1. The Atomic Theory of Matter
Chapter Two
The Atomic Theory of Matter

2. Models of the Atom a Historical Perspective

3. Early Greek Theories
Democritus
400 B.C. - Democritus thought matter could not be divided indefinitely.
This led to the idea of atoms in a void.
fire
air
water
earth
Aristotle
350 B.C - Aristotle modified an earlier theory that matter was made of four “elements”: earth, fire, water, air.
Aristotle was wrong. However, his theory persisted for 2000 years.

4. Alchemists
Relied on experimentation not the development of scientific theories.
Developed some useful techniques in distillation, evaporation, crystallisation and filtration
Aristotle rejected the idea of atoms, but accepted and refined the notion of four elements – water, air, fire and earth
One of the major goals of an alchemist was to discover a substance that would turn base metals such as iron, copper, and lead into gold.

5. Chemistry becomes Science
Robert Boyle
Attempted to isolate Aristotle’s four elements
Defined an element as being a substance that could not be broken down into simpler substances
Was able to distinguish clearly between elements, compounds and mixtures

6. Chemistry becomes Science
Antoine Lavoisier
Changed chemistry from a qualitative to a quantitative science through his tin experiment
Showed that the mass of the products in a reaction is equal to the mass of the reactants – proving the law of conservation of mass

7. John Dalton 1800 -Dalton proposed a modern atomic model
based on experimentation not on pure reason.
All matter is made of atoms.
Atoms of an element are identical.
Each element has different atoms.
Atoms of different elements combine in constant ratios to form compounds.
Atoms are rearranged in reactions.
His ideas account for the law of conservation of mass (atoms are neither created nor destroyed) and the law of constant composition (elements combine in fixed ratios).

8. The Structure of Atoms What is an atom made of ?
Discovery of subatomic particle
i) The Discovery of Electrons
*1897 J.J. Thomson- cathode ray experiment
Thomson’s experiment involved the use of cathode-ray tube. When a sufficiently high voltage is applied across the electrode, an electric current flows through the tube from negatively charged electrode ( the cathode) to the positively charged electrode (the anode).

10. Later experiments had shown that cathode-ray can be deflected by electric or magnetic field. Because the beam is produced at a negative electrode and is deflected toward a positive plate, Thomson proposed that cathode rays are negatively charged fundamental particles found in all atoms which, we now called electrons.
Furthermore, because electrons are emitted from electrodes made of many different metals, all these substances must contain electrons. By careful measuring the amount of deflection caused by electric and magnetic fields of known strength, He established the ratio of mass to electric charge for cathode ray that is, m/e = x10-9 g/coulomb.

11. Review Complete revision questions page 32 (1 – 2)
Summarise how subatomic particles were originally discovered through investigations on electricity
What is a cathode ray?

12. Adding Electrons to the Model
Materials, when rubbed, can develop a charge difference. This electricity is called “cathode rays” when passed through an evacuated tube (demos).
These rays have a small mass and are negative.
Thompson noted that these negative subatomic particles were a fundamental part of all atoms.
Dalton’s “Billiard ball” model ( )
Atoms are solid and indivisible.
Thompson “Plum pudding” model (1900)
Negative electrons in a positive framework.
The Rutherford model (around 1910)
Atoms are mostly empty space.
Negative electrons orbit a positive nucleus.

13. Radioactivity
1896 Antoine Becquerel discovered uranium ores emit invisible rays through working with some photographic plates
Marie and Pierre Curie experimented with radioactivity discovering that uranium, radium and polonium were disintegrating over time and emitting radiation

14. Ernest Rutherford Rutherford shot alpha () particles at gold foil.
Zinc sulfide screen
Thin gold foil
Lead block
Radioactive substance
path of invisible -particles
Most particles passed through. So, atoms are mostly empty.
Some positive -particles deflected or bounced back!
Thus, a “nucleus” is positive & holds most of an atom’s mass.

15. Rutherford’s Scattering Experiment
1909 Ernest Rutherford
~ use  particle to study the inner structure of atoms. When he directed a beam of -particles at a thin gold foil, he found that
 The majority of -particles penetrated the foil undeflected.
 Some  particles experienced slightly deflections.
 A few (about one in every 20,000) suffered rather serious
deflections as they penetrated the foil.
 A similar number did not pass through the foil at all, but bounced back in the direction from which they had come.

16. Discovery of the Neutron
James Chadwick
Proposed the existence of uncharged particles through his experiments with Beryllium foil
Showed that neutrons had almost the same mass as protons

17. The Structure of Atoms Therefore
Modern picture of an atom, then, consist of three types of particles-electrons, protons and neutron.
Electric Charge Mass
Particle SI (C ) Atomic SI (g) amu Located
Electron x x x outside nucleus
Proton x x in nucleus
Neutron x in nucleus

18. Conclusion:
Modern physics has revealed successively deeper layers of structure in ordinary matter. Matter is composed, on a tiny scale, of particles called atoms. Atoms are in turn made up of minuscule nuclei surrounded by a cloud of particles called electrons. Nuclei are composed of particles called protons and neutrons, which are themselves made up of even smaller particles called quarks. Quarks are believed to be fundamental, meaning that they cannot be broken up into smaller particles.

19. Review
Complete the revision questions page 36 (3 – 9). Check and review your answers.

20. Isotopes Frederick Soddy
Found ‘two kinds of thorium’ and isolated the difference in the number of neutrons.
The Mass Spectrometer (provides information about;)
The number of isotopes in a given sample of element
The relative isotopic mass of each isotope
The percentage abundance of the isotopes

21. Most elements consist of a mixture of isotopes
Distinguishing the Atoms: atomic numbers, mass numbers, relative isotopic masses and relative atomic masses
Most elements consist of a mixture of isotopes
The relative atomic mass (Ar)of an element represents the average mass of one atom, taking into consideration the number of isotopes of the element, their relative isotopic mass (RIM) and their relative abundance.

22. Calculating Isotopic Mass
(RIM first isotope x abundance)+(RIM second isotope x abundance)
Ar = ______________________________________________________
100

23. Calculating Isotopic Mass
Using the data in the table calculate the RIM of Cl 37
Isotope
RIM
% Abundance
Cl 35 (Z = 17)
34.97
75.80
Cl 37 (Z = 17)
unknown
24.20
(RIM first isotope x abundance)+(RIM second isotope x abundance)
Ar (Cl) = ______________________________________________________
100
Check your answer and workings on page 38

24. Review
Complete the revision questions page 39 (10 – 18). Check and review your answers.

25. Electrons in an atom
Modern atomic theory is concerned primarily with electron arrangement around the nucleus and has helped the understanding of the chemical properties and behaviour of atoms.
In chemical reactions, when atoms react bonds are broken and atoms rearrange themselves to form new bonds.
The formation of new bonds involves the redistribution of electrons.

26. Bohr’s model Electrons orbit the nucleus in “shells”
Electrons can be bumped up to a higher shell if hit by an electron or a photon of light.
There are 2 types of spectra: continuous spectra & line spectra. It’s when electrons fall back down that they release a photon. These jumps down from “shell” to “shell” account for the line spectra seen in gas discharge tubes (through spectroscopes).

27. Neils Bohr and the Hydrogen Spectrum
If atoms emit only discrete wavelengths, they must have discrete energies.
Each orbit corresponds to a different energy level
An electron can move between energy levels but cannot remain between levels
Quantum jump – moving from one energy level to another
A specific quantity of energy (a photon) is associated with each quantum jump made by an electron

28. Neils Bohr and the Hydrogen Spectrum
An electron can move from a lower energy level to a higher level by absorbing energy (eg flame)
When the electron falls back to a lower level, energy is released as a consequence of the quantum jump.
The energy given out is the difference in energy between the two energy levels and will be associated with a specific wavelength.
Each wavelength corresponds to a coloured line in the emission spectrum.

29. The Hydrogen Spectrum

30. Quantum Mechanical Model
In the Quantum Mechanical Model an electron around a nucleus may be visualised as a cloud of negative charge.

31. Quantum Mechanical Model
The energy levels of electrons are designated by principal quantum numbers, n, and are assigned specific values : n = 1, 2, 3, 4, 5 etc. These principal quantum numbers may be referred to as shells and are also called the K, L, M, and N shells.
Within each shell, several different subshells exist.
The number of subshells equals the shell number
Eg Shell number 2, has 2 subshells
Each subshell corresponds to a different electron cloud shape
Subshells are represented by the letters: s, p, d, f

32. Electron Configuration
The way in which electrons are arranged around the nucleus – generally lowest energy first
Note that the 4s subshell is filled before the 3d subshell (3d is of a higher energy)

33. Electron Configuration
Exited States
When an atom moves to a higher energy level than the ground state by absorbing energy, its electron configuration changes
The outermost electron moves to a higher energy level subshell
Eg Ne (ground state) 1s2 2s2 2p6
Eg Ne (exited state) 1s2 2s2 2p5 3s1

34. Electron Configuration
The group and period in which an element is found is easily read from the electron configuration
(is the number of valence electrons in the last shell being filled)
Eg Be (Z = 4) 1s2 2s2
(the highest shell number being filled)
Group ? _______
Ar (Z = 18) 1s2 2s2 2p6 3s2 3p6
Period ? _______

35. Electron Configuration
Group number is found by using the total number of electrons in the last shell to count across the periodic table
The period is the number of the last shell occupied by electrons
Work through Sample Problems 2.2 and 2.3

36. Electron structure
Nucleus
Consider an atom of Potassium: K 19 39
Potassium has 19 electrons. These are arranged in shells…
The inner shell has __ electrons
The next shell has __ electrons
The next shell has the remaining __ electron
Electron structure
= 2,8,8,1

37. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Hydrogen H
1 electron
0 electron
Helium He
2 electron

38. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Lithium Li
2 electron
1 electron
0 electron
Beryllium Be

39. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Boron B
2 electron
3 electron
0 electron
Carbon C
4 electron

40. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Nitrogen N
2 electron
5 electron
0 electron
Oxygen O
6 electron

41. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Fluorine F
2 electron
7 electron
0 electron
Neon Ne
8 electron

42. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Sodium Na
2 electron
8 electron
1 electron
0 electron
Magnesium Mg

43. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Aluminium Al
2 electron
8 electron
3 electron
0 electron
Silicon Si
4 electron

44. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Phosphorus P
2 electron
8 electron
5 electron
0 electron
Sulphur S
6 electron

45. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Chlorine Cl
2 electron
8 electron
7 electron
0 electron
Argon Ar

46. How the shells fill with electrons
Element
Shell 1
Shell 2
Shell 3
Shell 4
Potassium
2 electron
8 electron
1 electron
Calcium Ca

47. The First Twenty Elements
Hydrogen 1,0,0,0
Helium ,0,0,0
Lithium ,1,0,0
Beryllium 2,2,0,0
Boron ,3,0,0
Carbon ,4,0,0
Nitrogen ,5,0,0

48. First 20 Elements continued
Oxygen ,6,0,0
Fluorine ,7,0,0
Neon ,8,0,0
Sodium ,8,1,0
Magnesium 2,8,2,0
Aluminium 2,8,3,0
Silicon ,8,4,0

49. First 20 Elements continued
Phosphorus ,8,5,0
Sulphur ,8,6,0
Chlorine ,8,7,0
Argon ,8,8,0
Potassium ,8,8,1
Calcium ,8,8,2

50. Review Complete the revision questions pages 44, 45 (21 – 26)
Work through the Sample Problems 2.4, 2.5
Complete the revision questions page 45 (27 – 31)

51. Atomic numbers, Mass numbers
There are 3 types of subatomic particles. We already know about electrons (e–) & protons (p+). Neutrons (n0) were also shown to exist (1930s).
They have: no charge, a mass similar to protons
Elements are often symbolized with their mass number and atomic number
E.g. Oxygen: O 16 8
These values are given on the periodic table.
For now, round the mass # to a whole number.
These numbers tell you a lot about atoms.
# of protons = # of electrons = atomic number
# of neutrons = mass number – atomic number
Calculate # of e–, n0, p+ for Ca, Ar, and Br.

52. Atomic
Mass p+ n0 e– Ca 20 40 20 20 20 Ar 18 40 18 22 18 Br 35 80 35 45 35

53. Bohr - Rutherford diagrams
Putting all this together, we get B-R diagrams
To draw them you must know the # of protons, neutrons, and electrons (2,8,8,2 filling order)
Draw protons (p+), (n0) in circle (i.e. “nucleus”)
Draw electrons around in shells He Li
3 p+
4 n0
2e– 1e–
Li shorthand
3 p+
4 n0
2 p+
2 n0
Draw Be, B, Al and shorthand diagrams for O, Na

54. Be B Al O Na 8 p+ 11 p+ 8 n° 12 n° 4 p+ 5 n° 5 p+ 6 n° 13 p+ 14 n°
2e– 8e– 1e– Na
8 p+
8 n°
2e– 6e– O

55. Isotopes and Radioisotopes
Atoms of the same element that have different numbers of neutrons are called isotopes.
Due to isotopes, mass #s are not round #s.
Li (6.9) is made up of both 6Li and 7Li.
Often, at least one isotope is unstable.
It breaks down, releasing radioactivity.
These types of isotopes are called radioisotopes
Q- Sometimes an isotope is written without its atomic number - e.g. 35S (or S-35). Why?
Q- Draw B-R diagrams for the two Li isotopes.
A- The atomic # of an element doesn’t change Although the number of neutrons can vary, atoms have definite numbers of protons.

56. For more lessons, visit www.chalkbored.com
6Li 7Li
3 p+
3 n0
2e– 1e–
3 p+
4 n0
2e– 1e–
For more lessons, visit

57. Chapter 2 Review Practice your electron dot diagrams here
Complete the multiple choice questions pages 48, 49
Consider each of the review questions