Atoms and reactions

Use these Atoms & Reactions Study Notes to learn atomic structure, electron shells and orbitals, isotopes, ionisation energy, and key reaction types for Section III.

10 min

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.

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

A comparison of the atomic structures of Carbon-12 and Carbon-13. On the left, Carbon-12 is shown with 6 neutrons (red spheres) and 6 protons (purple spheres), with an atomic mass of 12.000. On the right, Carbon-13 has 7 neutrons and 6 protons, with an atomic mass of 13.003. Both atoms have 6 electrons (yellow spheres) orbiting around the nucleus.

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.

Electrons occupy specific energy levels outside the nucleus of an atom.

The arrangement of electrons in these quantised energy levels in an atom is referred to as its electron structure or electronic configuration.

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.

A diagram illustrating the atomic structure, showing a blue nucleus at the center with three concentric circles representing electron orbits. The orbits are labeled with increasing energy levels: n=1, n=2, and n=3. An arrow points outward, indicating increasing energy and distance from the nucleus.

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.

A table displaying the relationship between quantum shell numbers and the corresponding number of electrons. The first column lists quantum shell numbers (n) from 1 to 4, and the second column shows the number of electrons that can occupy each shell: 2 for n=1, 8 for n=2, 18 for n=3, and 32 for n=4.

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.

Illustration showing the 3D shape of an s-orbital, represented as a pink sphere with labeled axes (X, Y, Z) and a nucleus at the center. To the right, a simplified 2D representation of the s-orbital is shown as a plain circle.

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.

Illustration showing three 2p orbitals on the Cartesian plane, labeled X, Y, and Z. The top section displays a combined view of the orbitals, while the bottom section presents the individual shapes of the 2px, 2py, and 2pz orbitals.

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.

A diagram illustrating atomic structure, featuring a central sub-shell represented by a pink sphere, surrounded by various orbitals depicted as shapes in different orientations. The orbitals are labeled with axes (X, Y, Z) and categorized into shells and sub-shells.

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.

A diagram illustrating the shapes of atomic orbitals: s, p, d, and f. The s orbital is shown as a spherical shape at the top, followed by three p orbitals represented as two-lobed shapes. Below are five d orbitals with various lobular configurations, and at the bottom, seven f orbitals with more complex shapes. Each orbital is labeled with its corresponding quantum number and axes are indicated.

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

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.

A table displaying the principal quantum numbers, corresponding sub-shells, and the maximum number of electrons that can occupy each shell. The table includes four rows for quantum numbers 1 to 4, detailing the sub-shells (1s, 2s, 2p, etc.) and their respective electron capacities.

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.

A diagram illustrating energy levels, labeled from Level 1 to Level ∞, with red arrows indicating transitions between levels. Below, a bar graph shows a range from high energy to low energy.

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’.

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.

A graph illustrating the relationship between successive ionisation energy and ionisation energy in kJ mol-1. The graph shows a significant increase in ionisation energy when removing an electron from a quantum shell closer to the nucleus, indicated by a labeled section highlighting three electrons in the outer shell.

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.

A line graph showing the first ionisation energy (in kJ mol-1) plotted against atomic number (Z) for elements from hydrogen (H) to sodium (Na). The graph indicates a large drop in ionisation energy from period 1 to period 2 and from period 2 to period 3, with notable peaks at helium (He) and neon (Ne).

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.

A graph showing the first ionisation energy (in kJ mol⁻¹) plotted against atomic number (Z) for elements from hydrogen (H) to sodium (Na). The graph features peaks and troughs, indicating variations in ionisation energy, with annotations explaining factors affecting electron attraction, such as nuclear shielding and electron-electron repulsion.

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).

Balancing equations is a process to make sure there are the same number of atoms on each side of the symbol equation therefore obeying the law of conservation of mass.

When balancing equations ensure:

the only numbers changes are the coefficients. These are placed in front of the formulae
the formulae are not altered in any way, such as changing to to increase the number of oxygen atoms.

When balancing equations, we must:

identify the number of atoms of each element in the reactants and products
add the appropriate coefficient in front of the formulae to balance the number of atoms on each side
make sure that the final balanced equation has the same number of each type of atom on both sides.

To convert a balanced (stoichiometric) symbol equation into an ionic equation we must:

split compounds composed of aqueous ions into the individual ions, including any multiples
leave intact in the equation any substance that does not contain aqueous ions or is covalently bonded
cancel out spectator ions which appear on both sides of the equation and do not undergo any change.

Addition reactions are a type of chemical reaction where two or more molecules combine to form a single product.

A key example of an addition reaction is electrophilic addition to an alkene.

Elimination reactions remove atoms or small groups from a molecule, resulting in the formation of a double bond.

A key example of an elimination reaction is the elimination of water from an alcohol to form an alkene.

Substitution reactions involve replacing one atom or group with another.

Key examples of substitution reactions are electrophilic substitution on an aromatic ring or nucleophilic substitution on a haloalkane or alcohol.

Oxidation reactions involve:

The loss of electrons
Increase in oxidation state
Addition of oxygen (sometimes)
Removal of hydrogen (sometimes).

where:

= oxidising agent

Formulae for oxidation reactions may include the symbol for an oxidising agent, .

Reduction reactions involve:

The gain of electrons
Decrease in oxidation state
Removal of oxygen (sometimes)
Addition of hydrogen (sometimes).

where:

= reducing agent

Formulas for reduction reaction may include the symbol for a reducing agent, .

Hydrolysis reactions involve the breakdown of a compound by the addition of water, resulting in two products.

A key example of hydrolysis is the breakdown of an ester into a carboxylic acid and alcohol by water.

The term hydrolysis is also applied to other reactions where water is responsible for the breaking of a bond. In the hydrolysis of halogenoalkanes the water is a nucleophile breaking the bond.

Polymerisation reactions are reactions in which monomers are chemically bonded to form polymers.

where:

represents the monomer
is the number of monomers that combine to form the polymer
There are two types of polymerisation:

Addition polymerisation:

Involves unsaturated monomers that contain double bonds. The double bonds open up, allowing the monomers to join together to form a long chain polymer.

Condensation polymerisation:

Involves monomers that contain more than two functional groups. When these monomers react, a small molecule is eliminated as a by-product.

Free radical substitution of alkanes to produce haloalkanes.

There are three steps: initiation, propagation, and termination.

Diagram illustrating the free radical substitution mechanism, including initiation, propagation, and termination steps. The process involves chlorine gas (Cl2) and methane (CH4) under UV light, showing the formation of radicals and products like HCl and chlorinated hydrocarbons.

Initiation
Formation of halogen free radicals from a non radical molecule.
Conditions required: UV light.

Propagation
Step 1: The halogen radical reacts with the alkane, to produce an alkyl radical and .
Step 2: The alkyl radical then reacts with the halogen to produce a haloalkane and regenerates another halogen radical.

Termination
Two radicals combine, resulting in a non radical molecule being formed and no further reaction.

Atoms and reactions

On this page

The structure of the atom

Orbitals

Stoicheometric equations

Types of reaction