Atomic structure and the periodic table (Topic 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.

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.

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:

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 for a larger molecule is complex because the molecule breaks into many fragments during the process.

The last major peak in the spectrum is called the molecular ion peak ( peak). This peak represents the molecular ion, which has a single positive charge and provides the relative molecular mass, of the compound.

The molecular ion peak is found to the far right of the spectrum and should not be confused with the highest peak, the base peak which is linked to the most stable fragment ion.

The spectrum for ethanol () shows a molecular ion peak at produced by .

Ionisation energy relates to the energy required to remove an electron from an atom or ion. The analysis of successive ionisation energies of an element provides evidence for the existence of quantum shells.

Ionisation energy increases with successive ionisations as the attraction of the outermost electron to the nucleus increases.

The increase in ionisation energy with successive ionisation is not steady.

This indicates that the reduced ionic radius and reduced electron shielding is more significant for some successions.

Elements in the same group show the largest jumps between the same successions; the elements have the same number of electrons in the highest energy level.

The change in the quantum shell the electron is being removed from accounts for this significant increase in ionisation energy.

The existence of quantum shells is supported by data from the first ionisation energies of elements of increasing atomic number.

Large drops in first ionisation energy are seen as you move between periods of the periodic table.

Electrons in higher energy quantum shells are further from the nucleus and experience more nuclear shielding, reducing the effective nuclear attraction. This reduces the energy required to remove an electron.

There is a general trend of increasing first ionisation energy across a period.

Ignoring subshells, as you move across the period you expect the nuclear charge to increase, causing the atomic radius to decrease, and the attraction to the outermost electron to increase; the first ionisation energy should increase.

Exceptions in the general trend can be explained by the existence of subshells.

The drop in first ionisation energy from group 2 to group 3 (13) is due to the p subshell sitting at a higher energy level than the full s subshell. It is easier to remove an electron from the p subshell.

When subshells are half-filled each atomic orbital within the subshell is singly occupied. Moving along the period, following a half-filled subshell, a doubly occupied orbital is introduced. The impact of the increased electron–electron repulsion is greater than that of the increase in nuclear charge resulting in a weaker attraction and a drop in first ionisation energy from group 5 (15) to group 6 (16).

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.

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.

The maximum number of electrons that can fill a quantum shell is given by the formula e = 2n²

The number of electrons needed to fully occupy the first four quantum shells is shown below.

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.

The number of subshells per shell in an atom increases with the increase in principal quantum number.

The number of electrons per shell can be calculated by considering full occupancy of all the atomic orbitals present.

Within a shell, the atomic orbitals are grouped as subshells, distinguished by the letters s,p,d and f.

The s subshell has one atomic orbital, the p subshell has three orbitals, the d subshell has five orbitals while the f subshell has a total of seven atomic orbitals.

As the energy level of the shell increases, the size of the atomic orbitals within the subshells increases but their shapes remain similar.

An atomic orbital has a three-dimensional shape indicating the area in which an electron is likely to be found.

The s orbital has a symmetrical, spherical shape, with the nucleus located at the centre.

The s orbitals of all shells have the same spherical shape but differ in the size of their radius.

The p orbitals are described as dumb-bell shaped, with the nucleus located between the two halves.

There are three degenerate p orbitals in each p subshell. These orbitals have an identical shape and energy, however, their orientation in space differs. They lie at right angles to one another along the X, Y, and Z axes of a Cartesian plane.

The three p orbitals are thus distinguished as p_x, p_y, and p_z.

Box notation goes beyond subshell notation as it describes which atomic orbitals contain electrons. Each atomic orbital is drawn as a box and each electron as an arrow.

All electrons are identical negatively charged particles. Two electrons close to each other experience strong electrostatic repulsions.

To minimise electron–electron repulsion, when filling degenerate orbitals (such as three 2p orbitals), electrons first singly occupy each orbital before pairing occurs.

Two electrons in the same orbital must be in opposite spin (clockwise and anticlockwise) denoted by an upwards (↑) and a downwards (↓) pointing arrow.

This minimises the repulsive effect between the two electrons.

The electronic configuration of oxygen is 1s² 2s² 2p⁴.

Common mistakes to avoid using box notation are:

• doubly filling degenerate orbitals from left to right without singly filling first
• pairing electrons with the same spin

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.

Ions are formed by the loss or gain of electrons by the atom of an element.

The electronic configuration of an ion can be written using the same set of rules as that for the element and then adjusting for the increase or decrease in electron count.

For ionisation, it is important to adjust from the element rather than filling the atomic orbitals directly using the ions’ electron count. This is because, when filled, the 4s orbital becomes higher in energy than the 3d orbital.

Electrons must be taken from the 4s subshell before being removed from the 3d subshell.

Vanadium 2+, , and scandium, , both have an electron count of 21 but their electron configuration is different.

The existence of quantum shells was supported by experimental evidence involving characteristic atomic emission spectra of individual elements.

When electrons are excited, they move up an energy level. When they fall back to their original energy level they emit radiation at specific wavelengths, producing a line spectrum unique to each element.

These discrete spectral lines reveal that only certain energy transitions are possible; electrons in an atom must occupy quantized energy levels. These are the ‘electron shells’.

The periodic table is a fundamental tool in chemistry that organises the elements based on their atomic number, creating a systematic arrangement that reveals patterns in their properties and reactivity.

These repeating patterns are known as periodicity.

The periodic table is divided into rows called periods.

There are a total of seven periods.

Elements in the same period have the same number of electron shells.

Columns in the periodic table are known as groups.

There are 18 groups in the periodic table.

Elements within the same group have similar chemical properties because they have the same number of outer electrons.

All of the elements in the periodic table belong to blocks based on their electron configuration.

s-block: the outermost electron is in the s-subshell; for example, the electron configuration for potassium is

p-block: the outermost electron is in the p-subshell; for example, the electron configuration for aluminium is

d-block: the outermost electron is in the d-subshell; for example, the electron configuration for vanadium is

f-block: the outermost electron is in the f-subshell; for example, the electron configuration for cerium is

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.

Although across a period, the general trend is that the first ionisation energy increases, there are some exceptions.

There is a drop in first ionisation energy between Group 2 and 3, and . This is linked to the outermost electron sitting in the p-subshell, which is at a higher energy level than the s-subshell.

There is another drop in first ionisation energy between Group 5 and 6, and . This is linked to electron–electron repulsion in the half filled p-subshell. The fourth electron in the p-subshell is paired in an orbital and so experiences more electron-electron repulsion than the unpaired p-subshell electrons in group five elements.

The general trend across periods is an increase in melting point from Group 1 to Group 4. There is then a sharp decrease in melting point between Group 4 and Group 5.

The melting points from Group 5 onwards are comparatively low.

The trend exists due to the transition from giant lattices to simple molecules, held together by weak intermolecular forces.

There is an increase in melting point between phosphorus () and sulfur () in period 3. This is due to the increased molecular size and therefore the number of electrons is higher in compared to , creating to more significant temporary dipoles. This leads to stronger intermolecular forces that require more energy to overcome.

The atomic radius is taken from atoms in their most stable elemental form. It gives information about the distance between the nucleus and the outer shell electrons.

Across periods in the periodic table, atomic radii decrease from left to right.

The increasing atomic number across the period, and therefore the increased nuclear charge, attracts the electrons more tightly to the nucleus.

Electrons are added to the same electron shell, so there is no significant increase in nuclear shielding across the period. This means the stronger nuclear attraction directly reduces the atomic radius.