0620 · 2.3Isotopes

Defining isotopes

0:00Two atoms can belong to the same element without having the same mass. Carbon twelve and carbon thirteen are both carbon, even though their nuclei contain different numbers of particles. These different versions of an element are called isotopes. By the end, you will understand their symbols, their chemical properties, and their contribution to relative atomic mass. We need three particles to explain this difference. Protons have a positive charge, neutrons have no charge, and electrons have a negative charge.

0:27The protons and neutrons occupy the nucleus at the centre of an atom. Electrons are arranged in shells around it, and a neutral atom has equal numbers of protons and electrons. The number of protons determines which element an atom belongs to. Every carbon atom has six protons, while every chlorine atom has seventeen. Changing the number of protons would mean changing the element itself. Adding or removing neutrons, however, leaves the element unchanged because its proton number stays the same.

0:56Consider carbon twelve and carbon thirteen more closely. Carbon twelve has six protons and six neutrons, whereas carbon thirteen has six protons and seven neutrons. The extra neutron increases the mass but leaves the identity unchanged. Isotopes are different atoms of the same element that have the same number of protons but different numbers of neutrons. That definition tells us exactly what to compare when identifying isotopes. Matching proton numbers show that the atoms belong to the same element.

1:25If their neutron numbers differ, they are different isotopes. Simply noticing different masses is not enough, because atoms of different elements can also have different masses.

Reading isotope and ion symbols

1:35An isotope symbol gives us the information needed to count its particles. The number at the bottom left is the proton number, also called the atomic number. The number at the top left is the mass number. Mass number is the total number of protons and neutrons in the nucleus. For carbon twelve, the bottom number tells us there are six protons. The top number, twelve, is the total of protons and neutrons.

1:59Subtracting six from twelve leaves six neutrons. Because this symbol represents a neutral atom, there are also six electrons to balance the positive charges of its protons. The name carbon twelve refers to the mass number, not the neutron number. A proton and a neutron have almost the same mass. Electrons contribute very little compared with these nuclear particles. This explains why changing the neutron count changes the mass even though the element remains carbon.

2:26We can use the same method for ions, but we must also read the charge at the top right. The symbol shown represents a chloride ion with a 1− charge, made from chlorine thirty-five. Its proton number is seventeen, so it contains seventeen protons. Thirty-five minus seventeen gives eighteen neutrons in its nucleus. A neutral chlorine atom would have seventeen electrons, balancing its seventeen protons. A 1− charge means there is one extra electron compared with that neutral atom.

2:54This ion therefore has eighteen electrons. Its negative charge comes from the excess electron, not from its neutrons, because neutrons have no charge. Always work out protons and neutrons before using the charge to count electrons. The charge does not alter the proton number or mass number shown. Losing or gaining electrons makes an ion, while changing the number of neutrons gives a different isotope. These are separate differences, so an isotope can be represented as either an atom or an ion.

Chemical properties and relative atomic mass

3:23This next part, and the calculation after it, are for the Extended paper. Isotopes of the same element have the same chemical properties. Chemical behaviour depends on electrons, particularly those in the outer shell. Adding a neutron changes the nucleus, but it does not change the number of electrons in a neutral atom. Both neutral carbon twelve and neutral carbon thirteen have six electrons. Their electronic configuration is two electrons in the first shell and four in the second.

3:50Electronic configuration means the arrangement of electrons in shells. Because this arrangement is identical, these isotopes have the same chemical properties despite their different masses. In an explanation, connect the idea all the way through. The isotopes have the same number of electrons and therefore the same electronic configuration. That gives them the same chemical properties. Saying only that they have the same number of protons identifies the element, but does not fully explain the chemical behaviour.

4:18An element can contain a mixture of isotopes, so its relative atomic mass represents an average. More precisely, it is a weighted average that accounts for how common each isotope is. Relative abundance describes the proportion of atoms belonging to each isotope in the sample. Imagine a simplified chlorine sample containing seventy-five percent chlorine thirty-five and twenty-five percent chlorine thirty-seven. The lighter isotope is three times as common as the heavier one.

4:45The average must therefore lie closer to thirty-five than to thirty-seven. Simply finding the midpoint would incorrectly assume that the two isotopes were equally abundant. Relative atomic masses compare atomic masses with one twelfth of the mass of a carbon twelve atom. They have no units because they are relative values. For this calculation, use the supplied isotope masses of thirty-five and thirty-seven. We will combine them according to their abundances rather than giving each isotope equal influence.

5:13Imagine one hundred atoms with exactly the stated proportions. Seventy-five have relative mass thirty-five, so their combined contribution is seventy-five multiplied by thirty-five, giving two thousand six hundred and twenty-five. The other twenty-five atoms have relative mass thirty-seven. Twenty-five multiplied by thirty-seven gives nine hundred and twenty-five. Adding the contributions gives a total relative mass of three thousand five hundred and fifty. Divide by one hundred because we are averaging across one hundred atoms.

5:43The relative atomic mass is therefore thirty-five point five. This answer is between the two isotope masses and closer to the more abundant isotope, as expected. The general method is to multiply each isotope’s relative mass by its abundance, then add the contributions and divide by the total abundance. Percentage abundances total one hundred. If a question gives relative numbers of atoms instead, divide by their sum. Always use the isotope masses supplied rather than replacing more precise values with mass numbers.

6:13The decimal answer does not mean an individual chlorine atom contains part of a neutron. Mass number describes the whole number of protons and neutrons in one atom. Relative atomic mass describes the weighted average for an element. A sample can therefore have a decimal average even though each atom contains whole particles.

Exam mistakes and recap

6:30Common exam mistakes include reversing the numbers in an isotope symbol and forgetting that mass number includes both protons and neutrons. Another mistake is treating a negative ion as having fewer electrons. Extra negatively charged electrons produce a negative charge. For abundance calculations, avoid the simple midpoint unless the isotopes are equally abundant. An exam might show carbon twelve and carbon thirteen and ask you to state the particle numbers for carbon thirteen.

6:59It might then ask you to explain their identical chemical properties. Read the lower number first, subtract it from the upper number, and check for a charge. Give the particle counts and the explanation as separate clear points. Carbon thirteen has six protons. Thirteen minus six gives seven neutrons. As a neutral atom, it also has six electrons. The isotopes have the same number of electrons and therefore the same electronic configuration.

7:25This means they have the same chemical properties. The main connection is simple: proton number determines the element, while neutron number distinguishes its isotopes. Mass number counts protons and neutrons, and an ion’s charge tells you about its electron count. Identical electron arrangements explain identical chemical properties. Relative atomic mass combines isotope masses according to their abundances, giving a weighted average rather than the mass of one atom.