Atoms are small - very small. Imagine a grapefruit, with all of its atoms as nitrogen atoms. If you were to blow up the grapefruit so it was the size of the Earth, how big do you think each of the atoms would be? The size of the UK? A football field? A London taxi? They would only be the size of a blueberry. You can imagine that this makes finding out some of an atom’s simplest properties, such as its mass, a little bit difficult. To represent mass, we use something called mass number, a value that is particularly important when looking at isotopes.
An atom’s mass number, also known as A, is the combined total number of protons and neutrons in its nucleus.
You’ll remember from "Fundamental particles" that atoms have three main subatomic particles. In actual fact they have a few more than that, but we only really care about the proton, neutron and electron right now. You might also remember that while protons and neutrons have very similar masses, the mass of an electron is almost non-existent. So when calculating an atom’s mass, we can ignore its number of electrons. The table below gives you a reminder of the relative masses and charges of the three fundamental particles:
Fig. 1 - A table showing the relative masses and charges of protons, neutrons and electrons
But knowing an atom’s mass number isn’t very helpful, as it could contain multiple different combinations of protons and neutrons. For example, take an atom with a mass number of 14. Does it have eight protons and six neutrons, or six protons and eight neutrons, or perhaps seven of each? This is where the atomic number comes in handy.
An atom’s atomic number, Z, is just the number of protons in its nucleus.
Atomic number also tells us exactly which element in the periodic table it is a part of, as all atoms in an element have the same number of protons.
Why do we use the letter A to represent mass number and the letter Z to represent atomic number? Well, A comes from the German word Atomgewicht, meaning atom weight, whilst Z comes from the German word Zahl, which means number.
The number of fundamental particles in an atom
To find out exactly how many protons, neutrons and electrons an atom contains, you need to know its mass number and atomic number. But lucky you - the periodic table gives you this information for free! Let’s take carbon as an example.
Fig. 2 - Carbon
The larger number, 12, is its mass number and the smaller number, 6, is its atomic number. What can we learn from this? Well, because its atomic number is 6, carbon must have six protons. Neutral atoms contain the same number of electrons as protons, so that means carbon also has six electrons. You’ll remember that the mass number is simply the number of protons and neutrons combined, so to find just the number of neutrons we can subtract the number of protons, the atomic number, from the mass number. Here, 12 - 6 = 6. This carbon atom has six neutrons.
We could also show mass number and atomic number by writing \(_6^{12}C\).
Let’s look at another example.
Fig. 3 - Hydrogen
Here, hydrogen has a mass number of 1 but also an atomic number of 1, to the nearest whole number. It must therefore have one proton and one electron. 1 - 1 = 0, and so it has no neutrons.
What is an isotope?
You’ll notice that the mass number for hydrogen is not a whole number. In fact, hardly any of the mass numbers on the periodic table are. Why is this the case? Well, mass numbers account for isotopes of an element. We’ll explore what that means down below, but for now, let's take a look at what exactly an isotope is.
Isotopes are atoms of the same element with the same number of protons and electrons but different numbers of neutrons.
This means that isotopes have different atomic weights and mass numbers but the same atomic number. Isotopes occur naturally due to the differing stabilities of different arrangements of protons and neutrons in the nucleus.
Isotopes of the same element have the same chemical properties as they have the same electron configuration. However, they have different physical properties because they have different masses. We represent isotopes by a small superscript mass number to the left of the chemical symbol. For example, 
has an atomic mass of 12 but 
has an atomic mass of 13, meaning it contains an extra neutron.
Isotopes of hydrogen
Hydrogen has three naturally occurring isotopes:
, simply known as hydrogen, is the most abundant isotope of hydrogen. Each atom has one proton, one electron and no neutrons.
, also known as deuterium, has the same number of protons and electrons as hydrogen but has one neutron.
, also known as tritium, also has one proton and one electron but has two neutrons.
Radioactivity
The number of neutrons in an atom is determined by the relative stability of the nucleus. Lighter elements such as carbon are at their most stable when they have equal numbers of protons and neutrons in their nuclei, whereas heavier elements prefer having slightly more neutrons. This keeps their nuclei stable. Atoms with too many or too few neutrons do exist in nature but they are often unstable. This means that they will decay and release radiation. We call these atoms radioactive isotopes. To become more stable they can either emit an alpha particle, which is a package of two protons and two neutrons, or a beta particle, which is a fast moving electron. This occurs if a neutron turns into a proton and an electron. Both types of radioactive decay change the mass number to atomic number ratio of the atom, making it more stable.
What is an ion?
We know that all atoms in an element have the same number of protons, and we’ve just learnt that they can have different numbers of neutrons. But what happens if they have different numbers of electrons?
An ion is an atom that has gained or lost one or more electrons to form a charged particle.
However, ions still have the same number of protons, so the same atomic number. A negative ion has gained electrons whereas a positive ion has lost electrons, as electrons are negatively charged.
Ions of the same element have different chemical properties because they have different electron configurations. Atoms tend to want to have a full outer shell of electrons, and so their number of electrons can dramatically change their reactivity. For example, you won’t find sodium atoms anywhere in nature but you will find positive sodium ions, \(Na^+\). This is because sodium atoms readily react to lose two electrons so that they have a full outer shell.
Ions are represented by a small superscript number to the right of the chemical symbol showing the charge on the ion, or by roman numerals. For example, Fe³⁺ has lost 3 electrons from the neutral Fe atom to form a positive ion with a charge of +3. This can also be represented by iron (III).
Let’s take the lithium ion, \(Li^+\), as an example. Lithium has an atomic number of 3 and so has three protons. If this was an uncharged atom, we would also expect it to have three electrons. However, this ion is positively charged, meaning it has lost an electron. Therefore, \(Li^+\) has just two electrons.
What is relative atomic mass?
So why are the mass numbers of elements not whole numbers? Well, this is because mass numbers in the periodic table use relative atomic mass.
Relative atomic mass is the average mass of an atom of an element in a sample compared to 1/12th of the mass of a ¹²C atom, taking into account the abundance of different isotopes.
We measure relative atomic mass on the carbon-12 scale, where ¹²C has a mass of exactly 12.
To work out the relative atomic mass of a sample, represent the percentage abundance of each isotope as a decimal. Multiply it by the isotope’s mass and add all the values up. This is your relative atomic mass.
For example, a sample of chlorine may contain 75% 
and 25% 
.
0.75 x 35 = 26.25
0.25 x 37 = 9.25
26.25 + 9.25 = 35.5
Hence the relative atomic mass of chlorine is 35.5.
For a more detailed look at mass numbers, check out the article "Relative Mass".
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