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  • Slide 1
  • AP CHEMISTRY CHAPTER 4 SOLUTION STOICHIOMETRY
  • Slide 2
  • Electrons arent shared evenly (oxygen is more electronegative) Electrons spend more time close to O than to H.
  • Slide 3
  • This uneven distribution of charge makes water polar. Because of this, water is a good solvent. The positive end (H) attracts negative ions or the negative end of another polar molecule. The negative end of water (O) attracts the positive ions or the positive end of another polar molecule.
  • Slide 4
  • When water surrounds an ionic crystal, the H end attracts the anion and the O end attracts the cation. This process is called hydration.
  • Slide 5
  • Hydration causes salts (ionic compounds) to dissolve. H 2 O also dissolves polar covalent substances such as C 2 H 5 OH. H 2 O doesnt dissolve nonpolar covalent substances because there is not enough attraction between the water and the nonpolar molecule.
  • Slide 6
  • Hydration
  • Slide 7
  • Show the association of the ions with some water molecules when 1 formula unit of KCl dissolves in excess water. K+ Cl -
  • Slide 8
  • A solution is a homogeneous mixture. In a solution, a solute dissolves in the solvent. If the solute ionizes in the solution, electricity can be conducted and the solute is said to be an electrolyte. If the solute ionizes 100% or nearly 100%, it is called a strong electrolyte. Lesser ionization occurs with weak electrolytes.
  • Slide 9
  • Svante Arrhenius determined that the extent to which a solution can conduct an electrical current depends directly on the number of ions present.
  • Slide 10
  • Solubility- is usually shown as g/given volume solvent or moles/given volume solution
  • Slide 11
  • Strong electrolytes 1.soluble salts 2.strong acids completely ionize HCl(aq), HNO 3 (aq), H 2 SO 4 (aq)
  • Slide 12
  • Ex. Show how HCl dissociates when dissolved in water. HCl H + + Cl -
  • Slide 13
  • Acid (Arrhenius) a substance that produces H + ions in water solution
  • Slide 14
  • 3.strong bases- completely ionize -contain OH -bitter taste and slippery feel -NaOH, KOH
  • Slide 15
  • Weak electrolytes -only ionize slightly (weak acids and bases) HC 2 H 3 O 2 H + + C 2 H 3 O 2 99% 1%
  • Slide 16
  • Slide 17
  • Ammonia (NH 3 ) -weak base NH 3 + H 2 O NH 4 + + OH
  • Slide 18
  • Molarity (M) = moles of solute liters of solution Molarity is the most common unit of concentration used in Chemistry. We may also see mM (millimolar) = moles of solute mL of solution
  • Slide 19
  • Ex. Calculate the molarity of a solution made by dissolving 23.4g of sodium sulfate in enough water to form 125 mL of solution. 23.4 g Na 2 SO 4 1 mol Na 2 SO 4 = 0.165 mol Na 2 SO 4 142.06g Na 2 SO 4 0.165 mol = 1.32 M 0.125 L
  • Slide 20
  • Ex. How many grams of Na 2 SO 4 are required to make 350 mL of 0.50 M Na 2 SO 4 ? 0.350L 0.50 mol Na 2 SO 4 142.06g Na 2 SO 4 = 24.9g 1 L 1 mol Na 2 SO 4
  • Slide 21
  • Ex. What volume of 1.000 M KNO 3 must be diluted with water to prepare 500.0 mL of 0.250 M KNO 3 ? Dilution problem (M 1 V 1 = M 2 V 2 ) (1.000M)(V 1 ) = (0.250M)(500.0mL) V 1 = 125 mL Remember, this formula can only be used for dilution! Never use it for a chemical reaction (stoichiometry)!
  • Slide 22
  • Read procedure for using volumetric flasks and types of pipets. We will be using both in several labs this year.
  • Slide 23
  • Slide 24
  • Lets Review Equation Writing from Chemistry I Some reactions fit neatly into a certain category of reaction type, some do not.
  • Slide 25
  • DECOMPOSITION REACTIONS
  • Slide 26
  • Reaction where a compound breaks down into two or more elements or compounds. Heat, electrolysis, or a catalyst is usually necessary.
  • Slide 27
  • A compound may break down to produce two elements. Ex. Molten sodium chloride is electrolyzed. 2NaCl(l) 2Na + Cl 2
  • Slide 28
  • A compound may break down to produce an element and a compound. Ex. A solution of hydrogen peroxide is decomposed catalytically. 2H 2 O 2 2H 2 O + O 2
  • Slide 29
  • A compound may break down to produce two compounds. Ex. Solid magnesium carbonate is heated. MgCO 3 MgO + CO 2
  • Slide 30
  • Metallic carbonates break down to yield metallic oxides and carbon dioxide.
  • Slide 31
  • Metallic chlorates break down to yield metallic chlorides and oxygen.
  • Slide 32
  • Hydrogen peroxide decomposes into water and oxygen.
  • Slide 33
  • Sulfurous acid decomposes into water and sulfur dioxide.
  • Slide 34
  • Carbonic acid decomposes into water and carbon dioxide.
  • Slide 35
  • Hydrated salts decompose into the salt and water. Na 2 CO 3 H 2 O Na 2 CO 3 + H 2 O
  • Slide 36
  • ADDITION REACTIONS Also known as Synthesis, Combination, or Composition
  • Slide 37
  • Two or more elements or compounds combine to form a single product.
  • Slide 38
  • A Group IA or IIA metal may combine with a nonmetal to make a salt.
  • Slide 39
  • A piece of lithium metal is dropped into a container of nitrogen gas. 6Li + N 2 2Li 3 N
  • Slide 40
  • Two nonmetals may combine to form a molecular compound. C + O 2 CO 2
  • Slide 41
  • When an element combines with a compound, you can usually sum up all of the elements on the product side. Ex. PCl 3 + Cl 2 PCl 5 This is a trick that works because the common positive oxidation states of P are + 3 and +5.
  • Slide 42
  • Two compounds combine to form a single product. Ex. Sulfur dioxide gas is passed over solid calcium oxide. SO 2 + CaO CaSO 3
  • Slide 43
  • A metallic oxide plus carbon dioxide yields a metallic carbonate. (Carbon keeps the same oxidation state)
  • Slide 44
  • A metallic oxide plus sulfur dioxide yields a metallic sulfite. (Sulfur keeps the same oxidation state)
  • Slide 45
  • A metallic oxide plus water yields a metallic hydroxide. A nonmetallic oxide plus water yields an acid.
  • Slide 46
  • Double Replacement (metathesis)
  • Slide 47
  • Two compounds react to form two new compounds. No changes in oxidation numbers occur. All double replacement reactions must have a "driving force" that removes a pair of ions from solution.
  • Slide 48
  • Formation of a precipitate: A precipitate is an insoluble substance formed by the reaction of two aqueous substances. Two ions bond together so strongly that water can not pull them apart. You must know your solubility rules to write these net ionic equations!
  • Slide 49
  • Simple Rules for Solubility 1.Most nitrate (NO 3 ) salts are soluble. 2.Most alkali (group 1A) salts and NH 4 + are soluble. 3.Most Cl , Br , and I salts are soluble (NOT Ag +, Pb 2+, Hg 2 2+ ) 4.Most sulfate salts are soluble (NOT BaSO 4, PbSO 4, HgSO 4, CaSO 4 ) 5.Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH) 2, Ca(OH) 2 are marginally soluble) 6.Most S 2 , CO 3 2 , CrO 4 2 , PO 4 3 salts are only slightly soluble.
  • Slide 50
  • SOLUBILITY SONG To the tune of My Favorite Things from The Sound of Music My Favorite Things Nitrates and Group One and Ammonium, These are all soluble, a rule of thumb. Then you have chlorides, theyre soluble fun, All except Silver, Lead, Mercury I. Then you have sulfates, except for these three: Barium, Calcium and Lead, you see. Worry not only few left to go still. We will do fine on this test. Yes, we will! Then you have the--- Insolubles Hydroxide, Sulfide and Carbonate and Phosphate, And all of these can be dried!
  • Slide 51
  • Ex. Solutions of silver nitrate and lithium bromide are mixed. AgNO 3 (aq) + LiBr(aq) AgBr(s) + LiNO 3 (aq)
  • Slide 52
  • Formation of a gas: Gases may form directly in a double replacement reaction or can form from the decomposition of a product such as H 2 CO 3 or H 2 SO 3. H 2 CO 3 H 2 O and CO 2 H 2 SO 3 H 2 O and SO 2 NH 4 OH NH 3 and H 2 O
  • Slide 53
  • Ex. Excess hydrochloric acid solution is added to a solution of potassium sulfite. 2HCl(aq) + K 2 SO 3 (aq) H 2 SO 3 H 2 O(l) + SO 2 (g) + 2KCl(aq) Remember that sulfurous acid decomposes into water and sulfur dioxide!
  • Slide 54
  • Ex. A solution of sodium hydroxide is added to a solution of ammonium chloride. Remember that ammonium hydroxide does not exist. NaOH(aq) + NH 4 Cl(aq) NaCl(aq) + NH 4 OH NH 3 (g) + H 2 O(l)
  • Slide 55
  • Formation of a molecular substance: When a molecular substance such as water or acetic acid is formed, ions are removed from solution and the reaction "works".
  • Slide 56
  • Ex. Dilute solutions of lithium hydroxide and hydrobromic acid are mixed. LiOH(aq) + HBr(aq) LiBr(aq) +H 2 O(l)
  • Slide 57
  • Ex. Gaseous hydrofluoric acid reacts with solid silicon dioxide. 4HF(g) + SiO 2 (s) SiF 4 (g) + 2H 2 O(l) This reaction occurs when glass is etched.
  • Slide 58
  • Single Replacement
  • Slide 59
  • Reaction where one element displaces another in a compound. One element is oxidized and another is reduced. A + BC B + AC
  • Slide 60
  • Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Each halogen will displace less electronegative (heavier) halogens from their binary salts.
  • Slide 61
  • Activity Series of Nonmetals Most Active F 2 Cl 2 Br 2 Least Active I 2
  • Slide 62
  • Ex. Chlorine gas is bubbled into a solution of potassium iodide. Cl 2 (g) + 2KI(aq) 2KCl(aq) + I 2 (s)
  • Slide 63
  • Active metals replace less active metals or hydrogen from their compounds in aqueous solution. Use an activity series or a reduction potential table to determine activity. The more easily oxidized metal replaces the less easily oxidized metal.
  • Slide 64
  • Group I,II,+III Transition Metals Hydrogen Jewelry Metals
  • Slide 65
  • Ex. Magnesium turnings are added to a solution of iron(III) chloride. 3Mg(s) + 2FeCl 3 (aq) 2Fe(s)+3MgCl 2 (aq)
  • Slide 66
  • Ex. Sodium is added to water. 2Na(s) + 2H 2 O(l) 2NaOH(aq) + H 2 (g) Alkali metal demo
  • Slide 67
  • COMBUSTION REACTIONS
  • Slide 68
  • -Elements or compounds combine with oxygen to produce the oxides of each element. The oxide of H is H 2 O, oxide of S is usually SO 2, oxide of C is CO 2, etc.
  • Slide 69
  • Hydrocarbons or alcohols combine with oxygen to form carbon dioxide and water.
  • Slide 70
  • Nonmetallic hydrides combine with oxygen to form oxides and water.
  • Slide 71
  • Nonmetallic sulfides combine with oxygen to form oxides and sulfur dioxide.
  • Slide 72
  • Ex. Carbon disulfide vapor is burned in excess oxygen. CS 2 + 3O 2 CO 2 + 2SO 2
  • Slide 73
  • Ex. Ethanol (C 2 H 5 OH) is burned completely in air. C 2 H 5 OH + 3O 2 2CO 2 + 3H 2 O
  • Slide 74
  • Diethyl ether (C 2 H 5 OC 2 H 5 ) is burned in air. C 2 H 5 OC 2 H 5 + 6O 2 4CO 2 + 5H 2 O
  • Slide 75
  • Selective precipitation- process by which ions are caused to ppt one by one in sequence to separate mixtures of ions. Qualitative analysis- process of separating and identifying ions
  • Slide 76
  • Ex. Separate Ag +, Ba 2+, Fe 3+ 1.Add Cl to remove Ag + as AgCl. 2.Add SO 4 2- to remove Ba 2+ as BaSO 4. 3.Add OH or S 2- to remove Fe 3+ as Fe(OH) 3 or Fe 2 S 3.
  • Slide 77
  • Ex. Separate Pb 2+, Ba 2+, Ni 2+ 1.Add Cl to remove Pb 2+ as PbCl 2. 2.Add SO 4 2 to remove Ba 2+ as BaSO 4. 3.Add OH or S 2 to remove Ni 2+ as Ni(OH) 2 or NiS.
  • Slide 78
  • Quantitative analysis- determines how much of a component is present. Gravimetric analysis- quantitative procedure where a ppt containing the substance is formed, filtered, dried & weighed.
  • Slide 79
  • Ex. The zinc in a 1.2000g sample of foot powder was precipitated as ZnNH 4 PO 4. Strong heating of the ppt yielded 0.4089 g of Zn 2 P 2 O 7. Calculate the mass percent of zinc in the sample of the foot powder. 0.4089gZn 2 P 2 O 7 1 mol Zn 2 P 2 O 7 2 mol Zn 65.37g = 304.7 g 1 mol Zn 2 P 2 O 7 1 mol Zn 0.1754g Zn 100 = 14.62% Zn 1.200g sample
  • Slide 80
  • Ex. A mixture contains only NaCl and Fe(NO 3 ) 3. A 0.456g sample of the mixture is dissolved in water, and an excess of NaOH is added, producing a precipitate of Fe(OH) 3. The ppt is filtered, dried, & weighed. Its mass is 0.128g. Calculate: a. the mass of the iron b. the mass of Fe(NO 3 ) 3 c. the mass percent of Fe(NO 3 ) 3 in the sample 0.128g Fe(OH) 3 1 mol Fe(OH) 3 1 mol Fe 55.85g Fe= 0.0669g Fe 106.9g Fe(OH) 3 1 mol Fe(OH) 3 1 mol Fe 0.0669g Fe 1 mol Fe 1 mol Fe(NO 3 ) 3 241.9g Fe(NO 3 ) 3 = 0.290g Fe(NO 3 ) 3 55.85g Fe 1 mol Fe 1 mol Fe(NO 3 ) 3 0.290g 100 = 63.6% Fe(NO 3 ) 3 0.456g
  • Slide 81
  • Acid-Base Reactions
  • Slide 82
  • Bronsted-Lowry acid-base definitions: acid- proton donor base- proton acceptor
  • Slide 83
  • When a strong acid reacts with a strong base the net ionic rxn is: H + (aq) + OH (aq) H 2 O(l) When a strong acid reacts with a weak base or a weak acid reacts with a strong base, the reaction is complete (the weak substance ionizes completely.) HC 2 H 3 O 2 (aq) + OH (aq) H 2 O(l) + C 2 H 3 O 2 (aq)
  • Slide 84
  • neutralization reaction - acid-base rxn When just enough base is added to react exactly with the acid in a solution, the acid is said to be neutralized.
  • Slide 85
  • Volumetric Analysis titration- process in which a solution of known concentration (standard solution) is added to analyze another solution (analyte).
  • Slide 86
  • Titrations are most often used for acids and bases, but can be used for other types of reactions, also.
  • Slide 87
  • titrant- solution of known concentration (usually in buret) equivalence point or stoichiometric point- point where just enough titrant has been added to react with the substance being analyzed
  • Slide 88
  • Indicator - chemical which changes color at or near the equivalence point End point- point at which the indicator changes color
  • Slide 89
  • Ex. 54.6 mL of 0.100 M HClO 4 solution is required to neutralize 25.0 mL of an NaOH solution of unknown molarity. What is the concentration of the NaOH solution? HClO 4 + NaOH H 2 O + NaClO 4 0.0546 L HClO 4 0.100 mol HClO 4 1 mol NaOH = 1 L HClO 4 1 mol HClO 4 0.00546 mol NaOH 0.00546 mol NaOH = 0.218 M NaOH 0.025L
  • Slide 90
  • Writing net-ionic equations 1. molecular equation -overall reaction stoichiometry 2. complete ionic equation -all strong electrolytes are represented as ions 3. net ionic equation -spectator ions are not included
  • Slide 91
  • 1. NaCl(aq) + AgNO 3 (aq) NaNO 3 (aq) + AgCl(s) 2. Na + (aq) + Cl (aq) + Ag + (aq) + NO 3 (aq) Na + (aq) + NO 3 (aq) + AgCl(s) 3. Cl (aq) + Ag + (aq) AgCl(s)
  • Slide 92
  • Oxidation-Reduction Reactions Redox Rxns - reactions in which one or more electrons are transferred.
  • Slide 93
  • Electronegativity - attraction for shared electrons most electronegative F>O>N=Cl elements Phone Call These are most likely to have negative oxidation numbers.
  • Slide 94
  • Rules for Assigning Oxidation States 1. Oxidation state of an atom in an element = 0 2. Oxidation state of monatomic element = charge 3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1) 4. H = +1 in covalent compounds 5. Fluorine = 1 in compounds 6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion
  • Slide 95
  • Review oxidation state rules on page 167. N 2 O PBr 3 HPO 3 2- P 4 O 6 NH 2 - +1-2 +3-1 +1+3-2 +3 -2 -3 +1
  • Slide 96
  • Noninteger states are rare, but possible. Fe 3 O 4 8/3 -2 O = 4(-2) = -8 Fe = 8/3 = 2 2/3 or Fe 2+, Fe 3+, Fe 3+
  • Slide 97
  • Oxidation - loss of electrons - increase in oxidation number Reduction - gain of electrons - decrease in oxidation number
  • Slide 98
  • OIL RIG Oxidation Is Loss (of e ), Reduction Is Gain (of e )
  • Slide 99
  • LEO the lion goes GER Lose Electrons = Oxidation, Gain Electrons = Reduction
  • Slide 100
  • Oxidizing agent - electron acceptor - substance that is reduced Reducing agent - electron donor - substance that is oxidized The terms oxidizing agent and reducing agent are not tested on the AP test.
  • Slide 101
  • 2KI + F 2 2KF + I 2 +1-1 0 +1-1 0 oxidized I reduced F OA F2F2 RA KI
  • Slide 102
  • 2PbO 2 2PbO + O 2 +4 -2 +2 -2 0 oxidized: O reduced: Pb OA: PbO 2 RA: PbO 2
  • Slide 103
  • Balancing redox reactions by the half-reaction method 1.Write skeleton half-reactions. 2.Balance all elements other than O and H. 3.Balance O by adding H 2 O. 4.Balance H by adding H +.
  • Slide 104
  • 5.Balance charge by adding e - to the more positive side. 6.Make the # of e lost = # of e gained by multiplying each half-rxn by a factor. 7.Add half-reactions together. 8.Cancel out anything that is the same on both sides.
  • Slide 105
  • 9.If the reaction occurs in basic solution, add an equal number of hydroxide ions to both sides to cancel out the hydrogen ions. Make water on the side with the hydrogen ions. Cancel water if necessary. 10.Check to see that charge and mass are both balanced.
  • Slide 106
  • Practice: Sn 2+ + Cr 2 O 7 2 Sn 4+ + Cr 3+ (acidic solution) Sn 2+ Sn 4+ Cr 2 O 7 2 Cr 3+ Sn 2+ Sn 4+ Cr 2 O 7 2 2Cr 3+ Sn 2+ Sn 4+ Cr 2 O 7 2 2Cr 3+ +7H 2 O Sn 2+ Sn 4+ Cr 2 O 7 2 + 14H + 2Cr 3+ +7H 2 O Sn 2+ Sn 4+ +2e Cr 2 O 7 2 + 14H + +6e 2Cr 3+ +7H 2 O 3(Sn 2+ Sn 4+ +2e ) Cr 2 O 7 2 + 14H + +6e 2Cr 3+ +7H 2 O 3Sn 2+ + Cr 2 O 7 2 + 14H + +6e - 3Sn 4+ +6e + 2Cr 3+ +7H 2 O 3Sn 2+ + Cr 2 O 7 2- + 14H + +6e - 3Sn 4+ +6e - + 2Cr 3+ +7H 2 O 3Sn 2+ + Cr 2 O 7 2 + 14H + 3Sn 4+ + 2Cr 3+ +7H 2 O
  • Slide 107
  • MnO 4 2- + I - MnO 2 + I 2 (basic solution) MnO 4 2- MnO 2 I - I 2 MnO 4 2- MnO 2 2I - I 2 MnO 4 2- MnO 2 + 2H 2 O 2I - I 2 MnO 4 2- +4H + MnO 2 + 2H 2 O 2I - I 2 MnO 4 2- +4H + + 2e - MnO 2 + 2H 2 O 2I - I 2 +2e - MnO 4 2- +4H + + 2e - +2I - MnO 2 + 2H 2 O +I 2 +2e - MnO 4 2- +4H + +2I - MnO 2 + 2H 2 O +I 2 MnO 4 2- +4H + + 4OH - +2I - MnO 2 + 2H 2 O +I 2 + 4OH - MnO 4 2- + 4H 2 O +2I - MnO 2 + 2H 2 O +I 2 + 4OH - MnO 4 2- + 2H 2 O +2I - MnO 2 + I 2 + 4OH -
  • Slide 108
  • OXIDATION-REDUCTION TITRATIONS Most common oxidizing agents: KMnO 4 & K 2 Cr 2 O 7
  • Slide 109
  • Potassium permanganate used to disinfect ponds and fish in Egypt.
  • Slide 110
  • MnO 4 - in acidic solution: MnO 4 - + 8H + + 5e Mn 2+ + 4H 2 O Purple colorless When you titrate with MnO 4 , the solution is colorless until you use up all of the reducing agent (substance being oxidized).
  • Slide 111
  • In calculations, work redox titrations like acid-base titrations. You must have a balanced reaction to know the mole ratio.
  • Slide 112
  • Slide 113
  • Slide 114
  • Slide 115