# Relative Atomic Mass Of Oxygen

The formula weight is simply the weight in atomic mass units of all the atoms in a given formula. Formula weights are especially useful in determining the relative weights of reagents and products in a chemical reaction. These relative weights computed from the chemical.

Related Topics: More Chemistry Lessons
In this lesson, we will learn
• how to calculate the relative formula mass or relative molecular mass
• how to use the relative molecular mass to calculate the percent mass of an element in a compound
• how to use the relative molecular mass to calculate the percent mass of water in a compound

Atomic mass is the measure of a mass of one atom in relation to 1/12 mass of Carbon 12 atom and measured in atomic mass unit(amu) or Dalton unit or grams/mol.Relative atomic mass is also the. The formula weight is simply the weight in atomic mass units of all the atoms in a given formula. Formula weights are especially useful in determining the relative weights of reagents and products in a chemical reaction. These relative weights computed from the chemical equation are sometimes called equation weights. Oxygen-17 atom is the stable isotope of oxygen with relative atomic mass 16.999131. The least abundant (0.038 atom percent) isotope of naturally occurring oxygen.

The following diagram shows how to calculate the relative molecular mass or relative formula mass. Scroll down the page for more examples and solutions.
What is relative formula mass and relative molecular mass?

The relative formula mass of a substance is the sum of the relative atomic masses of the elements present in a formula unit. The symbol for relative formula mass is Mr.
If the substance is made of simple molecules, this mass may also be called the relative molecular mass.

How to calculate the relative formula mass?

Example:

What is the relative mass formula of hydrogen gas? (Relative atomic mass: H = 1)

Solution:

The formula for hydrogen gas is H2. Each molecule contains 2 hydrogen atoms.
The relative mass formula of hydrogen gas is
Mr(H2) = 2 × Ar(H) = 2 × 1 = 2

Example:

What is the relative mass formula of water? (Relative atomic masses: H = 1, O = 16)

Solution:

The formula for water is H2O. Each molecule contains 2 hydrogen atoms and 1 oxygen atom.
The relative mass formula of water is
Mr(H2O) = 2 × Ar(H) + Ar(O) = 2 × 1 + 16 = 18

Example:

What is the relative mass formula of sodium chloride? (Relative atomic masses: Na = 23, Cl = 35.5)

Solution:

Sodium Chloride is an ionic solid with the formula Na+Cl-.
The relative mass formula of sodium chloride is
Mr(NaCl) = Ar(Na) + Ar(Cl) = 23 + 35.5 = 58.5
How to find the Percent Mass of Elements in a compound?
How to use Mr to calculate the percent mass of an element in a compound?

Example:

What percentage of the mass of ammonium nitrate is nitrogen? (The formula for ammonium nitrate is NH4NO3, Relative atomic masses: H = 1, O = 16, N = 14)

Solution:

Mr(NH4NO3) = (2 × 14) + (4 × 1) + (3 × 16) = 28 + 4 + 48 = 80 Mass of nitrogen in the formula = 28
Mass of nitrogen as a fraction of the total = Mass of nitrogen as percentage of total mass

Example:

What percentage of the mass of O in NaNO3? (Relative atomic masses: Na = 23, O = 16, N = 14)

Solution:

Mr(NaNO3) = 23 + 14 + (3 × 16) = 23 + 14 + 48 = 85

Mass of O in the formula = 48
Mass of O as a fraction of the total =
Mass of O as percentage of total mass
How to find percent by mass and percent composition?
Example:
Find the percent composition by mass of potassium dichromate (K2Cr2 O7)
Step 1: Find the molar mass of the compound.
Step 2: Divide the total mass of each element by the molar mass and multiply by 100 to find the % mass.
• Show Step-by-step Solutions
How find the percent of an element in a compound?
Finding the percent by mass means finding the mass of the elements in the compound and adding the masses for the total mass.
Example:
You have a 7364 milligrams sample of SO2. How many grams of sulfur are in the sample?
How to calculate the Mass Percent of an Element in a Compound?
Mass Percent Composition of an Element in a Compound
To calculate the mass percent composition (or simply, the mass percent) of an element in a compound, we divide the mass of the element in 1 mol of the compound with the mass of 1 mol of the compound and multiply by 100%.
Examples:
1. What is the mass percent of carbon in carbon dioxide?
2. What is the mass percent of oxygen in carbon dioxide? • Show Step-by-step Solutions
How to find the Percent Mass of Water in a compound?

## Relative Atomic Mass Example

How to use Mr to calculate the percent mass of water in a compound?

Example:

What percentage of the mass of magnesium sulfate is water? (Given that the formula for magnesium sulfate is MgSO4•7H2O, Relative atomic masses: H = 1, O = 16, S = 32, Mg = 24)

Solution:

Mr(MgSO4•7H2O) = 24 + 32 + (4 × 16) + (7 × 18) = 246
Mr(H2O) = 2 × Ar(H) + Ar(O) = 2 × 1 + 16 = 18
Mass of water in the formula = 18 × 7 = 126
Mass of water as a fraction of the total =
Mass of hydrogen as percentage of total mass
How to calculate the Water of Crystallization?
This video outlines how to determine the percent, by mass, of water trapped in a hydrate's crystal lattice.
Example:
What is the percent of water, by mass, in the hydrate CuSO4•2H2O How to calculate the percentage of water in the formula of a hydrated compound?
Example:
Determine the % water in the following hydrates:
a) CuSO4•5H2O
b) CaCl2•2H2O
c) KAl(SO4)2•12H2O

## Relative Atomic Mass Of Oxygen 16

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Taken from http://www.sizes.com/units/atomic_mass_unit.htm

## History of the atomic mass unit

Stanislao Cannizzaro (1826–1910), the pioneer in this field, adopted the hydrogen atom as a standard of mass and set its atomic weight at 2. Others accepted the idea of using a specific atom as a standard of mass, but preferred a more massive standard in order to reduce experimental error.

As early as 1850, chemists used a unit of atomic weight based on saying the atomic weight of oxygen was 16. Oxygen was chosen because it forms chemical compounds with many other elements, simplifying determination of their atomic weights. Sixteen was chosen because it was the lowest whole number that could be assigned to oxygen and still have an atomic weight for hydrogen that was not less than 1.

The 0=16 scale was formalized when a committee appointed by the Deutsche Chemische Gesellschaft called for the formation of an international commission on atomic weights in March 1899. A commission of 57 members was formed. Since the commission carried on its business by correspondence, the size proved unwieldy, and the Gesellschaft suggested a smaller committee be elected. A 3-member International Committee of Atomic Weights was duly elected, and in 1903 issued its first report, using the 0=16 scale.5

### Taking isotopes into account

The discovery of isotopes complicated the picture. In nature, pure oxygen is composed of a mixture of isotopes: some oxygen atoms are more massive than others.

This was no problem for the chemists’ calculations as long as the relative abundance of the isotopes in their reagents remained constant, though it confirmed that oxygen’s atomic weight was the only one that in principle would be a whole number (hydrogen’s, for example, was 1.000 8).

Physicists, however, dealing with atoms and not reagents, required a unit that distinguished between isotopes. At least as early as 19276 physicists were using an atomic mass unit defined as equal to one-sixteenth of the mass of the oxygen-16 atom (the isotope of oxygen containing a total of 16 protons and neutrons).

In 1919, isotopes of oxygen with mass 17 and 18 were discovered.7 Thus the two amu’s clearly diverged: one based on one-sixteenth of the average mass of the oxygen atoms in the chemist’s laboratory, and the other based on one-sixteenth of the mass of an atom of a particular isotope of oxygen.

## Relative Atomic Mass Ar

In 1956, Alfred Nier (at the bar in the Hotel Krasnapolski in Amsterdam) and independently A. Ölander8, both members of the Commission on Atomic Masses of the IUPAP, suggested to Josef Mattauch that the atomic weight scale be based on carbon-12. That would be okay with physicists, since carbon-12 was already used as a standard in mass spectroscopy. The chemists resisted making the amu one-sixteenth the mass of an oxygen-16 atom; it would change their atomic weights by about 275 parts per million. Making the amu one-twelfth the mass of a carbon-12 nucleus, however, would lead to only a 42 parts per million change, which seemed within reason.

Mattauch set to work enthusiastically proselytizing the physicists, while E. Wichers lobbied the chemists.9 In the years 1959–1961 the chemists and physicists resolved to use the isotope carbon-12 as the standard, setting its atomic mass at 12.