Dalton was proved incorrect and his theory was modified

• Atoms that have the same number of protons are always atoms of a specific element.
• Example: Carbon

BUT
• atoms can have different numbers of neutrons and still be an atom of a specific element.

• This is because elements can have isotopes
(basically atoms of the same element with a different number of neutrons in their nuclei).

• When Dalton stated his atomic theory in the early 1800’s, he assumed that all of the atoms of a given element were identical.

• Over 100 years after Dalton, James
Chadwick discovered that the nuclei of most atoms contains neutrons as well as protons.

• Dalton’s theory now states:
– All atoms of the same element contain the same number of protons and electrons, but atoms of a given element may have different numbers of neutrons.

• Hydrogen – 1
– Also written H-1
– Also known as protium
– Hydrogen has an atomic number of 1, so it has 1 proton
– The hyphen notation above tells us that the mass number of H-1 is 1
• Number of neutrons = mass number – atomic number
• So H-1 must have 0 neutrons

• Hydrogen – 2
– Also written H-2
– Also known as
Deuterium
– Hydrogen has an atomic number of 1, so it has 1 proton
– The hyphen notation above tells us that the mass number of H-2 is 2
• Number of neutrons = mass number – atomic number
• So H-2 must have 1 neutron

• Hydrogen – 3
– Also written H-3
– Also known as
Tritium
– Hydrogen has an atomic number of 1, so it has 1 proton
– The hyphen notation above tells us that the mass number of H-3 is 3
• Number of neutrons = mass number – atomic number
• So H-3 must have 2 neutron

• For chlorine found on the periodic table (the most common form of chlorine that is found in nature) Chlorine-35
• For Chlorine-37
• Chlorine-35 = 18 neutrons, Chlorine-37 = 20 neutrons
• For Cobalt found on the periodic table Cobalt-59
• For Cobalt-60
• Cobalt-59 = 32 neutrons, Cobalt-60 = 33 neutrons

Average atomic mass is the atomic mass that appears on the periodic table.
For example –
Copper has an average atomic mass of
63.55 amu.

Yet, in nature, most elements are found as mixtures of two or more isotopes. For example, copper consists of
– 69.17% copper-63 which has a relative atomic mass of 62.94 amu
AND
– 30.83% copper-65 which has a relative atomic mass of 64.93 amu

To find the average atomic mass, multiply the decimal equivalent of the percent (for example
69.17% = 0.6917) of each isotope by the respective relative atomic mass and add the results.
Isotope
Relative abundance in nature
Relative atomic mass
Copper – 63
69.17%
62.94 amu
Copper – 65
30.83%
64.93 amu
(0.6917 X 62.94 amu) + (0.3083 X 64.93 amu) = 63.55 amu

Practice Calculating Average Atomic Mass
• Boron – 10 is found 19.9% of the time in nature and has a relative atomic mass of
10.013 amu
• Boron – 11 is found 80.1% of the time in nature and has a relative atomic mass of
11.009 amu
• Calculate the average atomic mass of Boron

Practice Calculating Average Atomic Mass
Boron
Isotope
Relative abundance in nature
Relative atomic mass
Boron – 10
19.9%
10.013 amu
Boron – 11
80.1%
11.009 amu
(0.199 X 10.013) + (0.801 X 11.009) =
10.81 amu

Practice Calculating Average Atomic Mass
• Magnesium – 24 is found 78.99% of the time in nature and has a relative atomic mass of
23.985042 amu
• Magnesium – 25 is found 10.00% of the time in nature and has a relative atomic mass of
24.985837 amu
• Magnesium – 26 is found 11.01% of the time in nature and has a relative atomic mass of
25.982593 amu
• Calculate the average atomic mass of Magnesium

Practice Calculating Average Atomic Mass
Magnesium
Isotope
Relative abundance in nature
Magnesium – 24
78.99%
Relative atomic mass 23.985042 amu
Magnesium – 25
10.00%
24.985837 amu
Magnesium – 26
11.01%
25.982593 amu
(0.7899 X 23.985042) + (0.1000 X 24.985837) +
(0.1101 X 25.982593) =
24.306 amu