1. Chapter 2 & 3: The Atomic Theory of Matter and Numerical connections:

2. History of Atomic Theory

3. Structure of the atom Protons:- Found in nucleus Positive charge (+ve) Mass = 1 Number of protons is known as the Atomic number. Each element has a different number of protons. Protons are made form quarks and gluons. Neutrons:- Found in nucleus No charge Mass = 1 Neutrons act to stabilise the nucleus and hold the protons together with the strong force. The Relative atomic mass is the number of protons + number of neutrons. Neutrons are made form quarks and gluons. Electrons:- Found in (orbits) shells around the nucleus. Each shell and sub-shell represent a particular energy level that the electrons are found at. These energy levels can be explained by quantum mechanics. The number of electrons equals the number of protons unless the atom forms an ion.

4. Isotopes Mass Spectrometer Most elements form isotopes An Isotope is an element that has atoms that have a different number of neutrons. The RAM on the periodic table is an average for that element taking into account the abundance of the isotopes for that element. Isotope calculations are an important part of chemistry. Page 38 of your text gives you the formula and an example of these types of calculations. A mass spectrometer can be used to detect the presence and number of isotopes present in a sample of an element. Mass spectrum of Bromine molecules

5. Shells and Sub Shells

6. The Mole The mole is a comparable measurement of elements and compounds used by chemists. Atoms are so small that an easier method of counting and comparing them was needed. The mole was developed based on the theory that exactly 12 grams of Carbon-12 has experimentally be determined to have 602000000000000000000000 atoms. 6.02 x 10 23 is known as Avagadro’s number.

7. The mole continued... One mole (n) of any element is equal to its RAM and contains 6.02 x 10 23 atoms of that element. For compounds the RAM’s of the elements present are added together, known as the Molar Mass (M) Eg: One mole of H 2 O = 1 + 1 + 16 = 18 gmol -1 In most cases you are not dealing with 1 mole of the chemical you are using!

8. Mole calculations n = m M Number of mol = mass of chemical in grams divided by its Molar Mass. m = n x MMass of a chemical in grams = number of mol x Molar mass. M = m n Molar Mass of unknown chemical = mass in grams ÷ number of mole Particles = n x N A Particles present = number of mole x Avogadro's number. n = Particles N A Number of mole = Particles present ÷ Avogadro’s Number.

9. Percentage composition Percentage composition tells us the amount by mass of each element present in a compound. Molar mass = (39.1 x 2) + (52 x 2) + (16 x 7) = 294.2 gmol -1 % Oxygen = (16 x 7) x 100 294.2 1 = 38%

10. Empirical formula The simplest whole number ratio of atoms present in a compound. H 2 O C 2 H 3 O 4 C 6 H 12 O 6 You must be able to perform calculations to determine empirical formula. Refer to page 59-60 of text for method and examples.

11. Molecular formula The molecular formula is the actual formula of a chemical which is not always in its simplest form (empirical formula) Glucose has the molecular formula C 6 H 12 O 6. Its Empirical formula is CH 2 O. An explanation and examples of calculations of Molecular formula are found on pages 61 – 63 of your text book.

12. What do I need to be able to work out? Electron arrangement – really important Number of mole – most important Mass from number of mole – also important Particles present using N A – not used much % composition – not used much Empirical formula – important Molecular formula - important

13. The faster the circles move the more stressed you are. If they don’t move you are either very relaxed or dead. If this last chapter is doing your head in please stick with it. Lots of Chemistry students don’t really get this stuff until year 12 or even Uni. It will be revised heaps and is repeated in Unit 2, 3 and 4.