1. Unit 4 Atoms, Bonding, and Chemical Reactions Ch 18, 19, and 24

2. CHAPTER 18

3. Structure of the Atom Protons, neutrons, electrons Quarks – small particles that make up protons and neutrons

4. Models Dalton - sphere Thompson – electrons existed Rutherford – nucleus containing + charge surrounded by empty space containing electrons Bohr – electrons travel in orbits around nucleus with protons and neutrons Electron Cloud – electrons not in fixed orbits, but in a cloud around the nucleus Did The Rabbit Bite Eeyore?

5. Using the periodic table Atomic number = # protons Smaller # on periodic table On periodic table, # protons = # electrons Atomic mass = # protons + # neutrons Larger # on period table # neutrons = mass- atomic #

6. What is the atomic number of zinc? How many electrons does tungsten have? How many neutrons does scandium have? What is the atomic mass of carbon? How many protons does astatine have?

7. Isotopes Same number of protons (same element) Different number of neutrons Therefore, different atomic mass

8. Periodic Table structure Periods (left to right) = increasing number of protons and electrons Groups (up and down) = similar reactive properties

9. Energy Levels 1 st Level = hold max of 2 e- 2 nd Level = hold max of 8 e- 3 rd Level = hold max of 8 e-

10. Drawing Old School Way Shows all the electrons on all the energy levels Ex: draw flourine Electron Dot diagram Only shows the outermost electron (the valence electron) Ex: draw flourine

11. Trends Left to right, down to up: Increasing electro negativity Increasing ionization energy Decreasing atomic radius

12. Cheats on your periodic table On your periodic table, write in the Group numbers This is the number of electrons in the outermost level How many outermost electrons does Boron have?

13. Bonding Atoms want a full outermost shell This is when they are most stable Noble Gases (far right of table) already have full outermost shells Other elements want to give up or gain e- to make a full outermost shell If elements lose an e-, they become positively charged If elements gain an e-, they become negatively charged

14. Cheats on your periodic table Oxidation numbers Group 1 = becomes +1 charged ex: Li+1 Group 2 = becomes +2 charged ex: Mg+2 Group 3 = becomes +3 charged ex: Al+3 Group 5 = becomes -3 charged ex: N-3 Group 6 = becomes -2 charged ex: O-2 Group 7 = becomes -1 charged ex: F-1

15. Identify the oxidation numbers for each element: NaCl CaO N2O SiO2

16. CHAPTER 19

17. Ionic vs. Covalent bonding Ionic Total transfer of e- Between metal and nonmetal On both sides of your stairstep line Ex: NaCl Covalent Sharing e- Between a nonmetal and nonmetal Both to the right of the stairstep line Ex: CO

18. Identify if it is ionic or covalent SiO2 LiF NaCl C12H22O11 HCl

19. Polar vs. Nonpolar Polar Atoms have diff electro negativity Electrons not shared equally Ex: HCl Cl is more electronegative then H, therefore stronger negative charge Nonpolar Atoms have same electro negativity Electrons shared equally Ex: Cl2 Same electro negativity

20. CHAPTER 24

21. Chemical Rxns Reactants --> Products Conservation of mass Mass is converted into different forms but never created or destroyed

22. Symbols used in chemical equations ssolid lliquid ggas aqaqueous, dissolved in water

23. Coefficients and subscripts 4H + 02  2H20 Notice how this is balanced Use the distributive property 4 H on left 4 H on right 2 O on left 2 O on right

24. Chemical Equations Balancing equations Subscripts remain the same Coefficient applies to each element Ex: 2HO = 2H + 2O

25. UNL’s tricks to balance! 1. Start with compound with the greatest diversity of atoms 2. Leave pure elements alone until end (usually O or H) 3. If rule #1 doesn’t help, start with the compound farthest left 4. All coefficients must be whole numbers. This may require multiplying by the LCM to get rid of fraction. 5. # atoms of each element must be balanced on both sides of the equation

26. Balance these equations HgO  Hg + O2 Li + H2O  H2 + LiOH Mg + O2  MgO

27. Types of Reactions Synthesis A + B --> AB Ex: 2H2 + O2  2H2O Decomposition AB --> A + B Ex: 2H2O  2H2 + O2 Single Displacement A + BC --> AB + C Ex: Cu + 2AgNO3  Cu(NO3)2 + 2 Ag Double Displacement AB + CD --> AC + BD Ba(NO3)2 + K2SO4  BaSO4 + 2KNO3

28. Energy Exchanges Exergonic rxn = releases energy (EXITs) Ex: glow sticks (releases light) Exothermic rxn = releases heat Ex: burning wood Endergonic rxn = requires energy (moves IN) Endothermic rxn = requires heat Ex: activating a cold pack

29. Catalysts vs. Inhibitors Catalysts Speed up rxns Same product is formed Catalyst remains unchanged and separate from product Enzymes lower the activation E, making the rxn require less E to occur Ex: enzymes break down fruit (looks brown) Inhibitor Prevents rxn from occurring Same product is formed Inhibitor remains unchanged and separate from product Ex: lemon juice keeps fruit from browning