1. Atomic Structure

2. History of the atom
Robert Boyle (1st to define an element) believed all matter was made up of tiny particles
150 years later John Dalton proposed his atomic theory based on experiments he carried out

3. Dalton’s Atomic Theory:
All matter consists of tiny particles called atoms
All atoms are indivisble so cannot be split into simpler particles
Atoms cannot be created or destroyed

4. Discovery of the electron:
William Crookes worked with discharge tubes studying if gases conduct electricity

5. When electricity passes through gases they form invisible rays that travel from the negative electrode (cathode) to the positive electrode (anode)
Crookes called these rays cathode rays

6. John Thomson studied cathode rays
He found they were rays of negative particles that were lighter than a hydrogen atom (the smallest particle known then) so particles smaller than atoms must exist
He also discovered the e/m (charge to mass) ratio of the electron
Cathode rays are streams of negatively charged particles called electrons

7. Milikan’s Oil Drop experiment:
Robert Milikan carried out an experiment to measure the negative charge of an electron
He sprayed a fine mist of oil between oppositely charged plates
He passed electricity through the oil droplets to ionise them (make them charged)

9. He focused on one oil drop and adjusted the charge on the plates until the droplet stayed suspended in mid air
The electrical force on the plate was then equal to the gravitational force on the droplet
He could then calculate the charge on the electron

10. Thomson’s Plum Pudding Model:
He knew atoms were neutral
He knew they contained negative particles called electrons
He thought there must be a positive charge to balance the negative one

11. Thomson thought the atom looked like a plum pudding
A large sphere of positive charge with negative electrons stuck in it
Like fruit in a plum pudding

13. Discovery of the nucleus:
Ernest Rutherford carried out the Gold Foil experiment
He pointed alpha particles at sheets of gold foil
He used a Geiger-Muller tube to detect them on the other side

15. Based on Thomson’s plum pudding model Rutherford made the following assumptions:
1) most particles would pass straight through because the atom is full of empty space
2) some particles would be slightly deflected by electrons

16. What actually happened (his results):
1) most particles passed straight through
2) some were deflected at small angles
3) some were deflected at very large angles
4) some were deflected back on their paths

17. Expected:
Actual:

18. Rutherford’s conclusions were:
1) most of the atom is empty space
2) the atom does contain electrons
3) the atom has a large positive mass at its centre – called the nucleus

19. Discovery of the proton:
Rutherford continued his experiment using lighter elements like N and F
He found that these elements released positive particles when hit with alpha particles
He called these positive particles protons

20. Discovery of the neutron:
Chadwick (one of Rutherford’s researchers) carried out the same experiment with Be
He found that it released neutral particles when hit with alpha particles
He called these neutral particles neutrons

21. Sub-atomic particles:
Mass:
Charge:
Location:
Electron
1/1836 a.m.u. -1
Outside nucleus
Proton
1 a.m.u. +1
Nucleus
Neutron

22. Atomic Structure.
Atoms consist of a nucleus at the centre, containing protons and neutrons, with electrons orbiting the nucleus in shells.
Atomic number = Number of protons in the nucleus.
Mass number = Number of protons plus number of neutrons in nucleus.

23. No. electrons = no. protons No. neutrons = mass no. – atomic no.
No. protons = atomic no.
No. electrons = no. protons
No. neutrons = mass no. – atomic no.
Element:
No. electrons:
No. protons:
No. neutrons H 1 He 2 Li 3 4
----

24. Isotopes
Isotopes are atoms of the same element that have different mass numbers due to the number of neutrons in the nucleus,
e.g., H-1, H-2 and H-3 or C-12 and C-14
Relative atomic mass (Ar) is the average mass number of the isotopes of the element compared to 1/12th the mass of the carbon-12 isotope.
Mass spectrometer is used to measure Ar.

26. Mass Spectrometer Principle –
Positive ions are separated according to their mass in a magnetic field
The lighter the ions the more they will be deflected by the magnetic field
1) Vaporisation:
Sample is vaporised and injected into sample inlet

27. 2) Ionisation: 3) Acceleration:
Sample passes through the electron gun where it is
bombarded with electrons causing electrons in the
sample atoms to fall off so they become positively
charged
3) Acceleration:
Ions pass through oppositely charged plates where
they are attracted to the negative plate and are
accelerated to high speeds

28. 4) Separation:
Ions pass into a magnetic field where they are
separated based on their mass. Lighter particles
are easy to deflect compared to heavier ones. The
magnetic field can be adjusted so that all isotopes
present can be measured.
5) Detection:
The detector responds to the number of ions that
hit it which are then registered on a chart called a
mass spectrum – relative abundance peaks are
shown for the different relative atomic masses

29. Mass Spectrum for chlorine

30. Mass spectrum for aspirin

31. Uses of the mass spectrometer:
Identify presence of isotopes
Measure relative abundance of isotopes
Measure relative atomic & relative molecular mass
- Identify unknown compounds

32. Calculations: Q:
A mass spectrometer is used to measure a
sample of boron. It is found that the sample
contains 18.7% B-10 and 81.3% B-11.
Calculate the relative atomic mass of boron. A:
18.7 x 10 = 187
81.3 x 11 = 894.3
100 atoms =
1 atom = Ar =

33. Q:
A sample of chlorine consists of 75% Cl- 35
and 25% Cl-37. Calculate the average mass
of an atom of chlorine. A:
75 x 35 = 2625
25 x 37 = 925
100 atoms = 3550
1 atom = Ar = 35.5