Atomic structure (3.1.1)

The structure of the atom:

The nucleus contains the protons and neutrons.

The nucleus is surrounded by shells (orbitals) which contain the electrons.

Subatomic particles can be described by their relative charge and relative mass.

Dalton (1800s):

• discovered that atoms can be combined to make compounds
• believed atoms cannot be broken into smaller parts
• disproved by J.J. Thomson’s cathode ray tube experiments.

Thomson (1897):

• discovered the electron
• believed that atoms were a sphere of positive mass embedded with electrons: the plum pudding model
• was disproved by Rutherford’s gold foil experiment.

Rutherford (1905):

• discovered the nucleus with his gold foil experiment.

Rutherford’s gold foil experiment:

• A very thin piece of gold foil was bombarded with alpha particles.
• Most particles went straight through (A).
• Some were deflected (B) and a few were reflected (C).
• This proved that the vast majority of the mass of an atom is concentrated in a small central volume and that the central mass was positively charged.

Bohr (1913):

• discovered energy levels
• related atomic spectra to electronic energy levels.

The atomic number represents the number of protons in the nucleus of an atom.

It is a unique identifier for each element and is a fundamental factor in the organisation of the periodic table.

Elements are arranged in increasing order of atomic number from left to right.

Isotopes are atoms of the same element with different numbers of neutrons and different masses.

Carbon-12 and carbon-13 are isotopes of carbon.

• Both identify as carbon, with an atomic number of six.
• Carbon-12 has six neutrons and a relative atomic mass of 12 whereas carbon-13 has seven neutrons and a relative atomic mass of 13.

A simple time of flight (TOF) mass spectrometer is used to identify the mass of ions.

1. A sample is vapourised.
2. The vapourised sample is bombarded with high speed electrons that knock out an electron in the molecule, causing its ionisation.
3. An electric field is applied, which accelerates each ion toward a drift tube.
4. The ions all drift with the same kinetic energy, but their velocities differ depending on their mass-to-charge ratios (m/z). Ions with a smaller m/z travel faster and reach the detector first.
5. The mass spectrometer detects the ions’ time of flight.
6. An electric current proportional to the relative abundance of ions in the beam produces a peak on the mass spectrum at the time it was detected.

The raw data for a mass spectrum is the signal intensity paired with the time of flight.

The data is processed using information from calibration.

Signal intensity can be converted to relative abundance. This sits on the axis.

The conversion is not always normalised, (made so the peaks sum to ), so it is important to check.

Time of flight is converted to mass to charge ratio (m/z). This sits on the axis.

The mass spectrum below shows the relative abundance of the two isotopes of boron.

Isotope 1: m/z , relative abundance
Isotope 2: m/z , relative abundance

The relative atomic mass of boron can be calculated from this data:

Mass spectra of organic molecules will exhibit a tiny peak at 1 mass unit greater than the molecular ion.

This is called the M+1 peak and is caused by the presence of the carbon-13 isotope which comprises of all carbon atoms.

The of propanone, , is 58.

In the spectrum of propanone, the molecular ion peak, , is at 58 and a smaller M+1 peak at 59 shows the proportion of propanone featuring carbon-13.

Principal energy levels or principal quantum shells (n) are numbered based on their distance from the nucleus.

The first principle quantum shell is closest to the nucleus and described as n = 1.

n is an integer value which increases as the shells move further from the nucleus.

Shells are divided into subshells which can be further divided into atomic orbitals.

The atomic orbital is a three-dimensional region around the nucleus of an atom, where there is a maximum probability of finding an electron.

Each orbital can have a maximum of 2 electrons, with opposite spins.

Make use of your periodic table to remember the order in which the subshells are filled.

List the subshells in order as you move in increasing atomic number from hydrogen to the element of interest.

The energy levels of 3d and 4s subshells overlap.

The energy of 4s lies slightly below that of 3d therefore 4s is filled before 3d.

However, once filled, the energy of 3d falls below the 4s energy level.

When a transition metal is ionised, the 4s electrons are easier to remove than the 3d electrons.

Each period of the periodic table starts with one electron in a new electron shell.

The number before the orbital indicates the shell number, which also corresponds to the period number.

In period 2, the 2s subshell is filled first with two electrons.

After the 2s subshell is filled, electrons then occupy the 2p subshell, which can hold up to six electrons.

The trend in electron configuration across period 2 is shown below.

In the period 3, the 3s subshell is filled first with two electrons.

Once the 3s subshell is filled, electrons begin to occupy the 3p subshell, which can hold up to six electrons.

The trend in electron configuration across period 3 is shown below.

Successive ionisation energies for an unknown period 3 element are shown below:

There is a steady rise between the first ionisation energy and second ionisation energy for the unknown element.

Between the second ionisation energy and third ionisation energy, there is comparatively a sharp rise indicating that the third electron is being removed from a new energy level.

From this information, we can therefore deduce that element X is in Group 2.

As it is in period 3, element X must be magnesium.