1. Review

2. Various models of the atom Dalton’s – “billiard ball” model Thomson – “raisin bun” model Rutherford – “Electron Cloud” model Bohr – best model to date

3. Molecular Elements Elements ending in –gen are diatomic. Pg. 33 – should know this table

4. Changes Pg. 47

5. Classifying Chemical Reactions Pg. 58

6. Significant Digits Precision is the place value of the last measurable digit and is determined by the instrument 12.0 g vs 12 g vs 10 g Non-zero numbers are significant Zeros to the left are not significant Zeros to the right are significant

7. Compounds Valence Electrons Lewis Symbols Electronegativity Types of Bonds Bonding Capacities

8. Pg. 88

9. Predicting Shape Summary

10. Rules of polar and nonpolar

11. Electronegativity Differences Pg. 100

12. Summary Intermolecular forces are the attraction and repulsion of positive and negative charges All molecules have London forces – with momentary diploes Polar molecules have dipole-dipole forces Hydrogen bonding occurs when a hydrogen is attracted to the lone pair from an adjacent molecule IMF affect many various physical properties

13. Physical properties While solid has definite shape and volume; it does have various other properties that can be very different. Such as density, colour, melting points, etc Pg. 119

14. Summary

15. Unit Test Review Summary Lewis Formulas Bonding Theory (types) Electronegativity Molecular Formulas (Lewis, structural, sterochemical) Bonding capacity VSEPR Bond and molecular Polarity Intermolecular Forces – Isoelectronic Compounds Boiling points Properties of Liquids and Solids

16. Units Pg. 149

17. Summary STP: 0 ˚ C and 101.325 kPa SATP: 25 ˚ C and 100. kPa 101.325 kPa = 1 atm = 760 mm Hg T (K) = t ( ˚ C) + 273 P 1 V 1 = P 2 V 2 V 1 / T 1 = V 2 / T 2

18. Properties of Gases Kinetic molecular theory explains a lot about the properties of gases – Gases are compressible - the distance between the molecules – Gas pressure – the amount of collisions that occur – Boyle’s Law – Charles’ Law

19. Ideal Gas Law

20. Solvent and Solute Solvent – The component of the solution present in greater amount. Solvent dissolves the solute Solute – The component of the solution present in lesser amount. The solute is dissolved by the solvent.

21. Acids and Bases PG 199

22. Summary

23. Strength of Acids Strong Acids:  ionize (splits up into ions) almost 100% in water  mostly ions in solution  amount of HCl present is negligeable HCl(aq)  H + (aq) + Cl - (aq) Weak acids:  ionize poorly in water  not many of these ions present in solution  mostly acetic acid (HC 2 H 3 O 2 ) HC 2 H 3 O 2 (aq) C 2 H 3 O 2 - (aq) + H + (aq) NOTE: strong acids are strong electrolytes and will conduct electricity better than weak acids.

24. pH < 7 acidic pH = 7 neutral pH > 7 basic

25. Expressing hydronium concentrations in scientific notation isn’t very convenient. The pH scale was developed to make the expression of H 3 O + concentration more convenient. [H 3 O + ] is the concentration in mol/L pH = -log[H 3 O + ]

26. Back of book

27. Net Ionic Equations The chemical reaction equation can be written with just the chemicals which will react and form precipitates.

28. Applications of Stoichiometry To determine the overall experiment the percent yield is used. This lets us know about the experimental uncertainties – Measurements, purity, washing, estimations

29. Summary

30. Limiting Reactions There is only a certain amount of one of the reacting chemicals and when all of it is used up the reaction stops – this is the limiting reagent The other reactant is more abundant and is called the excess reagent

31. Limiting and Excess Reagent Having a reactant with excess means that more of the substance is present then required for the reaction. We our purposes we will use the value of 10 % for the excess amount required to make the reaction occur. Sample problem 8.1 pg 320

32. Interpreting Titration pH Curves Pg. 334