Final Exam Review

1. Final Exam Review

2. Atomic Models

3. Dalton Atomic Theory • 1. Elements are made of atoms. • 2. All atoms of same elements are identical. • 3. Atoms of different elements are different. • 4. Compounds have definite compositions. • 5. Atoms cannot be created or destroyed in chemical • reactions.

4. Rutherford’s Gold Foil Experiment

5. Most of the particles went straight through, a few of them bounced back.

6. Rutherford’s Atom • 1. Atoms are mostly empty space. • 2. Atoms have a small but massive nucleus with a positive charge. • 3. Surrounding the nucleus are electrons, which have very little mass.

7. Atoms, Ions, and Isotopes • Atom: • Made of electrons, protons, & neutrons • Smallest unit of an element • Neutral charge • Ion: • Atom that has lost or gained electrons • Has a + or - charge • Isotope: • Atoms of the same element that have differing numbers of neutrons

8. Subatomic Particles # Protons = Atomic # # Electrons = # Protons (if neutral) # Neutrons = Mass - # Protons If an atom gains electrons, it becomes a negative ion. If an atom loses electrons, it becomes a positive ion. Metals tend to lose electrons and nonmetals tend to gain electrons.

9. Periodic Table e-Configuration

10. Periodic Trends Biggest Radius Most Reactive Metal

11. Periodic Trends Most Electronegative Most Reactive Non-Metal

12. Periodic Trends Highest Ionization Energy

13. More on Periodic Table • Alkali metals = column 1 • Alkali earth metals = column 2 • Halogens (quite reactive) = column 17 • Noble gases (nonreactive to barely reactive) = column 18

14. When you write the electron configuration for any element, just follow this pattern and remember to stop at the element you’re representing. 1s2 1 6p6 2p6 3p6 4p6 5p6 4d10 6d10 5d10 3d10 5f14 4f14 2s2 2 3 3s2 4s2 4 5s2 5 6s2 6 7s2 7

15. 8A 1A 2A 3A 4A 5A 6A 7A • Elements in the 1A-7A groups are called the representative elements outer s or p filling

16. Each sublevel contains the following orbitals: • Sublevel s = 1 orbital 2 electrons • Sublevel p = 3 orbitals 6 electrons • Sublevel d = 5 orbitals 10 electrons • Sublevel f = 7 orbitals 14 electrons

17. Energy LevelSublevels (Type of Orbitals) • n = 1 1s (1) • n = 2 2s (1) 2p (3) • n = 3 3s (1) 3p (3) 3d (5) • n = 4 4s (1) 4p (3) 4d (5) 4f (7) • n = 5 5s (1) 5p (3) 5d (5) 5f (7)

18. Bonding 1) Ionic Compound: Metal and non metal, transfer of electrons 2) Covalent Compound: 2 nonmetals, sharing of electrons Covalent Compounds can be polar or non-polar. Polar: electrons are not shared equally. Non-polar: electrons are equally shared.

19. Metallic Bonding • Metallic bonds are described as a cluster of positive ions surrounded by a sea of their shared valence electrons.

20. Molecular Geometry Bent or angular Linear Tetrahedral Pyramidal vs planar Boron is a moron, it only needs 6 electrons!

22. Polar vs. Non-Polar Bent or pyramidal Angular = Bent

23. IntERmolecular Forces - (Forces Between Molecules) 1. Dipole - Dipole This is when the positive end of one polar molecule is attracted to the negative end of another molecule 2. Hydrogen Bonding Special type of dipole-dipole force. Hydrogen bonding involves oxygen, nitrogen or fluorinebonded to hydrogen. 3. London Disperson Forces These forces are said to be weak and short lived .

24. Acids vs. Bases • Acids Bases(alkalis) • Turn blue litmus red Turn red litmus blue • taste sour taste bitter • corrode metals slippery feel • Provide H+ ions Provide OH- ions • Conducts electricity Conducts electricity

25. Naming Acids & Bases For Binary Acids: • Contain hydrogen + nonmetal • Does not contain oxygen • Naming follows the pattern of “hydro-stem-ic acid” • Ex. HCl is hydrochloric acid For ternary or oxyacids: • Contains hydrogen and polyatomic anion. • When naming, you must recognize the polyatomic ion in the formula. HClO Hypochlorous HClO2 Chlorous HClO3 Chloric HClO4Perchloric

26. Arrhenius Definition Acids release hydrogen ions in water: HCl H+ + Cl- Bases release hydroxide ions in water: NaOH Na+ + OH-

27. Bronsted-Lowry Definition • Acids are PROTON DONORS • Bases are PROTON ACCEPTORS • When HCN Dissolves in water a reaction occurs: • HCN + H2O ----> H3O+ + CN- • HCN is a Bronsted Acid • Water is acting as a Bronsted Base.

28. Acids/Bases • Conjugate acid-base pairs differ by • a hydrogen ion. • HCN + H2O ----> H3O+ + CN- • Acidbaseconj. Acidconj. base

29. Molarity • Molarity (M): • - Measure of concentration • - Defined as moles of solute per Liters of solution. • - M = mol/L • Ex: What is the molarity of a solution obtained by dissolving 24.5g of H2SO4 in enough water to make 1.50L of solution?

30. Neutralization • An Acid/ base reaction is called a neutralization reaction. • Ex: NaOH + HCl --> NaCl +HOH • base + acid --> a salt + water

31. Titration Equation: (Moles A) MBVB = (Moles B) MAVA Dilution Equation: M1V1 = M2V2

32. Formulas for calculating pH and pOH. pH = - log [H3O+] pOH = - log [OH-] pH + pOH = 14

33. Formulas for calculating concentration. [H3O+] = 10-pH [OH-] = 10 –pOH [H3O+][OH-] = 10-14

34. pH Practice What is the pH of a solution having [OH-] = 1 x 10-3? pOH = 3 so 14-3=11, pH = 11 2. What is the concentration of a solution with a pH = 3.5? 3.5 mol/L b. 5.0 x 10-3 mol/L c. 3.2 x 10-11 mol/L d. 3.2 x 10-4 mol/L Try without a calculator D is the correct answer

35. Hydrolysis • Are each of the following salts acidic,basic or neutral? • Al(NO3)3 • CaSO4 Acid Neutral

36. Solution formed From salt

37. Multi-Step Reactions • Depletion of ozone has multiple steps. • 1. NO + O3 ---> NO2 + O2 • 2. NO2 + O ---> NO + O2 • Steps are known as reaction mechanism. • In a multi step reaction the slowest step is the rate determining step.

38. Equilibrium • A reversible chemical reaction is one that can occur in either direction: • A + B AB • Reactions that do not go to completion reach a point called dynamic equilibrium. • Dynamic Equilibrium : • Rate of forward reaction = Rate of Reverse reaction. • The concentration of reactants & products is constant.

39. Le Chatlier’s Principle • If a “stress” is applied to a rxn at equilibrium, the • equilibrium shifts in the direction to relieve the stress. • Stressors that cause equilibrium to shift: • 1. Concentration • 2. Pressure • 3. Volume • 4. Temperature Trick: “Same side opposite; Opposite side same.”

40. Temperature Temperature will effect both forward and reverse reactions. 2A + B2 <---> 2AB + Heat (Exothermic) Heat + C2 + 2D <---> 2CD (Endothermic)

41. Equilibrium Constant Keq Keq = [products] / [reactants] For gaseous and aqueous solutions If Keq > 1 forward rxn favored, goes to right If Keq < 1 reverse rxn favored, goes to left If Keq = 1 both are equally favored

42. Ionization constant of an acid (Ka) is the constant that describes how much an acid dissociates (ionizes) in water. • H2C2O4(aq) <---> H+(aq) + HC2O4-(aq) • Ka = [H+][HC2O4-] • [H2C2O4] % Ionization = [H+]x 100 [acid]

43. Isomers of Hexane hexane 2-methyl pentane 3-methyl pentane 2,2-dimethyl butane 2,3-dimethyl butane

44. Alkanes, Alkenes, and Alkynes Alkanes CnH2n+2 single bonds Saturated Alkenes CnH2n double bonds Unsaturated Alkynes CnH2n-2 triple bonds Unsaturated

45. Aromatics • Aromatic Hydrocarbons are based on the compound benzene, C6H6. • Benzene’s structure: a six sided ring with alternating double bonds Benzene can be represented as: or

46. Naming Aromatics • 1. If the ring is the parent chain with other things attached it is called benzene and all other rules apply CH2CH2CH2CH2CH2CH3 Hexyl benzene

47. The Phenyl Group • The benzene ring can be used as a branch in an alkane. In this case it is called the phenyl group. (you will know it is “phenyl if it is attached to the middle of a chain) CH3CHCH2CH2CH2CH3 This compound is 2-phenylhexane:

48. Ester Reaction Alcohol + Acid  Ester + Water

49. Good Luck!