1. ATOMIC STRUCTURE
Sargus

2. (greek for indivisible)
HISTORY OF THE ATOM
Democritus develops the idea of atoms
460 BC
he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called
ATOMA
(greek for indivisible)

3. Antoine Lavoisier 1743 -1794 Law of Conservation of Mass
Law of Definite Proportions
Law of Multiple Proportions

4. HISTORY OF THE ATOM ATOMS John Dalton 1808
suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them
ATOMS

5. Dalton’s Postulates All matter is composed of atoms
The atoms of a given element are all identical* and different from atoms of another element
Atoms cannot be created, divided, or destroyed by chemical processes.
Chemical reactions change the way atoms are combined

6. Humphrey Davy
Developed electrolysis where a battery was used to apply an electric field to decompose compounds into elements
Showed that atoms are held together by electrical forces

7. Faraday and Stoney
Michael Faraday’s work on electricity also showed that electrical charge is associated with chemical bonds
In 1874 George Stoney determines that discrete units of electrical charge are associated with atoms, and in 1891 he coins the term “electrons”.

8. HISTORY OF THE ATOM Joseph John Thompson 1898
Performs various experiments in which he manipulates a cathode ray tube to demonstrate that:
-cathode rays are made up of electrons
-electrons have mass
-electrons have a negative charge. He determined the charge to mass ratio of electrons

9. HISTORY OF THE ATOM PLUM PUDDING MODEL 1904
Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge
like plums surrounded by pudding.
PLUM PUDDING
MODEL

10. Robert Millikan 1909
Performs the famous “Millikan oil drop experiment” in which he experimentally determined the charge of an electron × 10−19 coulomb
Using Thompson’s charge to mass ratio allows for the determination of an electron’s mass.

11. Cautionary tale in science R. Feynman
We have learned a lot from experience about how to handle some of the ways we fool ourselves. One example: Millikan measured the charge on an electron by an experiment with falling oil drops, and got an answer which we now know not to be quite right. It's a little bit off because he had the incorrect value for the viscosity of air. It's interesting to look at the history of measurements of the charge of an electron, after Millikan. If you plot them as a function of time, you find that one is a little bit bigger than Millikan's, and the next one's a little bit bigger than that, and the next one's a little bit bigger than that, until finally they settle down to a number which is higher.
Why didn't they discover the new number was higher right away? It's a thing that scientists are ashamed of - this history - because it's apparent that people did things like this: When they got a number that was too high above Millikan's, they thought something must be wrong - and they would look for and find a reason why something might be wrong. When they got a number close to Millikan's value they didn't look so hard. And so they eliminated the numbers that were too far off, and did other things like that. We've learned those tricks nowadays, and now we don't have that kind of a disease.

12. HISTORY OF THE ATOM Ernest Rutherford 1910
oversaw Geiger and Marsden carrying out his famous experiment.
they fired Helium nuclei at a piece of gold foil which was only a few atoms thick.
they found that although most of them passed through. About 1 in 10,000 hit

13. HISTORY OF THE ATOM
gold foil
helium nuclei
helium nuclei
They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back.

14. HISTORY OF THE ATOM However, this was not the end of the story.
Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus.
He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction
However, this was not the end of the story.

15. HISTORY OF THE ATOM Niels Bohr
studied under Rutherford at the Victoria University in Manchester.
Niels Bohr
1913
Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.
His attempts in explaining the atomic emission spectrum of hydrogen led to the refutation of the idea of orbits and the development of the quantum model of the atom.
Bohr theorized that electrons occupy quantized energy levels within the atom.

16. Bohr’s Atom
electrons in orbits
nucleus

17. HELIUM ATOM + - + - Shell proton neutron electron
What do these particles consist of?

18. ATOMIC STRUCTURE
Particle
Charge
Mass
proton
+ ve charge 1
neutron
No charge 1
electron
-ve charge
nil

19. number of electrons = number of protons
ATOMIC STRUCTURE He 2
Atomic number
the number of protons in an atom 4
Atomic mass
the number of protons and
neutrons in an atom
number of electrons = number of protons

20. ATOMIC STRUCTURE
Electrons are arranged in Energy Levels or Shells around the nucleus of an atom.
first shell a maximum of 2 electrons
second shell a maximum of 8 electrons
third shell a maximum of 8 electrons

21. 1. Electronic Configuration
ATOMIC STRUCTURE
There are two ways to represent the atomic structure of an element or compound;
1. Electronic Configuration
2. Dot & Cross Diagrams

22. ELECTRONIC CONFIGURATION
With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. For example;
Nitrogen
configuration = 2 , 5 7
2 in 1st shell
5 in 2nd shell N
= 7 14

23. ELECTRONIC CONFIGURATION
Write the electronic configuration for the following elements; 20 11 8 Na O Ca a) b) c) 16 23 40
2,8,8,2
2,8,1
2,6 17 14 5 Cl Si B d) e) f) 11 35 28
2,8,7
2,8,4
2,3

24. N DOT & CROSS DIAGRAMS Nitrogen
With Dot & Cross diagrams elements and compounds are represented by Dots or Crosses to show electrons, and circles to show the shells. For example; X
Nitrogen N 7 X X N X X 14 X X

25. DOT & CROSS DIAGRAMS O Cl
Draw the Dot & Cross diagrams for the following elements; X 8 17 X O Cl a) b) X 35 X 16 X X X X X Cl X X X X X X X O X X X X X X X X X X

26. SUMMARY The Atomic Number of an atom = number of
protons in the nucleus.
The Atomic Mass of an atom = number of
Protons + Neutrons in the nucleus.
The number of Protons = Number of Electrons.
Electrons orbit the nucleus in shells.
Each shell can only carry a set number of electrons.