Atomic Structure and Isotopes

Atomic structure

The nucleus is made up of protons and neutrons, which are the heaviest particles. Electrons are approximately 2000 times lighter than a proton or neutron and weigh practically nothing, therefore most of the mass comes from the tiny nucleus. Protons are positively charged whereas neutrons have no charge, making the nucleus positively charged overall. Electrons are negatively charged, which enables the nucleus to keep hold of its electrons, due to the attraction between the opposite charges. In an atom, the number of electrons is equal to the number of protons in the nucleus. Atoms have no overall electrical charge.

Each atom has an atomic number and a mass number, which are displayed on the Periodic table as the small number and larger number next to each chemical symbol. The atomic number tells us the number of protons which is unique for each element. If the number of protons changes (and therefore the atomic number changes), which could occur during radioactive decay, for example, then the identity of the element also changes. The mass number tells us the number of both protons and neutrons in the nucleus. Remember that most of an atom’s mass comes from the nucleus and that the nucleus is made up of both protons and neutrons. Two atoms of the same element can have different mass numbers because they are isotopes.

Isotopes

Isotopes are atoms that have the same number of protons but different numbers of neutrons. Therefore they have the same atomic number but a different mass number. For example, two isotopes of chlorine are chlorine-35 and chlorine-37, in which chlorine-37 has two additional neutrons in its nucleus to give it a higher mass number.

The actual mass of an atom is so small that it’s more useful to compare its mass to another atom and express this as a relative mass. The mass of an atom is always compared to the mass of carbon-12. Make sure you know the following definitions:

The mass numbers that we see in the periodic table are an average of all the masses of all isotopes of a particular element. We call this the relative atomic mass (RAM) and can be calculated using the following formula:

Mass spectrometry

We can use mass spectrometry to provide information about which isotopes are present in a sample. Mass spectrometry takes place in the following steps:

1.       Ionisation – the sample is ionised when high energy electrons are fired at it from an ‘electron gun’. The electron gun is made up of a hot wire filament with a current running through it which emits electrons, forming 1+ ions. The 1+ ions are attracted to a negatively charged electric plate where they are accelerated.

2.       Acceleration – the positive ions are accelerated using an electric field so they all have the same kinetic energy.

3.       Detection – positive ions hit a negatively charged plate and gain electrons. This generates a movement of electrons, therefore an electric current that is measured. The size of the current gives a measure of the number of ions hitting the plate.

Here’s an example of a mass spectrum produced by a sample of bromine, which exists as bromine-79 and bromine-81 isotopes.

The y-axis gives the abundance of ions where the height of each peak is proportional to the relative isotopic abundance (so the taller the peak, the more of the isotope is present in the sample). The x-axis gives the mass/charge ratio (m/z) of the isotope. Given that the charge of the ions is always +1, the m/z is effectively the mass of each ion.

In the mass spectrum above, we can see that the sample consists of:

Relative formula mass

The relative formula mass (Mr) of a compound is the sum of the relative atomic masses of the atoms in the numbers shown in the formula.