Atomic Structure and Isotopes (OCR A-Level Chemistry A): Revision Notes

Atomic Structure and Isotopes

The nuclear atom

All matter is made up of atoms. At GCSE level, you learned about the nuclear model of the atom, which describes the structure of atoms in terms of subatomic particles.

An atom consists of:

• A central nucleus containing two types of subatomic particles: protons and neutrons
• Electrons occupying regions called shells that surround the nucleus

Properties of subatomic particles

Atoms and their subatomic particles have extremely small masses. Rather than working in grams, chemists find it more convenient to compare the masses of subatomic particles using relative masses.

Mass

The relative masses of the three subatomic particles are:

• A proton has a relative mass of 1
• A neutron has virtually the same mass as a proton, with a relative mass of 1
• An electron has negligible mass, approximately $\frac{1}{1836}$ the mass of a proton

While accurate measurements show that a neutron has a slightly greater mass than a proton (by a factor of 1.001375), this difference is so small that chemists usually assume protons and neutrons have identical masses. The tiny mass of electrons means they contribute almost nothing to the overall mass of an atom.

Charge

Subatomic particles also carry electric charge:

• A proton has a positive charge with a relative charge of $1+$
• An electron has a negative charge with a relative charge of $1-$
• A neutron is neutral and has no charge (relative charge of $0$)

The charge on a proton is equal in magnitude but opposite in sign to the charge on an electron. In actual values, a proton carries a charge of $+1.602177 \times 10^{-19}$ C (coulombs), while an electron carries $-1.602177 \times 10^{-19}$ C. However, it's much easier to work with relative charges of $1+$ and $1-$.

Summary of subatomic particles

Particle	Abbreviation	Relative charge	Relative mass
proton	$\text{p}^+$	$1+$	$1$
neutron	$\text{n}$	$0$	$1$
electron	$\text{e}^-$	$1-$	$\frac{1}{1836}$

Building atoms

Several important principles govern how atoms are constructed:

Role of neutrons:

Neutrons act as the "glue" holding the nucleus together. Without neutrons, the electrostatic repulsion between positively charged protons would cause the nucleus to break apart. Most atoms contain the same number of neutrons as protons, or slightly more. As the nucleus gets larger, increasingly more neutrons are needed to hold it together stably.

Atomic number - the identity of an element

The atomic number (also called the proton number, symbol $Z$) is the number of protons in the nucleus of an atom. This is the most important property of an atom because it determines which element the atom belongs to.

Key facts about atomic number:

• Every atom of the same element contains the same number of protons
• Different elements contain atoms with different numbers of protons
• The periodic table lists elements in order of increasing atomic number
• Each element is assigned its own unique atomic number

The periodic table shows each element labeled with its atomic number, allowing you to quickly determine the number of protons (and therefore electrons) in any atom.

Isotopes

What are isotopes?

While every atom of an element must have the same number of protons, the number of neutrons can vary. This gives rise to isotopes.

The diagram above shows two isotopes of hydrogen:

• Normal hydrogen has 1 proton, 0 neutrons, and 1 electron
• Deuterium has 1 proton, 1 neutron, and 1 electron

Both are isotopes of hydrogen because they both have 1 proton (atomic number 1), but they differ in their neutron count and mass.

Representing isotopes

Chemists use a special notation to represent different isotopes, showing both the mass number and atomic number.

The notation includes:

• Mass number ($A$, also called nucleon number): written as a superscript to the upper left of the element symbol
• Atomic number ($Z$, also called proton number): written as a subscript to the lower left of the element symbol
• The chemical symbol in the middle

The relationships between these numbers are:

$A = \text{number of protons} + \text{number of neutrons}$

$Z = \text{number of protons}$

To find the number of neutrons in an isotope:

$\text{number of neutrons} = A - Z$

Example: oxygen isotopes

Oxygen has three naturally occurring isotopes: oxygen-16, oxygen-17, and oxygen-18.

Alternative ways of writing isotopes

You may encounter isotopes written in different formats:

• Full notation: $^{16}_8\text{O}$
• Mass number only: $^{16}\text{O}$ or $\text{O-16}$ or simply "oxygen-16"

Isotopes and chemical reactions

However, isotopes do have slightly different physical properties:

• Higher-mass isotopes have higher melting points, boiling points, and densities
• These differences arise from the greater mass but don't affect chemical reactivity

Heavy water - a practical example

An interesting example of isotopes in practice is heavy water. In normal water ($\text{H}_2\text{O}$), nearly all hydrogen atoms are the $^1\text{H}$ isotope. In heavy water, all hydrogen atoms are replaced with deuterium ($^2\text{H}$, also symbolized as D), giving the formula $\text{D}_2\text{O}$.

Comparing normal and heavy water:

Physical property	Normal water ($\text{H}_2\text{O}$)	Heavy water ($\text{D}_2\text{O}$)
Melting point / °C	0.00	3.80
Boiling point / °C	100.00	101.40
Density / g cm⁻³	1.00	1.11

Tritium ($^3\text{H}$ or T) is a third isotope of hydrogen containing two neutrons, which forms "super-heavy water" ($\text{T}_2\text{O}$).

Ions

What are ions?

An ion is a charged atom in which the number of electrons is different from the number of protons.

There are two types of ions:

Cations (positive ions):

• Atoms with fewer electrons than protons
• Have an overall positive charge
• Formed when atoms lose electrons
• Examples: $\text{Na}^+$, $\text{Mg}^{2+}$, $\text{Al}^{3+}$

Anions (negative ions):

• Atoms with more electrons than protons
• Have an overall negative charge
• Formed when atoms gain electrons
• Examples: $\text{Cl}^-$, $\text{O}^{2-}$, $\text{N}^{3-}$

Writing ions

Ions are always shown with their overall relative charge written as a superscript to the right of the element symbol. For example:

• $^{24}_{12}\text{Mg}^{2+}$ represents a magnesium ion with a $2+$ charge
• $^{35}_{17}\text{Cl}^-$ represents a chlorine ion with a $1-$ charge

Atomic structure of ions

Let's examine the composition of two common ions:

Remember!