Atoms, Isotopes & Relative Atomic Mass (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching Chemistry, and this video is about atoms, isotopes and relative atomic mass. An atom's identity and its mass both come down to counting the same three particles: the protons determine which element it is, and the neutrons, which can vary, along with the protons, determine its mass.

Every atom is built from protons, neutrons and electrons. The number of protons decides which element the atom is, and every atom and ion of that element has the same number of them. The number of neutrons can differ from one atom of an element to another, and because it can differ, those atoms have different masses. So an element has one fixed identity and a mass that has to be averaged.

This video will cover the inside of an atom, which is where these particles sit and how their masses and charges compare. How to count them, using atomic number and mass number to work out the protons, neutrons and electrons in an atom or an ion. Then we will look at atoms of the same element that differ in their number of neutrons. These are called isotopes, and the relative atomic mass that comes from them.

All matter is made of atoms, the smallest part of an element that can take part in a chemical reaction. An atom is mostly empty space. At its centre is a very small, dense nucleus containing protons and neutrons, which together are called nucleons. The nucleus is positively charged overall, because protons are positive and neutrons are neutral. Negatively charged electrons sit in orbitals in the space around it, and the electrostatic attraction between that positive nucleus and these negative electrons is what holds the atom together.

These particles are far too small for their masses and charges to be measured in conventional units like grams and coulombs, so chemists compare them with each other instead, using relative mass and relative charge. A proton has a relative charge of plus one and a relative mass of one. A neutron is neutral, and also has a relative mass of one. An electron has a relative charge of minus one and a relative mass of one over one thousand eight hundred and thirty-six, so its mass is almost negligible. Almost all of the atom's mass is therefore concentrated in the nucleus, while the electrons create a cloud of negative charge around it.

Our model of the atom developed over time, from Thomson's plum pudding model to the nuclear model we use now. Because only the protons and neutrons carry any real mass, the mass of an atom is a count of the particles in its nucleus.

The atomic number, symbol Z, is the number of protons in the nucleus, and it is the atomic number that determines which element you have. Every atom and ion of the same element has the same number of protons. The mass number, symbol A, is the total number of protons and neutrons. Subtract the atomic number from the mass number and you get the number of neutrons. A neutral atom has the same number of electrons as protons, but an ion has either gained or lost electrons: a positive ion has lost electrons, so it has fewer electrons than protons, and a negative ion has gained electrons, so it has more. Both the mass number and the atomic number are given for every element in the Periodic Table.

The atomic number is the fixed half of this: it never changes for an element, while the mass number can differ from one atom of that element to the next.

Isotopes are atoms of the same element that contain the same number of protons and electrons, but a different number of neutrons. We name them with the chemical symbol followed by the mass number: for example carbon-12 and carbon-14 are both carbon, with six neutrons and eight neutrons respectively. Because isotopes have the same number of electrons in their outer shells, and it is electrons that take part in chemical reactions, isotopes of an element show the same chemical properties. Neutrons are neutral, so they only add mass, which means isotopes differ in physical properties such as mass and density.

Because an element exists as a mixture of isotopes with different mass numbers, we don't quote a single mass for it. Instead we use relative atomic mass, A r, which uses the average mass of the element's isotopes. It is a weighted average: each isotope counts according to how much of it is present, which is its relative abundance. Those abundance values are either given to you, or read straight off a mass spectrum, like this one for boron. Due to the different masses of the isotopes their time of flight will differ allowing an electrical current to be measured for each one and therefore their abundance determined.

The protons keep the element the same, and averaging the neutrons gives that element its mass.

When you define mass number, state it as the total number of protons and neutrons in the nucleus. Never use the word 'mass' or 'average' in that definition, because that describes relative atomic mass, and examiners will not award the mark if you confuse the two. And when you are filling in a table of subatomic particles for an ion, remember to adjust the electron count: subtract electrons for a positive ion, add them for a negative ion. That is the most common error examiners see on this topic.

The number of protons defines which element an atom is. The mass number is protons plus neutrons, and taking the atomic number away from it gives you the neutrons. Isotopes of an element differ only in their neutrons, which is why they behave identically in chemical reactions but differ in mass. And because a sample contains a mixture of isotopes, the mass we use for an element is the relative atomic mass, the average of its isotopes weighted by how abundant each one is. The protons give the element its identity, and averaging the neutrons gives it its mass.