Atomic Structure (HSC SSCE Chemistry): Revision Notes

Atoms and Isotopes

What is an atom?

An atom is the smallest unit of matter that retains the properties of an element. Our modern understanding of atoms developed throughout the 19th and 20th centuries, building on the work of pioneering scientists like Lavoisier, Dalton, Thomson, Rutherford, Bohr, and many others.

Early chemists in the 18th and 19th centuries thought of atoms as solid, indivisible spheres. This simple model helped explain chemical reactions and compound formation, but it couldn't answer important questions like why atoms bond together or why some substances are more reactive than others.

Basic structure of an atom

An atom consists of two main regions:

1. The nucleus - an extremely small, dense core at the centre
2. The electron cloud - a large region surrounding the nucleus

The nucleus is the tiny central core of the atom. Although it's incredibly small, it contains almost all of the atom's mass (over 99.95%). The nucleus carries a positive electrical charge.

The electron cloud is the region around the nucleus where electrons move. This cloud makes up most of the atom's volume, but contributes very little to its mass. The electron cloud carries a negative electrical charge.

Understanding the size of an atom

To understand just how small the nucleus is compared to the whole atom, imagine this:

If the nucleus were the size of a cherry stone placed at the centre of a large stadium, the electrons would be like tiny fruit flies buzzing around anywhere in the stadium - on the field, in the stands, even at the very back. But unlike fruit flies, electrons move incredibly fast, able to zip from near the nucleus to the outer edge of the atom in the tiniest fraction of a second.

The term 'electron cloud' describes how electrons move randomly through the space around the nucleus. Because electrons are so small and have so much room to move, they never collide with each other despite their random motion.

Subatomic particles

Atoms are made of three types of smaller particles called subatomic particles:

Protons

A proton is a positively charged particle found in the nucleus. Each proton has:

• A charge of +1 (relative charge)
• A mass approximately equal to the mass of a hydrogen atom (set as 1 on the relative mass scale)

Neutrons

A neutron is a neutral particle (no charge) also found in the nucleus. Each neutron has:

• A charge of 0
• A mass approximately equal to that of a proton (1 on the relative mass scale)

Electrons

An electron is a negatively charged particle found in the electron cloud surrounding the nucleus. Each electron has:

• A charge of -1 (equal in magnitude but opposite to a proton's charge)
• A mass of only $\frac{1}{2000}$ that of a proton (electrons are extremely light)

Properties comparison

Particle	Symbol	Relative Mass	Relative Charge	Location
Electron	$e$ or $^0_{-1}e$	$\frac{1}{2000}$	-1	Electron cloud
Proton	$p$ or $^1_1p$	1	+1	Nucleus
Neutron	$n$ or $^1_0n$	1	0	Nucleus

Composition of simple atoms

Hydrogen - the simplest atom

The simplest atom is hydrogen. Its nucleus contains just one proton (and no neutrons), with one electron in the electron cloud surrounding it.

Building up the elements

All other elements are formed by adding more protons to the nucleus. When we add protons, we must also add electrons to keep the atom neutral. The number of neutrons varies, but for lighter elements, the number of neutrons is approximately equal to the number of protons.

Here's the composition of the first ten elements:

Element	Protons	Neutrons	Electrons	Mass Number
H	1	0	1	1
He	2	2	2	4
Li	3	4	3	7
Be	4	5	4	9
B	5	6	5	11
C	6	6	6	12
N	7	7	7	14
O	8	8	8	16
F	9	10	9	19
Ne	10	10	10	20

Atomic number and mass number

Two important numbers describe every atom:

Atomic number (Z)

The atomic number ($Z$) is the number of protons in the nucleus of an atom.

• The atomic number determines what element the atom is
• Each element has a unique atomic number
• For example: hydrogen has $Z = 1$, carbon has $Z = 6$, sodium has $Z = 11$

Mass number (A)

The mass number ($A$) is the total number of protons plus neutrons in the nucleus.

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

• For helium: $A = 2 + 2 = 4$
• For carbon: $A = 6 + 6 = 12$
• For sodium: $A = 11 + 12 = 23$

Calculating the number of neutrons

If you know the atomic number and mass number, you can work out the number of neutrons:

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

Isotopes

What are isotopes?

For most elements, not all atoms are identical. Some atoms of the same element have different numbers of neutrons in their nuclei.

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons in their nuclei.

Because isotopes of the same element have the same number of protons, they have:

• The same atomic number ($Z$)
• Different mass numbers ($A$)
• The same chemical properties
• Very similar physical properties

Isotope notation

We represent isotopes using the notation:

$^A_Z M$

where:

• $M$ is the chemical symbol for the element
• $A$ is the mass number (top number)
• $Z$ is the atomic number (bottom number)

Examples of isotopes

Even the simplest elements have isotopes:

• Hydrogen has three: $^1_1\text{H}$ (normal hydrogen), $^2_1\text{H}$ (deuterium), and $^3_1\text{H}$ (tritium)
• Oxygen has three: $^{16}_8\text{O}$, $^{17}_8\text{O}$, and $^{18}_8\text{O}$
• Carbon has three: $^{12}_6\text{C}$, $^{13}_6\text{C}$, and $^{14}_6\text{C}$

Relative abundance

The relative abundance of an isotope is the percentage of that isotope present in a naturally occurring sample of the element.

Common isotopes and their relative abundances

Element	Isotope	Mass Number	Relative Abundance (%)
Hydrogen	$^1\text{H}$	1	99.99
	$^2\text{H}$	2	0.01
Boron	$^{10}\text{B}$	10	19.6
	$^{11}\text{B}$	11	80.4
Carbon	$^{12}\text{C}$	12	98.89
	$^{13}\text{C}$	13	1.11
Magnesium	$^{24}\text{Mg}$	24	78.7
	$^{25}\text{Mg}$	25	10.1
	$^{26}\text{Mg}$	26	11.2
Chlorine	$^{35}\text{Cl}$	35	75.5
	$^{37}\text{Cl}$	37	24.5
Copper	$^{63}\text{Cu}$	63	69.1
	$^{65}\text{Cu}$	65	30.9

Unstable isotopes and radioactivity

What is radioactivity?

Some isotopes are radioactive, meaning they spontaneously emit radiation. Radioactivity is the spontaneous emission of radiation that occurs with certain isotopes because these isotopes are unstable.

Radioisotopes

Radioactive isotopes (or radioisotopes) are isotopes that spontaneously emit radiation. They are also called unstable isotopes.

Scientists describe nuclei as either:

• Stable nuclei - do not emit radiation
• Unstable nuclei - radioactive, emit radiation

Types of radiation

Radioactive substances can emit three different types of radiation, named before their true identities were discovered:

Alpha (α) particles

Alpha particles are helium nuclei (two protons and two neutrons stuck together). They have:

• Symbol: $\alpha$ or $^4_2\text{He}$
• Charge: +2 (two protons)
• Mass: 4 (2 protons + 2 neutrons)
• Penetrating power: Low - stopped by a sheet of paper or a few centimetres of air

Beta (β) particles

Beta particles are high-speed electrons emitted from the nucleus. They have:

• Symbol: $\beta$ or $^0_{-1}e$
• Charge: -1
• Mass: $\frac{1}{2000}$ (very light)
• Penetrating power: Moderate - can pass through paper and up to 0.5 mm of aluminium, but stopped by 0.5 mm of lead

Gamma (γ) rays

Gamma rays are a type of electromagnetic radiation (like X-rays, light, or radio waves, but with much shorter wavelengths). They have:

• Symbol: $\gamma$
• Charge: 0 (no charge)
• Mass: 0 (electromagnetic radiation has no mass)
• Penetrating power: High - very penetrating; require several centimetres of lead or 15 cm of concrete to stop them

Properties summary

Emission	Symbol	Identity	Relative Charge	Relative Mass	Penetrating Power
Alpha	$\alpha$, $^4_2\text{He}$	Helium nucleus	+2	4	Low
Beta	$\beta$, $^0_{-1}e$	Electron	-1	$\frac{1}{2000}$	Moderate
Gamma	$\gamma$	Electromagnetic radiation	0	0	High

Nuclear equations

What are nuclear equations?

Nuclear equations show what happens when a radioactive nucleus disintegrates. They show:

• The original nucleus
• The particle emitted
• The new nucleus formed

Alpha decay

When a nucleus emits an alpha particle, it loses 2 protons and 2 neutrons.

Beta decay

When a nucleus emits a beta particle (electron), one of its neutrons transforms into a proton and an electron:

$^1_0n \rightarrow ^1_1p + ^0_{-1}e$

The electron is ejected from the nucleus, while the proton remains. This means:

• The atomic number increases by 1 (one more proton)
• The mass number stays the same (total nucleons unchanged)

Worked examples

Why are some isotopes unstable?

Inside the nucleus, there are two competing forces:

1. Electrostatic repulsion between positively charged protons (pushing them apart)
2. Mass-mass attraction between all nucleons (pulling them together)

Instability in large nuclei

For large nuclei (atomic mass greater than about 80), there are too many particles for the short-range attractive forces to overcome the electrostatic repulsions. This makes these nuclei unstable.

Instability in lighter nuclei

For lighter nuclei, stability depends on the ratio of neutrons to protons (n:p ratio). There's a "zone of stability" - a range of n:p ratios that produces stable nuclei.

• If the n:p ratio is too high (too many neutrons), the isotope is a beta emitter
• If the n:p ratio is too low (too few neutrons), the isotope is an alpha emitter

Half-life

What is half-life?

Radioisotopes decay at different rates. Some decay quickly, others very slowly. We measure this using half-life.

The half-life of a radioisotope is the time required for half the atoms in a sample to undergo radioactive decay.

Understanding half-life with an example

Common radioisotopes and their half-lives

Name	Radiation Emitted	Half-life
Cobalt-60	$\beta$, $\gamma$	5.3 years
Iodine-131	$\beta$, $\gamma$	8 days
Plutonium-239	$\alpha$, $\gamma$	$2.4 \times 10^4$ years (24,000 years)
Radium-226	$\alpha$, $\gamma$	$1.6 \times 10^3$ years (1,600 years)
Sodium-24	$\beta$, $\gamma$	15 hours
Uranium-238	$\alpha$, $\gamma$	$4.5 \times 10^9$ years (4.5 billion years)

Natural and human-made radioisotopes

Natural radioisotopes occur in nature, such as:

• Uranium isotopes
• Radium isotopes
• Carbon-14
• Hydrogen-3 (tritium)

Human-made radioisotopes are produced by:

• Nuclear reactors (as by-products or by deliberately placing materials in reactors)
• Particle accelerators like cyclotrons (by bombarding stable isotopes)

Remember!