Atomic Weight in Chemistry

The atomic weight—or more correctly stated: relative atomic mass—of an element is a number that is approximately equal to the average number of protons and neutrons of the isotopes of the element.

Since the most common isotope makes up a large percentage of the material, the atomic weight is close to the number of of protons and neutrons in that element. For example, Oxygen-16 is the most common isotope in nature, thus the atomic weight of Oxygen is approximately 16 (15.9994).

Several factors cause it not to be exactly that number. The atomic weight is used in chemistry as a relative unit to determine the mass of molecules involved in chemical reactions. It can also show the number of atoms or molecules in a mass of material.

This lesson will answer those questions.

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Atomic weight not an integer

The relative atomic weight or mass of an element is designated in atomic mass units (amu or u). It originally was supposed to be the sum of the protons and neutrons for a given element. Thus, since the nucleus of Oxygen typically consists of 8 protons and 8 neutrons, its atomic weight should be 16 u. Unfortunately, that is not the case.

Slightly different

The list in the Chemical Elements lesson shows that the mass of each element is slightly off from the integer or whole number one would expect. The atomic weight for Oxygen is listed as 15.9994 u.

Reasons

There are several reasons for this discrepancy.

First is that although scientists designated the proton and neutron as each being equal to 1 u, careful measurement showed that each was slightly different from that value. By defining Carbon-12 as having an atomic weight of 12.000 u, the values for the proton and neutron were established.

Another reason is that when the protons and neutrons gather together in the nucleus, some of their mass is lost. This is called the mass defect and is the basis for nuclear energy.

Finally, elements consist of atoms with different number of neutrons. The atomic weight for the element averages the the atomic weights of the various isotopes of the element. (See the lesson on Isotopes for more information.)

Oh yes, we forgot about the electrons. Although they weigh only about 0.001 u, they still make a contribution to the atomic mass.

(See the lesson on Atomic Weight in the Physics section for more information.)

Practical use

For all practical purposes, the atomic weight is rounded off to the nearest whole number or integer. In most uses in Chemistry, it is not necessary to use the exact value.

Weight of materials

One use of atomic weight is to find out the weight of molecules resulting in chemical reactions.

For example, consider what happens when Zinc is exposed to Sulfuric Acid. The result is Zinc Sulfate and Hydrogen gas:

Zn + H₂SO₄ → ZnSO₄ + H₂

The atomic weight of each element involved (rounding off) is:

• Zn = 65
• H = 1
• S = 32
• O = 16

Thus, the atomic weights involved (in parentheses) are:

Zn(65) + H₂SO₄(98) → ZnSO₄(161) + H₂(2)

You can see the relative weights of the parts of this chemical equation. Thus, you might mix 65 grams of Zn to 98 grams of H₂SO₄ to get 161 grams of ZnSO₄ plus some gas. Also note that the atomic weight on the left side of the equation equals that on the right side.

Number of atoms or molecules

The atomic weight is used to determine the number of atoms in a given weight. The same applies to the atomic weight of a molecule, also called its molecular weight.

A mole (mol) is defined as the number of grams of a substance that is equal to its atomic or molecular weight. For example, a mole of water (H₂O) is the molecular weight of water. Since H = 1 u and O = 16 u, then a mole of H₂O = 18 grams.

There is a number called Avogadro’s number that states the number of molecules in 1 mole.

A_N = 6.02 * 10²³ molecules

Thus, 18 grams of water contains 602,000,000,000,000,000,000,000 molecules.

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Science can be amazing

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Resources and references

The following resources provide information on this subject:

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