1. Midterm Part 2

2. 32. Valence electrons and atomic numbers. Atomic numbers increase as you move across the periodic table Number of valence electrons also increase Transition elements are exception

3. 33. What are the subatomic particles for an isotope. Ex. Carbon-12 & Carbon-14 6 protons for each Carbon-12 – Neutrons = 6 Carbon-14 – Neutrons = 8

4. 34. Oxidation Number Electrons lost or gained See periodic table Metals (positive) & nonmetals (negative)

5. 35. Fission and Fusion Fission: when 2 nuclei separate(chain reaction) Fusion: when 2 nuclei come together – Ex. Sun Both take small masses and convert into great amounts of energy

6. 36. Uses of nuclear applications Medical – Radioactive tracers – Radiation for cancer treatment Nuclear power plants Nuclear weapons

7. 37. Molecule and Atom Molecule: covalently bonded group of elements – Smallest part of a compound Atom : smallest piece of matter

8. 38. Types of Chemical Bonds Ionic: metal + nonmetal Covalent: nonmetal + nonmetal – Single, double, or triple bonds – Polyatomic ions

9. 39. Chemical bonding and stability To complete outer energy level * remember * 8 valence electrons Exception: Hydrogen and Helium Argon Atom

10. 40. How does covalent bonding cause stability? Ex. Group 14: 4 valence electrons – Would take too much energy to lose or gain 4 electrons – More chemically stable to just share

11. 42. Ionic & Covalent Bonding Ionic Bonds (crystalline) – Metal + Nonmetal – Ex. NaF, MgO, Al 2 O 3 Covalent Bonds (molecular) – Nonmetal + nonmetal – *Prefixes* – Ex. P 3 N 5, PCl 3

12. 43. Writing chemical formulas for Binary Ionic Compounds Write name of positive ion first, then add –ide to end of negative ion Ex. KCl – Potassium Chloride Ex. BaF 2 – Barium Fluoride

13. 44. Balanced chemical equation Balance using only coefficients Same number of atoms of each element on each side 2Mg(s) + O 2 (g)  2Mgo(s)

14. 45. Law of Conservation of Mass Reactants = products Matter is neither created nor destroyed.

15. 46. Endothermic and Exothermic Reactions Endothermic: energy is used, temperature decreases – Ex. Cold pack Exothermic: energy released, temp. increases – Fire

16. 47. Evidence of Chemical Reaction Evolution of gas Precipitate formation Changes in : – Temperature – Color – Smell

17. 48. Synthesis Equation A + B  C CaO(s) + H 2 O  Ca(OH) 2 (aq)

18. 49. Decomposition Equation C  A + B NH 4 NO 3  N 2 O + 2H 2 O

19. 50. Single and Double Displacement Reaction Single Displacement A + BC  AC + B Fe + CuSO 4  FeSO 4 + Cu Double Displacement AB + CD  AD + CB Ba(NO 3 ) 2 + K 2 SO 4  BaSO 4 + 2KNO 3

20. 51. Reaction Rates Affected by: – Temperature – Surface area – Add a catalyst (does not change amount of product)

21. Buoyancy (Archimede’s Principle) Bernoulli’s Principle Pascal’s Principle Heterogeneous Mixture Homogeneous Mixture Kinetic Theory Charles’s law Boyle’s Law

22. Exam Group Names Symbols (S, l, g, & aq)