2. Atomic Theory In 1808, the English Chemist John Dalton proposed the first theory of the nature of matter in stating that all matter was composed of atoms. Dalton based his theory on two other scientific principles The Law of Conservation of Mass The Law of Definite Proportions (constant composition)

3. The Law of Conservation of Mass 1789 – Antoine Lavoisier Mass of products in a chemical reaction is equal to the mass of the reactants. 1.00 g of carbon + 5.34 g of sulphur = 6.34g carbon disulphide

4. The Law of Constant Composition (Definite Proportion) 1799 – Joseph Proust In a pure compound the elements are always present in the same definite proportion by mass. E.g. 2.0160g hydrogen will combine with 15.9994 oxygen to produce water. And 4.032g hydrogen will combine with 31.998 g oxygen Same proportion by mass

5. Dalton’s Atomic Theory All matter is made up of small particles called atoms Atoms cannot be divided Atoms can neither be created nor destroyed in a chemical reaction All atoms of the same element are identical in mass, size and physical properties The properties of the atoms of one element differ from the atoms of all other elements Atoms combine in small whole number ratios to form compounds.

6. Modern Atomic Theory We know that not all of these are correct

7. Subatomic Particles Electrons Protons Neutrons

8. Discovery of the Electron Experiments with cathode ray tubes Contained gases at low pressure Electric current passed through the gas Caused a glow on the surface of the tube directly opposite the cathode

9. J.J.Thomson hypothesized that the glow was caused by a stream of particle called cathode rays To test it – placed an object between the cathode and the opposite end of the tube cast a shadow on the glass. A paddle wheel was moved by the rays (had mass) Cathode rays were deflected by magnetic fields Rays deflected by negatively charged objects

11. These observations led to the hypothesis that the particles that compose cathode rays are negatively charged J.J. Thomson was able to find the ratio of charge to mass of the particles. He found this was always the same regardless of the metal used to make the cathode, or the gas in the tube. He concluded that all cathode rays are composed of negatively charge particles, later named electrons

12. Millikan’s experiment Millikan found the mass of the electron is about one two-thousandth of the mass of a hydrogen atom. Confirmed electrons are negatively charged – huge charge for its small mass Inference 1: atoms are neutral, so there must be a positive charge to balance the negative. Inference 2: There must be more particles to account for the mass of the atom.

13. Rutherford’s experiment Bombarded alpha particles (positively charged particles with about four times the mass of hydrogen) into gold foil. 1 in 8000 particles was redirected back to the source.

14. He concluded that atoms contain a small dense positively charged nucleus that repels the positive alpha particles. He also concluded that most of the atom was empty space. The volume of the nucleus is very small compared to the volume of an atom – if the nucleus was the size of a marble, then the atom would be the size of a football field.

15. Neutrons were discovered in 1932 by James Chadwick. Their function may be to hold the nucleus together. Why don’t the positively charge protons in the nucleus repel each other and fly out? Because of a very strong force called nuclear force.

16. Atomic number Isotopes Mass number Relative atomic masses One atomic mass unit (amu) is exactly 1/12 of the mass of a carbon-12 atom or 1.660540 x 10 –27 kg Although isotopes have different mass they do not differ significantly in their chemical behavior. Average atomic mass

17. Isotope Notation Mass # Element Atomic # Nuclear symbol # of protons # of electrons # of neutrons Protium ( Hydrogen-1) 11H11H 110 Deuterium (Hydrogen-2) 12H12H111 Tritium (Hydrogen-3) Carbon-12 Carbon-14

18. Mass number and atomic mass Mass number - # of protons + # of neutrons Atomic mass unit –1 amu is exactly 1/12 the mass of a C-12 atom or 1.660540 x 10 -27 kg Atomic mass of any isotope is found by comparing it to carbon-12 For example, hydrogen-1 has an atomic mass of about 1/12 the atomic mass of carbon. The exact value of hydrogen-1 is 1.007825 amu.

19. Mass of subatomic particles Electron0.0005486 amu 9.109 x 10 -31 kg Proton1.007276 amu 1.673 x 10 -27 kg Neutron1.008665 amu 1.675 x 10 -27 kg

20. Average Atomic Mass The weighted average of the atomic mass of the naturally occurring isotopes of an element. For example copper is made of 69.17% of Cu-63 and 30.83% of Cu-64. The atomic mass of copper-63 (Cu-63) is 62.929598 amu And the atomic mass of copper-64 (Cu-64) is 64.927793 amu The weighted average is 0.6917 x 62.929598 + 0.3083 x 64.927793

21. Moles A mole is the amount of substance that contains as many particles as there atoms in exactly 12g of carbon-12. It’s a counting unit (like a dozen) The number of particles in one mole of a pure substance is Avogadro’s number 6.022 x 10 23 Molar mass – the mass of one mole of a pure substance is its molar mass.