1. 2.1 Atoms and Their Structure
Democritus – 460 – 370 BC
Greek philosophers thought all matter was made up of these 4 ‘elements’: earth, air, fire and water
World made up of empty space and tiny particles called atomos
atoms (atomos) means ‘indivisible’ and indestructible.
His work not based on scientific method.

2. Lavoisier, Antoine (1743 – 1794)
First Modern Chemist – ‘Father of Chemistry’
Concluded through experimentation that when a chemical reaction occurs
“Matter is neither created nor Destroyed”
Law of Conservation of Matter
Atoms are always ‘recycled’ in our Universe
Studied gases – Air is a ‘mixture’ O2 and N2
He Gave Oxygen its’ name.

3. Proust, Joseph
Known for the ‘Law of Definite Proportions’
Elements that make up “compounds are always found in a certain percentage by mass”
He observed that water is always 11% hydrogen and 89% oxygen by mass.

4. John Dalton’s Atomic Theory
(1766 – 1844)
All matter is made up of atoms.
Atoms are indivisible – (later proved wrong)
Atoms of the same element are exactly alike (later proved wrong)
and are different from atoms of other elements.

5. JJ Thomson
– 1897
Discovered the first subatomic particle – the electron.
Discovered the electron using a vacuum tube with a positive end (+ anode) and a negative end (- cathode).
Invisible, negatively charged particles now known as “electrons”.
“Plum Pudding Model”
Some atoms have different masses
He proved that the atom was not a ‘solid ball’.

6. Cathode Ray Tube – negatively charged particles attracted towards the positive end

7. Robert A. Millikan U.S. Physicist (1868 – 1953)
Known for the “Oil Drop Experiment”
Using the charge to mass ratio of an electron he calculated the mass of an electron.
Has a mass of 1/1,840 the mass of an hydrogen atom.
(Almost 2000 times lighter than a Proton)

8. Ernest Rutherford (1871 – 1937) Known for the Gold Foil Experiment
In 1909 Discovered the nucleus of an atom.
Nucleus - very dense, positively charged center of the atom.
Experiment: Fired stream of positively charged particles at a thin sheet of gold foil.
Determined that atoms are mostly empty space.
Evidence that the nucleus of an atom is P+ positive.

9. Internet Links – Animations
Example of JJ Thomson’s experiment
Example of Rutherford’s Experiment
Another Web site illustrating Gold Foil Experiment

10. Protons Positively charged (+) subatomic particles.
2000 times greater mass than an electron.
Just slightly less mass than a proton

11. Isotopes
Atoms of the same element that are chemically alike but differ in mass
Ne-20 and Ne-22

12. Neutrons – no charge
The discovery of isotopes led scientists to hypothesis that atoms must contain a third type of particle approximately equal in mass to a proton but neutral in charge.
Called Neutrons

13. James Chadwick English Physicist (1891-1974)
Discovered the neutron in 1932.
His discovery helped the development of splitting atoms in half.
Making possible the development of the atomic bomb. (fission of uranium 235)
Revolutionized the study of radioactive elements.

14. Electron Microscopy
Within 30 years of J.J. Thomson’s discovery of the electron scientists were producing images of objects using an electron beam.

15. Niels Bohr (1885-1962) Danish physicist
His work led to our understanding atomic structure and quantum theory, he received the Nobel Prize in Physics in 1922.
In 1913 he proposed the ‘planetary model’.
Electrons have energy of motion.
Electrons absorb energy and move to higher energy states.
Electrons give off that energy in the form of light when they fall back down to lower energy states.

16. Isotopes
Atoms of the same element can have different number of neutrons.
They have different masses.
Same chemical properties because they have the same
number of protons.
Example: 3 isotopes of Hydrogen. Protium H- 1
Deuterium H- 2
Tritium H-3

17. Atomic Number The number of protons in nucleus.
The Atomic Number determines what the element is.
Each element has its own number.
“Like a fingerprint”
Also tells how many electrons in a neutral atom of that element.
47 protons and also 47 electrons

18. Atomic Mass Number
The average number of protons and neutrons in any given sample.
A weighted average mass of all naturally occurring isotopes of that element.
Atomic Mass Number

19. Mass Number - sum (add) of protons and neutrons in the nucleus in an atom.
Whole number – not decimal
Used for isotopes
Compared to the Periodic Table the numbers are flipped upside down
Mass Number is also written like this:
Mass Number
C-12 and C-14

20. Isotopes of Carbon
Any sample in nature naturally contains a mixture of isotopes.
It is usually not a whole number.
Example: atomic mass of carbon =

21. Atomic Mass 1 μ = 1/12 the mass of a carbon-12 atom
Easy way to express the mass of atoms: Compare the mass to a known mass of carbon.
Atomic mass unit (amu), symbol = μ.
Carbon is the standard for atomic masses.
12 (amu) or 12 μ.
(6 p+ and 6 n0) = 12 amu
1 μ = 1/12 the mass of a carbon-12 atom
1 u is approximately the mass of a single proton or neutron.
Mass of a Proton = 1 amu
Mass of a Neutron ~ 1 amu
Symbol
Standard for mass
Mass of Carbon – 12
Unit

22. To Calculate the Atomic Mass of an Element:
1) Multiply the mass of each isotope by its abundance (percentage naturally found in nature) expressed as a decimal.
2) Then add the products.
Example: Atomic mass of carbon
( amu x ) + ( amu x ) = amu

23. How are atoms measured?
Mass spectrometry is an analytical technique that measures the mass-to-charge ratio of charged particles.

24. Hypothesis – testable prediction that explains observations
Theory – explanations based on many observations and supported by many experiments
Scientific Law - a fact of nature observed so many times that it is accepted as truth.
Theories Explain Laws -
Methods of Science