energy of magnesium. [2]2 marks

  2. Explain why the first ionisation energy of aluminium is lower than that of magnesium. [2]2 marks

  3. Successive ionisation energies of an element (kJ mol-1): 738, 1451, 7733, 10543. Deduce the group. [2]2 marks

Section 17 of 18Key terms

TermMeaning
Atomic numberThe number of protons in the nucleus of an atom.
Mass numberThe total number of protons and neutrons in the nucleus.
IsotopesAtoms of the same element with different numbers of neutrons.
Relative atomic massThe weighted mean mass of an atom of an element compared with one twelfth of the mass of an atom of carbon-12.
IonAn atom or group of atoms with an overall positive or negative charge.
TermMeaning
Mass spectrometryAn analytical technique that separates ions by mass-to-charge ratio to determine masses and abundances.
m/zThe mass-to-charge ratio of an ion.
Molecular ionThe ion formed when a molecule loses one electron, whose m/z gives the relative molecular mass.
Electrospray ionisationIonisation by forming protonated ions from a sample in solution at high voltage.
Electron impact ionisationIonisation by bombarding vaporised sample with high-energy electrons.
Time of flightThe time taken for an ion to travel through the drift region to the detector.
TermMeaning
ShellA group of orbitals with the same principal quantum number n.
SubshellA set of orbitals of the same type (s, p, d or f) within a shell.
OrbitalA region of space that can hold up to two electrons.
Valence electronAn electron in the outermost shell of an atom.
s-blockElements whose outermost electron occupies an s orbital.
d-blockElements whose highest-energy electrons are in a d subshell.
TermMeaning
First ionisation energyThe energy required to remove one mole of electrons from one mole of gaseous atoms.
Successive ionisation energiesThe energies needed to remove the 1st, 2nd, 3rd ... mole of electrons in turn.
Nuclear chargeThe total positive charge of the protons in the nucleus.
ShieldingThe repulsion of outer electrons by inner-shell electrons, reducing attraction to the nucleus.
Atomic radiusThe distance from the nucleus to the outer electrons.

Section 18 of 18Exam technique

1
Saying the mass number is the number of protons - it is protons plus neutrons.
2
Giving the electron a relative mass of 0 - it is 1/1836 (negligible), not zero.
3
Saying isotopes have different numbers of electrons - they differ in neutrons only.
4
Forgetting to adjust the electron count for ions.
5
Quoting relative atomic mass as a whole number without calculating the weighted mean.
1
Forgetting that the sample must be ionised to positive ions - neutral particles are not accelerated.
2
Stating that heavier ions are faster - all have the same kinetic energy, so lighter ions are faster.
3
Forgetting the extra proton (M + 1) in electrospray ionisation.
4
Omitting the vacuum from a description of the instrument.
5
Taking the tallest peak as the molecular ion - it is the peak with the highest m/z.
1
Removing 3d electrons before 4s when forming transition-metal ions - 4s is lost first.
2
Writing chromium as 3d44s2 and copper as 3d94s2.
3
Forgetting that each orbital holds a maximum of two electrons.
4
Pairing electrons in p orbitals before each orbital has one electron.
5
Confusing a shell with a subshell or an orbital.
1
Omitting state symbols (g) from ionisation energy equations.
2
Saying ionisation energy decreases down a group because nuclear charge decreases - it increases; the radius and shielding effects win.
3
Using 'it is easier to remove an electron from a full shell' for the Group 2/3 dip - the real reason is the p subshell is higher in energy.
4
Forgetting that successive values always increase.
5
Defining ionisation energy per atom rather than per mole.

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