Atomic Masses and Molecular Masses

Introduction to atomic masses and molecular masses

Given that the atom of an element is small, determining the actual mass of a single atom is difficult. The  IUPAC came up with a scale basing on C-12 Isotope where the Carbon was used as a base for calculating the atomic mass of all other elements. C-12 was recommended since it is a stable and very familiar element. Carbon-12 was assigned12.00 a.m.u by the IUPAC. In most cases, calculating the individual atomic masses is quite impossible hence, the scientists take the relative atomic masses of a group of atoms by weighing them as a whole. So the mass of the entire group divided by the total number of atoms participated will give the relative atomic mass.

 

Table Of Contents

1. Introduction to atomic masses and molecular masses

2. Definition of AMU

3. Relative Atomic mass      

4. AMUs of the most common elements

5. Molecular mass

6. Law of Conservation of Mass

7. Gay Lussac’s law

 

Definition of AMU

“The atomic mass unit is defined as the ratio of the average mass of one atom of the element to one-twelfth of the mass of one atom of C-12”. The atomic weight of an atom is a dimensionless number when it is divided by unified atomic weight or Daltons. This is called the relative isotopic mass. The atomic masses of elements vary from 1.008 amu for hydrogen up to 250 amu for elements that have a very high atomic number. Mass of molecules can be determined by adding the average atomic mass of each atom in the molecule.

 

Relative Atomic mass      

The RAM of an element is the sum of the product of the percentage abundance and the corresponding atomic masses of the isotopes. For example: calculate the RAM of the chlorine atom whose relative abundances and corresponding masses are given below.

Isotope

Relative abundances

37Cl

25

35Cl

75

 

For example;

 

AMUs of the most common elements

Element Name

Atomic Mass units( AMU)

Element Name

Atomic Mass units( AMU)

Aluminium

26.981

Molybdenum

95.94

Argon

39.948

Neon

20.180

Arsenic

74.922

Nickel

58.693

Barium

137.327

Nitrogen

14.007

Beryllium

9.012

Oxygen

15.999

Bismuth

208.980

Palladium

106.42

Boron

10.811

Phosphorus

30.974

Bromine

79.904

Platinum

195.084

Calcium

40.078

Potassium

39.098

Carbon

12.011

Radium

n/a

Chlorine

35.453

Radon

n/a

Cobalt

58.933

Rubidium

85.468

Copper

63.546

Scandium

44.956

Fluorine

18.998

Selenium

78.96

Gallium

69.723

Silicon

28.086

Germanium

72.64

Silver

107.868

Gold

196.967

Sodium

22.990

Helium

4.003

Strontium

87.62

Hydrogen

1.008

Sulfur

32.065

Iodine

126.904

Tantalum

180.948

Iridium

192.217

Tin

118.710

Iron

55.845

Titanium

47.867

Krypton

83.798

Tungsten

183.84

Lead

207.2

Uranium

238.029

Lithium

6.941

Xenon

131.293

Magnesium

24.305

Zinc

65.409

Manganese

54.938

Zirconium

91.22

 

Molecular mass

Molecular mass is the total atomic masses of the elements that are contained in a molecule. It is calculated by finding the product of the atomic masses by the corresponding number of atoms and adding them together. Example molecular mass of H2SO4 can be obtained as shown below

 (1x2) + 32 + (16x4)

 2 +32 + 64

= 98

 

Law of Conservation of Mass

The law states that the mass in an isolated system is neither created nor destroyed by chemical reactions or physical transformations. According to this law, the mass of the reactants is always equal to the mass of the products in any given chemical reaction. Consider the reaction between Sulphur and Oxygen.

S(s) + O2 (g)  -----------------> SO2 (g)

 

Gay Lussac’s law

The law was discovered by Gay Lussac’s in the year 1805. It summarizes the relationship between the reacting volumes of gases. The law states that when gases react, they do son in volumes that bear a simple ratio to one another and to the volumes of products if gaseous, provided temperature and pressure are kept constant.

For example,

2CO (g) + O2 (g)     ------------->  CO2 (g)

20cm3        10cm3             10 cm

2vol             1vol             1 vol

Thus the ratio becomes 2:1:1

In an experiment, 20cmof Sulfur (IV) Oxide are found to react completely with 10cmof Oxygen gas to produce 30cmof Sulfur (IV) Oxide. Determine the equation for the reaction.

2SO2 (g)   +   O2 (g)  -------------->    SO3 (g)

20cm          10cm3                 20cm3

2vol               1 vol                2vol

And the ratio is 2:1:2

Results from experiments have shown that equal volumes of all gases under the same conditions of temperature and pressure contain an equal number of molecules. This relationship leads to Avogadro’s Law which states that "Equal volumes of gases will contain an equal number of molecules." The Avogadro's constant is denoted by "L" is always 6.02× 1023

 





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