The Atomic World (VCE SSCE Chemistry): Revision Notes

The Atomic World

Introduction to atomic theory

Scientists have developed a deep understanding of the structure of atoms, which are the fundamental building blocks of all matter. Because atoms are incredibly small—too small to see even with the most powerful optical microscopes—much of our knowledge about atoms comes from theoretical models and indirect observations.

Dalton's atomic theory

In $1802$, English scientist John Dalton presented the first comprehensive atomic theory of matter. Dalton proposed that all matter is made up of tiny spherical particles called atoms, which he believed were indivisible (could not be divided) and indestructible (could not be destroyed).

Dalton accurately described elements as materials containing just one type of atom, and compounds as materials containing different types of atoms in fixed ratios. This means that, for example, pure oxygen contains only oxygen atoms, whilst water contains hydrogen and oxygen atoms always in the same proportion.

Viewing atoms

Dalton's atomic theory assumed that atoms are spherical in shape. However, there was no way to confirm this assumption until $1981$, when IBM researchers Gerd Binnig and Heinrich Rohrer developed the scanning tunnelling microscope (STM). This revolutionary technology allowed scientists to visualize individual atoms for the first time, confirming that atoms are indeed spherical.

These images show copper atoms and silicon atoms on a silicon chip surface, as visualized using scanning tunnelling microscopy. The coloured representations help us understand the arrangement and structure of atoms in materials.

Structure of atoms

Overview

Atoms consist of a small, positively charged nucleus at the centre, surrounded by a much larger cloud of negatively charged electrons. The nucleus itself is made up of two types of subatomic particle: protons (which are positively charged) and neutrons (which have no charge).

This simplified model shows a helium atom with two protons and two neutrons in the central nucleus, and two electrons orbiting around it. Whilst the diagram shows electrons in fixed orbits, evidence suggests they actually move throughout an area more like a cloud surrounding the nucleus.

Electrons

Electrons are negatively charged particles that form a cloud of negative charge around the nucleus. This electron cloud gives the atom its overall size and volume. Each electron is approximately $1800$ times smaller than a proton or neutron, which means electrons contribute very little to the total mass of an atom. However, the space occupied by the electron cloud is $10\,000$ to $100\,000$ times larger than the nucleus itself.

Electrons are held to the nucleus by electrostatic attraction—the force between opposite charges. Negative particles attract positive particles, so the negatively charged electrons are attracted to the positively charged protons in the nucleus. The charge on an electron is equal but opposite to the charge on a proton. We say electrons have a charge of $-1$, whilst protons have a charge of $+1$.

The electricity that powers lights and household appliances results from electrons moving through wires as an electric current. Sparks and lightning are also caused by electrons moving through air.

The nucleus

The nucleus of an atom is approximately $10\,000$ to $100\,000$ times smaller than the overall size of the atom. To put this in perspective, if an atom were the size of the Melbourne Cricket Ground, the nucleus would be about the size of a pea in the centre.

The subatomic particles in the nucleus—the protons and neutrons—are referred to together as nucleons. Protons are positively charged particles with a mass of approximately $1.673 \times 10^{-27}$ kg. Neutrons are almost identical in mass to protons, with a mass of approximately $1.675 \times 10^{-27}$ kg, but they have no electrical charge.

The table below summarises the properties of all three subatomic particles:

Particle	Symbol	Charge	Mass relative to a proton	Mass (kg)
proton	p	$+1$	$1$	$1.673 \times 10^{-27}$
neutron	n	$0$	$1$	$1.675 \times 10^{-27}$
electron	e	$-1$	$\frac{1}{1800}$	$9.109 \times 10^{-31}$

Rutherford's gold foil experiment

Between $1899$ and $1911$, physicist Ernest Rutherford conducted experiments that revealed atoms are mostly empty space. He fired a beam of alpha particles (helium nuclei) at an extremely thin piece of gold foil. You can imagine this experiment as throwing pebbles at an object in the dark—by listening to whether the pebbles hit something or pass through, you can work out the object's shape.

Element symbols

Atoms can be identified by how many protons they contain. An element is made up of atoms that all contain the same number of protons in their nucleus. Scientists have discovered $118$ different elements, with about $98$ of these occurring naturally. The remaining elements have only been observed in laboratory conditions.

Each element has a name and a unique chemical symbol. The chemical symbol consists of one or two letters, where the first letter is always capitalised and any subsequent letters are always lowercase.

Some common elements and their symbols include:

Element	Symbol	Element	Symbol
aluminium	Al	mercury	Hg
argon	Ar	nitrogen	N
carbon	C	oxygen	O
chlorine	Cl	potassium	K
copper	Cu	silver	Ag
hydrogen	H	sodium	Na
iron	Fe	uranium	U

In many cases, the chemical symbol corresponds to the element's name. For example, nitrogen has the symbol N, chlorine has the symbol Cl, and uranium has the symbol U.

Representing atoms

Atomic number

The number of protons in an atom's nucleus is known as the atomic number and is represented by the symbol $Z$.

All atoms that belong to the same element must have the same number of protons and therefore the same atomic number. For example, all hydrogen atoms have $Z = 1$, all carbon atoms have $Z = 6$, and all gold atoms have $Z = 79$.

Because all atoms are electrically neutral (have no overall charge), the number of electrons in an atom equals the number of protons. Therefore, the atomic number tells you both the number of protons and the number of electrons in a neutral atom. For example, carbon atoms with $Z = 6$ have six protons and six electrons.

Mass number

The total number of protons and neutrons in the nucleus is known as the mass number and is represented by the symbol $A$. The mass number represents the total mass of the nucleus, because electrons contribute negligibly to the atom's mass.

Nuclide notation

A standard way of representing an atom shows its atomic number and mass number in nuclide notation:

$^{A}_{Z}\text{X}$

where:

• $A$ = mass number (top)
• $Z$ = atomic number (bottom)
• X = element symbol

For example, an aluminium atom would be written as:

$^{27}_{13}\text{Al}$

From this representation, you can determine:

• The number of protons is $13$ (the atomic number, $Z$)
• The number of neutrons is $14$ (calculated as $A - Z = 27 - 13 = 14$)
• The number of electrons is $13$ (equals the number of protons in a neutral atom)

Calculating subatomic particles

Isotopes

All atoms belonging to the same element have the same number of protons in the nucleus and therefore the same atomic number, $Z$. However, not all atoms of the same element have the same mass number, $A$. This is because atoms of the same element can have different numbers of neutrons.

Isotopes are atoms that have the same number of protons (same atomic number) but different numbers of neutrons (and therefore different mass numbers). In other words, isotopes are different forms of the same element.

Hydrogen isotopes

Hydrogen provides a clear example of isotopes. Hydrogen atoms can have a mass number of $1$, $2$, or $3$. This means hydrogen atoms may contain just a single proton, a proton and a neutron, or a proton and two neutrons.

The three isotopes of hydrogen are given special names:

• Hydrogen-1 ($^{1}_{1}\text{H}$) or protium: one proton, no neutrons
• Hydrogen-2 ($^{2}_{1}\text{H}$) or deuterium: one proton, one neutron
• Hydrogen-3 ($^{3}_{1}\text{H}$) or tritium: one proton, two neutrons

All three isotopes have one proton and one electron, which means they are all hydrogen and have identical chemical properties.

Carbon isotopes

Carbon has three naturally occurring isotopes: carbon-$12$, carbon-$13$, and carbon-$14$. These can be represented as:

${}^{12}_{6}{C} \quad {}^{13}_{6}{C} \quad {}^{14}_{6}{C}$

All three carbon isotopes have six protons and six electrons, but they differ in their number of neutrons:

• Carbon-$12$ has $6$ neutrons
• Carbon-$13$ has $7$ neutrons
• Carbon-$14$ has $8$ neutrons

Properties of isotopes

Applications: studying climate change

Scientists use isotopes to study climate change by analyzing frozen air bubbles deep within ice that has remained frozen for thousands of years. Ice-core studies have been conducted using samples from Greenland, Antarctica, and mountain peaks like Mount Kilimanjaro in Tanzania.

Ions

Nuclide notation can also represent ions. Ions are atoms that have lost or gained one or more electrons, giving them an overall electrical charge. An atom that loses electrons becomes positively charged (because the positive charges in the nucleus now outnumber the negatively charged electrons). Similarly, an atom that gains electrons becomes negatively charged.

Atoms that lose electrons become positive ions (also called cations). Atoms that gain electrons become negative ions (also called anions).

Calculating subatomic particles in ions