In this lesson, we will learn about Isotopes, Isotope Notation, Atomic Mass Unit (amu), and how to calculate the Atomic Mass of an element.

Related Topics: More Lessons on Chemistry

**Isotopes**are atoms of the same element that have different number of neutrons. (In order for them to be atoms of the same element, their number of protons would be the same)

*Example: *

is an isotope of chlorine that has 17 protons and 18 neutrons

is an isotope of chlorine that has 17 protons and 20 neutrons

- Generally, isotopes behave the same way during chemical reactions.
- The extra neutrons just change the mass of the atom and its density.

What are Isotopes?

In this video, we'll learn about what isotopes are and how to write atomic number and mass number in isotope notation. We talk about a simple analogy with cars to explain this tutorial. Isotopes are versions of an atom or an element that have the same number of protons, but different numbers of neutrons. Isotopes and isotope notation are particularly important in nuclear chemistry.

Isotope Notation

Learn how to write atoms in isotope notation! In isotope notation, you can quickly show how many protons, neutrons, and electrons are in an atom. You put the atomic number, mass number, and net charge around the chemical element symbol. Isotope notation is particularly important in nuclear chemistry, because if you're doing fission, fusion, alpha decay, beta decay, positron emission, or electron capture, you want to be able to tell how many neutrons and protons are in the nucleus.

- Since the mass of an atom would be extremely small when measured in grams, it would be more convenient to measure the masses of atoms relative to a standard atom.
- The
**standard atom**chosen is (carbon-12) isotope. An atom of carbon-12 is taken to have a mass of 12**atomic mass unit**(amu). - Since one carbon-12 atom has 6 proton and 6 neutron,

mass of one proton (neutron) = mass of one carbon-12 atom

= 1 amu (atomic mass unit)

Atomic Mass Units

- The
**relative atomic mass**() of an element is the average mass of the naturally occurring atoms of the element. This quantity takes into account the percentage abundance of all the isotopes of an element which exist.*A*_{r} - The formula for relative atomic mass is

*Example:*

Given that the percentage abundance of is 75% and that of is 25%, calculate the *A _{r}* of chlorine.

*Solution: *

*Example:*

Bromine has two isotopes, Br-79 and Br-81. Both exist in equal amounts. Calculate the relative atomic mass of bromine.

*Solution: *

*Example:*

The neon element has three isotopes. They are 90.92% of , 0.26% of and 8.82% of

*Solution: *

Atomic Mass: Introduction

What is atomic mass? It is a weighed average of the different isotopes of an element. It is sometimes referred to as atomic weight, relative atomic mass, or average atomic mass. We look at how to calculate and determine the weighed average of elements using atomic mass units.

Atomic Mass: How to Calculate Isotope Abundance

How do you determine and calculate isotope abundance when you know the relative atomic mass (also known as atomic weight), as measured in amu or atomic mass numbers? Here we will go through the algebra and reasoning to figure out the amount of abundances of the isotopes, in percentages and in decimals

There are two stable isotopes of chlorine: Chlorine-35 (which weighs 34.97 amu) and Chlorine-37 (which weighs 36.97 amu). If the relative atomic mass of chlorine is 35.45, what is the abundance of each isotope?

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