1. Chapter 3 – VCE Chemistry
The Masses of Atoms
Chapter 3 – VCE Chemistry

2. Atomic Weights and Tables
Accuracy in Chemistry is vital when measuring substances.
The invention of measuring scales enabled more accuracy in scientific research.
John Dalton proposed that atoms could be compared with each by order of mass.
This form of measurement is known as a relative scale (ie. How heavy atoms are relative to other atoms).
Over time, tables of atomic weights were devised.
In 1826, Jons Berzelius a table of 50 known elements.
Berzelius purified, analysed and weighed more than 2000 compounds over a 10 year period.
Berzelius

3. Berzelius’s Atomic Weights

4. Weight and Mass
Mass (m) – How heavy any object is relative to 1 kilogram (kg).
Eg. A 60kg person is 60 times as heavy as a 1 kg object.
The mass of an atom is based on the sum of its subatomic particles.
The greater the number of particles, the heavier the atom.
Weight (W) – Is a measure the gravitational force acting on an object.
Force is measured in Newtons (N).
On earth, objects experience 9.8N/kg.
Eg. A 60kg person has a weight of 588N on the earth’s surface.
The international system of units such as m, W and N are called SI units.
This international prototype, made of platinum-iridium, for 1 kg.
Note. 1kg = 1dm3 of water.

5. Isotopes
Soddy
Studies of uranium and thorium found that different forms of these elements existed.
At that time, they thought that new elements were being formed.
However in 1913 Fredrick Soddy, termed these different elements isotopes.
Isotopes are chemically identical.
Uranium

6. Mass Spectra
Aston
The invention of the mass spectrograph by British scientist Francis Aston enabled chemists to isolate each isotope by mass and abundance.
Modern versions of the instrument are called mass spectrometers.
The output of these instruments is termed mass spectrum.

7. Confusing Evidence
The work of Rutherford concluded that the nucleus of an atom contained positively charged particles called protons.
Hydrogen has one proton, whereas Helium has two.
However, the mass of Helium was four times the mass of Hydrogen – not double.
Rutherford incorrectly predicted that the nucleus must also contain electrons to allow for the difference in mass.
Unknown to Rutherford

8. Discovering the Neutron
It wasn’t until 1932, that James Chadwick proved that a neutrally particle called neutrons exist in the nucleus.
Chadwick’s experiments investigated the effects of alpha particles on a sample of Beryllium.
He discovered that some of the particles were not deflected by magnetic or electrical fields.
He concluded that some particles have no charge. These particles within the atom are called neutrons.
Chadwick also measured the relative mass of neutrons.
Chadwick

9. Summary of Sub-atomic Particles

10. Mass Numbers and Symbols of Isotopes
The number of neutrons in isotopes vary, thus isotopes have different masses.
Protons and neutrons are called nucleons due to their location in the nucleus.
The mass number (A) is the sum of the relative masses of the nucleons.
The number of neutrons can be easily calculated by subtracting the mass number from the atomic number.

11. Questions 3.1
Complete questions 1-4 on P of Nelson VCE Chemistry.

12. The Relative Mass Scale
Experiments by Thomson and Chadwick led to the current accepted masses of subatomic particles.
These values are far to small to be useful in a laboratory – therefore we use a relative mass scale based on carbon-12.
The relative isotopic mass (RIM) of each element can therefore be used in comparison to carbon-12.

13. Relative Atomic Mass

14. Units of the Relative Mass Scale / RAM
Although not a SI unit, chemists use the symbol u (or amu) for relative mass.
1 u = 1/12 the mass of carbon-12.
An alternative unit is the dalton (Da), which was named after John Dalton.
Note. Until 1961, the constant was based on oxygen-16, not carbon-12.
The mass of an element is dependent on the mass and abundance of it’s isotopes present.
The mass of an element is therefore a weighted mean.
The weighted mean of the masses of the isotopes of an element on a scale which carbon-12 is 12 is known as the relative atomic mass (RAM) of the element.
The symbol for RAM is Ar.

15. Isotopic composition of some common elements

16. Mass Spectrum Analysis
Mass spectrum of atomic
chlorine.
A simplified mass
spectrum of
magnesium.

18. Atomic Mass Example Step 1 – There are 3 isotopes
Step 2 – Abundance = peak height/total peak height
Mg-24 = 39/47 Mg-25 = 4/47 Mg-25 = 4/47
Step 3 – The X shows the predicted RAM
Step 4 – Multiply abundance by mass for each isotope.
Ar(Mg) = 39/47x24 + 4/47x25 + 4/47x26
Answer = 24.3 amu

19. RAM and the Periodic Table

20. Questions 3.2
Complete questions 1-11 on P of Nelson VCE Chemistry.

21. Relative Molecular Mass
By knowing the mass of elements, we can calculate the relative mass of molecules.
The relative molecular mass [RMM] (Symbol Mr) is the mean mass of one molecule relative to carbon-12.

22. Calculating Relative Molecular Mass
Mr(H2) = 1.0 x 2 = 2.0 amu
Mr(O2) = 16.0 x 2 = 32.0 amu
Mr(CO2) = x 2 =44.0 amu
Mr(H20) = 1.0 x = 18.0 amu
Mr(NH3) = x 3 = 17.0 amu
Mr(C5H12) = 12.0 x x 12 = 72.0 amu

23. Questions 3.3
Complete questions 1-2 on P. 58 of Nelson VCE Chemistry.

24. Theoretical Percentage Composition
The theoretical percentage composition by mass of a compound gives us the percentage by mass each element contributes.
% element = relative mass of element / relative mass of molecule or compound x 100
Example – H20
%H = 2/18 x 100 = 11.1%
%O = 16/18 x 100 = 88.9%

25. Applying Percentage Composition
Composition can also be translated to mass instead of percentages.
100g of water for example is made up of 11.1g hydrogen and 88.9g oxygen.

26. Questions 3.4
Complete questions 1-4 on P. 61 of Nelson VCE Chemistry.

27. Review Questions
Complete questions 1-10 on P. 64 of Nelson VCE Chemistry.