Atoms are incredibly small. If we took all the atoms in a grapefruit and made them nitrogen atoms, and then blew up the grapefruit to the size of the Earth, each atom would be as small as a blueberry. This makes it hard to figure out some of the basic things about atoms, like their mass. To represent mass, scientists use something called mass number, which is really important when studying isotopes.
An atom’s mass number, or A, is the total number of protons and neutrons in its nucleus. Remember that atoms have three main subatomic particles: protons, neutrons, and electrons. When calculating an atom’s mass, we can ignore the number of electrons because they are so tiny. Here’s a table to remind you of the relative masses and charges of the three fundamental particles:
Understanding isotopes can get pretty complicated, but it all starts with knowing about mass number.
But knowing an atom’s mass number isn’t very helpful by itself because it could have different combinations of protons and neutrons. For example, an atom with a mass number of 14 could have eight protons and six neutrons, or six protons and eight neutrons, or maybe seven of each. This is where the atomic number comes in handy.
An atom’s atomic number, or Z, is just the number of protons in its nucleus. Atomic number also tells us which element in the periodic table the atom is a part of because all atoms in an element have the same number of protons.
You might be wondering why we use A for mass number and Z for atomic number. Well, A comes from the German word Atomgewicht, which means atom weight, while Z comes from the German word Zahl, which means number.
To find out exactly how many protons, neutrons, and electrons an atom has, you need to know its mass number and atomic number. Luckily, the periodic table provides this information for you! Take carbon as an example.
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. Since neutral atoms contain the same number of electrons as protons, carbon also has six electrons. To find the number of neutrons, we can subtract the atomic number from the mass number. Here, 12 - 6 = 6. This carbon atom has six neutrons.
We could also write the mass number and atomic number like this: .
Let’s look at another example.
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.
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.
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.
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.
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, . 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, , 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, has just two electrons.
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".
Here are some key takeaways about isotopes:
What is an isotope?
Isotopes are atoms of the same element with different atomic masses. This is because they have different numbers of neutrons.
Are all isotopes radioactive?
No. Most elements have at least one stable isotope.
How do you calculate relative atomic mass of an isotope?
To calculate relative atomic mass, first convert % abundance of each isotope into a decimal. Then multiply each decimal by the mass of its respective isotope, and add all these values up.
How many isotopes of hydrogen are there?
There are three isotopes of hydrogen: hydrogen, deuterium and tritium.
What is the half life of a radioactive isotope?
The half life of a radioactive isotope is the time taken for half of the radioactive atoms present in a sample to decay.
