2. AP Chemistry The Ultimate Chemical Equations Handbook

3. HOMEWORK  Do all exercises in this book on a separate sheet of paper.  DO NOT WRITE IN THE BOOK.

4. Chapter 1  Symbols and Nomenclature of the elements  There is interesting info where the elements got their name, but nothing we will cover.

5. Chapter 2 and 3 Naming Binary Compounds  First, determine if you have an ionic compound or a covalent compound.  A metal and a nonmetal will form an ionic bond.  Compounds with Polyatomic ions form ionic bonds.  Nonmetals bonding together or Nonmetals and a metalloid form covalent bonds.

6. Covalent bonding is very similar to ionic naming  You always name the one that is least electronegative first (furthest from fluorine)  Most electronegative last, and gets the suffix “-ide”.

7. Covalent bonding is very different from ionic naming  Ionic names ignored the subscript because there was only one possible ratio of elements.  Covalent gives several possibilities so we have to indicate how many of each atom is present in the name

8. Prefixes you have to know prefixmeaningprefixmeaning *mono-1hex-6 di-2hept-7 tri-3oct-8 tetr-4non-9 pent-5dec-10 * the first atom named does not get the prefix “mono-”, it just keeps its original name!

9. Examples  CO  carbon monoxide  CO 2  carbon dioxide  NI 3  nitrogen triiodide P4O6P4O6P4O6P4O6  tetraphosphorus hexoxide

10. Continuing I4O9I4O9I4O9I4O9  tetriodine nonoxide  S 2 F 10  disulfur decafluoride  IF 7  Iodine heptafluoride  Si 2 Cl 6  disilicon hexachloride

11. Naming ionic compounds  For monoatomic anions only  drop the ending and add “-ide”  so F -  fluoride  Cl -, O 2-, C 4-  chloride, oxide and carbide

12. Continuing…  cations keep the name of the element.  When naming compounds always name the positive (cation) first and the negative (anion) last.  so mixing ions of chlorine and sodium give you  sodium chloride  (positive) (negative)

13. Determining the formula of ions  Ionic compounds are neutral  You need to find the lowest number of each ion to make it neutral  for example:  Na + and O 2-  2 sodium for every one oxygen  Na 2 O

14. More examples  Al 3+ and O 2-  Al 2 O 3  K + and Cl -  KCl  the subscripts don’t effect the name if there is only one possibility  still (cation)(anion)  Aluminum oxide  Potassium chloride

15. Several atoms can form a couple of different ions.  All of these are metals that are not in group 1, 2 or aluminum.  for example iron can form Fe 2+ or Fe 3+  These are said as iron (II) and iron (III)  Cu + and Cu 2+ is Copper (I) and Copper (II)

16. Figuring out charge on these elements  If the ion is named, the charge is in the name.  If you have the formula, use the charges of the other ions present to determine the charge.  Remember  Alkali will always be +1  Alkaline Earth +2, Halogens -1, oxygen group -2  Aluminum will always be +3

17. Examples  Copper (II) chloride  Cobalt (III) sulfide  NiF 2  TiS 2

18. Polyatomic ions  Polyatomic Ions- many atoms in one ion  You can NOT break these apart in this section.  the “ide” suffix only applies to monoatomic anions

19. Common polyatomic ions AmmoniumNH 4 + Chlorate ClO 3 - AcetateCH 3 CO 2 - PerchlorateClO 4 - ChromateCrO 4 2- NitrateNO 3 - Permanganate MnO 4 - Dichromate Cr 2 O 7 2- NitriteNO 2 - Carbonate CO 3 2- Hydroxide OH - Sulfate SO 4 2- Phosphate PO 4 3- HypochloriteClO - SulfiteSO 3 2- ChloriteClO 2 - Oxalate CyanideCN - Thiocyanate SCN - Hydrogen carbonate HCO 3 - Hydrogen sulfate HSO 4 - Hydrogen sulfite HSO 3 - C 2 O 4 2- Thiosulfate S 2 O 3 2- IodateIO 3 - BrO 3 - Bromate SilicateSiO 3 2- AsO 4 3- Arsenate

20. YOU WILL HAVE TO MEMORIZE THESE!  This is one of the big differences from last year.  We will have a quiz just like the elements quiz last year over these!

21. Determining the formula of ions  Ionic compounds are neutral  Remember– don’t break a polyatomic ion apart  for example: Ammonium carbonate  NH 4 + and CO 3 2-  (NH 4 ) 2 CO 3

22. Chapter 4 acids and salts  Oxyanions- negative ions containing oxygen.  These have the suffix “-ate” or “-ite”  “-ate” means it has more oxygen atoms bonded, “-ite” has less  For example  SO 4 2- sulfate  SO 3 2- sulfite

23. Oxyanions  Oxyanions may contain the prefix “hypo-”, less than, or “per-”, more than.  For example  ClO 4 - Perchlorate  ClO 3 - Chlorate  ClO 2 - Chlorite  ClO - Hypochlorite

24. Acids  Certain compounds produce H + ions in water, these are called acids.  You can recognize them because the neutral compound starts with “H”.  For example HCl, H 2 SO 4, and HNO 3.  Don’t confuse a polyatomic ion with a neutral compound.  HCO 3 - is hydrogen carbonate (or bicarbonate), not an acid.

25. Naming acids  Does it contain oxygen?  If it does not, it gets the prefix “hydro-” and the suffix “-ic acid”  HCl  Hydrochloric acid  HF  Hydroflouric acid  HCN  Hydrocyanic acid

26. Naming Acids  If it does contain an oxyanion, then replace the ending.  If the ending was “–ate”, add “-ic acid”  If the ending was “–ite”, add “-ous acid”  H 2 SO 4 Sulfuric Acid  H 2 SO 3 Sulfurous Acid

27. Examples  HNO 3  HI  H 3 AsO 3

28. Chapter 5 Complex ions  Complex ion- transition metal ion with attached ligands  Fe(CN) 6 3-  Glance over this chapter. Skip the problems, this material is out of the test.

29. Ch 6 Organic  Alkanes- straight chain hydrocarbons with all single bonds  Alkenes- hydrocarbons with a double bond  Alkynes- hydrocarbons with a triple bond  Cyclic hydrocarbons- rings

30. Root words # of C atoms Meth1Hex6 Eth2Hept7 Prop3Oct8 But4Non9 Pent5Dec10

31. Name this molecule 4 ethyl octane And give its molecular formula C 10 H 22 4 propyl decane C 13 H 28

32. Name and give the formula Methyl cyclohexane C 7 H 14

33. Functional groups  halogenated Alcohols R-OH *R means any carbon chain -ol Carboxylic Acids R-C=O -OH -oic acid Aldehydes R=OR=O -al Ketones R-C-R =O -one at the edge NOT at the edge

36. Predicting organic reactions  Addition reactions occur by adding halogens or hydrogen to alkene or alkynes.  In the reaction, the new molecule takes the place of the double or triple bond.  Cl 2 + CH 3 -CH=CH 2  CH 3 -CClH- CClH 2

37. example  1- butene is reacted with fluorine  C 4 H 8 + F 2  C 4 H 8 F 2

38. Predicting organic reactions  Substitution reactions occur by adding halogens to an alkane.  In the reaction, the new molecule takes the place of a hydrogen.  Cl 2 + CH 3 -CH 3  CH 3 -CClH 2 + HCl  Cl 2 + C 2 H 6  C 2 ClH 5 + HCl

39. Predicting organic reactions  Combustion reactions occur when an organic compound is burned in oxygen.  The products of a complete combustion are water vapor and carbon dioxide.  C 6 H 12 O 6 + 6 O 2  6 H 2 O+ 6 CO 2

40. Predicting organic reactions  Esterification reactions  Made by reacting carboxylic acids with alcohols. R-C-O-H O= H-O-R + R-C-O-R O= + H-O-H Carboxylic acid alcohol Ester

41. Examples  Fluorine is added to 2 propene  Ethanol is burned in oxygen  Chlorine is added to propane  Ethanoic acid is reacted with 1- butanol

42. Ch 7 Predicting molecular equations  Synthesis Reactions pg 42  Metals and nonmetals combine to form salts. CHECK CHARGES OF IONS!!!!!  Metal Oxides and water form bases, bases have OH attached  Nonmetal oxides and water form acids  Metal oxides and nonmetal oxides form salts (you will make an oxyanion- just move one oxygen over)

43. Examples  Magnesium burns in oxygen  Calcium reacts with chlorine gas  Magnesiumvoxide reacts with water  Carbon dioxide is bubbled through water  Lithium oxide is added to carbon dioxide

44. Ch 7 Predicting molecular equations  Decomposition Reactions pg 42  CHECK CHARGES OF IONS!!!!!  Carbonates decompose into oxides and CO 2  Chlorates decompose into chlorides and O 2  Ammonium and a base makes ammonia and water  Hydrogen peroxide decomposes into water and oxygen.  A binary compound may break down to produce two elements.

45. Three that always appear on the AP test  Hydrogen peroxide decomposes into water and oxygen  Ammonium hydroxide decomposes into ammonia and water  Carbonic acid decomposes into carbon dioxide and water

46. Examples  Titanium (IV) chlorate decomposes  Copper (III) carbonate is heated  Magnesium chloride is electrolyzed  Carbonic acid is heated  Hydrogen peroxide decomposes

47. Ch 8 Single replacement reactions  A + BX  AX + B  You will have a chart of activity series  More active metals will replace less active metals from their compound in a solution  A less active element will have no reaction when added to a more active element!  Active metals replace hydrogen in water  Active metals replace hydrogen in acids  Active nonmetals replace less active nonmetals from their compounds in solutions

48. Activity series chart

49. Old chart from test

50. Examples  Zinc is added to a solution of cobalt (II) chloride  Cadmium is added to a solution of barium iodide  Lithium is added to a solution of copper (II) chlorate  Chromium is left in water

51. Examples  Potassium is added to sulfuric acid  Silver is added to hydrochloric acid  Chlorine gas is bubbled through a solution of sodium bromide

52. Chapter 9 Double replacement reactions  AY + BX  AX + BY  These reactions occur in solution  Remember in solution the ions are free floating. For a reaction to occur, the ions have to come together and leave their dissolved state.  Formation of a precipitate  Formation of a gas  Formation of a molecular species

53. Solubility Rules  Acids are soluble.  Compounds of: alkali metals, ammonium, and nitrate are soluble.  This needs to be memorized.

54. Other soluble compounds   All acetates are soluble except Fe 3+   All chlorates are soluble.   All binary compounds of the halogens (other than F) with metals are soluble, except those of Silver, Mercury(I), and Lead. Fluorides are insoluble except for rule 1 and 2.   All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury (I).

55. Insoluble compounds   Carbonates, oxalates, sulfites, chromates, oxides, silicates, and phosphates are insoluble.   Hydroxides are insoluble except Ba, Sr, and Ca   Sulfides are insoluble except for calcium, barium, strontium, magnesium.   The exception is with alkali metals or ammonium.

56. Net Ionic Equation  Determine what dissolved and precipitated  Mg(NO 3 ) 2(aq) +Na 2 CO 3(aq)  MgCO 3(s) + 2NaNO 3(aq)  Dissociate everything that is aqueous, not solid  Mg ++ + 2NO 3 - + 2Na + + CO 3 -- – MgCO 3(s) + 2 Na + + 2NO 3 -  Now cancel out everything that is the same on both sides of the equation  These are called spectator ions The remaining part is the net ionic equation Mg ++ + CO 3 --  MgCO 3(s)

57. Examples  Hydrochloric acid reacts with silver nitrate  Potassium carbonate reacts with calcium chlorate  Sodium chloride reacts with ammonium oxalate  Scandium acetate reacts with lithium chromate

58. Gases  H 2 S (hydrogen sulfide) is formed from any sulfide reacting with an acid  CO 2 (Carbon dioxide) is formed from any carbonate reacting with an acid, water is also produced  SO 2 (sulfur dioxide) is formed from any sulfite reacting with an acid, water is also produced  NH 3 (ammonia) is formed from ammonium reacting with a soluble hydroxide

59. Examples  Ammonium chloride reacts with calcium hydroxide  Sodium sulfide is combined with nitric acid  Ammonium carbonate is combined with barium chlorate  Lithium sulfite reacts with phosphoric acid

60. Formation of a molecular species  It is the same as precipitates or gases except a liquid is formed.  Acid base neutralization reactions will produce water.  NaOH + HNO 3  H 2 O (l) + NaNO 3 (aq)

61. Strong acids AcidformulaAcidFormula Hydrochloric acid HClSulfuric Acid H 2 SO 4 Hydrobromic acid HBrNitric AcidHNO 3 Hydriodic acid HIPerchloric Acid HClO 4 Chloric Acid HClO 3

62. Strong Bases NameFormulaNameFormula Sodium Hydroxide NaOHCalcium Hydroxide Ca(OH) 2 Potassium Hydroxide KOHStrontium Hydroxide Sr(OH) 2 Barium Hydroxide Ba(OH) 2 these make a lightning bolt on the periodic table!

63. Strong acids and bases  Strong acids and bases are not at equilibrium, there is no reverse reaction.  Strong acids and bases will never be formed in a net ionic equation.  All other acids/bases can be formed, and will be formed by reacting the appropriate ion with a strong acid/base.  *Most other bases are insoluble

64. Examples  Calcium hydroxide reacts with chloric acid  Hydrochloric acid reacts with calcium nitrite  Nitric acid reacts with sodium chlorite  Sodium chloride is mixed with sulfuric acid

65. Chapter 10 Aqueous Solutions and Ionic Equations  This chapter has already been covered  *Only dissociate soluble ionic compounds  Molecular equation  Na 2 S + CrCl 2  CrS + NaCl  Full Ionic Equation  2Na + +S 2- +Cr 2+ +2Cl -  CrS +Na + +Cl -  Net Ionic Equation  Cr 2+ +S 2-  CrS

66. Chapter 11 Redox Reactions  Redox or oxidation-reduction reactions are reactions that involve a transfer of electrons.  Oxidation is the loss of electrons.  Reduction is the gain of electrons.  (think of the charge, OIL RIG)  4 K + O 2 → 4 K + + 2 O 2-  Potassium get oxidized, oxygen get reduced

67. Using oxidation states  In the reaction…  2 Na +2 H 2 O →2 NaOH + H 2  0 +1 -2 +1 -2 +1 0  Note the changes  Sodium went from 0 to 1  2 of the hydrogen atoms went from +1 to 0 (the other two were unchanged)

68. Breaking into two half reactions  Sodium must have lost 2 electrons  2 Na → 2Na + + 2 e -  And Hydrogen gained two electrons  2 H 2 O +2 e - → 2 OH - + H 2  Sodium is oxidized, hydrogen is reduced in this reaction  Oxidation is an increase in oxidation state  Reduction is a decrease in oxidation state

69. Balancing Redox Equations  by Half Reactions Method or oxidation state method  The book does not separate these into half reactions, although it adds another step I think it makes it easier

70. Half reactions  Ce 4+ + Sn 2+ → Ce 3+ + Sn 4+  Half reactions  Ce 4+ + e - → Ce 3+  Sn 2+ → 2e - + Sn 4+  Electrons lost must equal electrons gained!  2 Ce 4+ +2 e - → 2 Ce 3+  Merge the two half reactions  2 Ce 4+ + Sn 2+ → 2 Ce 3+ + Sn 4+

71. Redox reactions in acidic solutions  It will be noted in the problem  Balance all elements except hydrogen and oxygen.  Balance oxygen by adding H 2 O (which is always prevalent in an acidic solution)  Balance hydrogen by adding H +  Then balance the charge adding electrons and proceed as normal.

72. Example  In an acidic solution  Cr 2 O 7 2- + Cl - → Cr 3+ + Cl 2  Half reactions  Cr 2 O 7 2- → Cr 3+  Cl - → Cl 2

73. Reduction side  Cr 2 O 7 2- → Cr 3+  Cr 2 O 7 2- → 2 Cr 3+  Cr 2 O 7 2- → 2 Cr 3+ + 7 H 2 O  Cr 2 O 7 2- + 14 H + → 2 Cr 3+ + 7 H 2 O  Cr 2 O 7 2- + 14 H + + 6 e - → 2Cr 3+ +7 H 2 O

74. Oxidation side  Cl - → Cl 2  2 Cl - → Cl 2  2 Cl - → Cl 2 + 2 e -  I have to equal 6 e - so multiply by 3  6 Cl - → 3 Cl 2 + 6 e -

75. Combine my half reactions  Cr 2 O 7 2- + 14 H + + 6 e - → 2 Cr 3+ + 7 H 2 O  6 Cl - → 3 Cl 2 + 6 e -  And you get  Cr 2 O 7 2- +14 H + +6Cl - → 2Cr 3+ +3 Cl 2 +7H 2 O  The electrons cancel out.

76. Example  In an acidic solution  MnO 4 - + H 2 O 2 → Mn 2+ + O 2

77. Balancing Redox Equations in a basic solution  Look for the words basic or alkaline  Follow all rules for an acidic solution.  After you have completed the acidic reaction add OH - to each side to neutralize any H +.  Combine OH - and H + to make H 2 O.  Cancel out any extra waters from both sides of the equation.

78. Example  We will use the same equation as before  In a basic solution  MnO 4 - + H 2 O 2 → Mn 2+ + O 2  2 MnO 4 - + 6 H + + 5 H 2 O 2 → 2 Mn 2+ + 5 O 2 + 8 H 2 O

79. Another example  In a basic solution  MnO 4 − + SO 3 2- → MnO 4 2− + SO 4 2-  Half reactions  MnO 4 − → MnO 4 2−  SO 3 2- → SO 4 2-

80. Predicting redox reactions  Single replacement reactions we went over a few chapters ago are all redox reactions.  Here are the same examples, this time write out the net ionic equation  The change will be the charge!!

81. Examples  Zinc is added to a solution of cobalt (II) chloride  Cadmium is added to a solution of barium iodide  Lithium is added to a solution of copper (II) chlorate  Chromium is left in water

82. Examples  Potassium is added to sulfuric acid  Silver is added to hydrochloric acid  Chlorine gas is bubbled through a solution of sodium bromide

83. Common Oxidizing/reducing agents Oxidizing agentProductReducing AgentProduct MnO 4 - in acidMn 2+ H2O2H2O2 O2O2 MnO 2 in acidMn 2+ Halogens (dilute basic) Hypohalite ion (hypochlorite) MnO 4 - in baseMnO 2 CrO 4 2- in acidCr 3+ halogens (conc basic) Halate ion (chlorate) Cr 2 O 7 2- in acidCr 3+ HClO 4 Cl - Free metalsMetal ions Na 2 O 2 NaOH H2O2H2O2 H2OH2OC 2 O 4 2- CO 2 H 2 SO 4 conc.SO 2 Sulfite or SO 2 SO 4 2- Free halogensHalide ion Free halogen HNO 3 conc.NO 2 NO 2 - NO 3 - HNO 3 diluteNO Metal ionsLower oxidation number Metal ionHigher oxidation number

84. IMPORTANT  For every redox equation, you have to have both an oxidizing agent and a reducing agent!!  Otherwise it doesn’t work.

85. Examples  Bromate reacts with bromide in an acidic solution  Permanganate reacts with oxalate in an acidic solution  Calcium metal reacts with permanganate in a sodium hydroxide solution

86. Example  Chromate reacts with chloride in an acidic solution  Chlorine reacts with permanganate in a concentrated solution of sodium hydroxide

87. Ch 12 Electrolysis in water  Electrolysis is a fairly simple process.  There are two plates in a solution, and an electric current is sent through.  The plates are the cathode, where reduction takes place, and the anode, where oxidation takes place.  cathode-reduction anode- oxidation

88. Rules for cathode reaction  A cation may be reduced to a metal  Cu + + 1 e -  Cu  Or water way be reduce to hydrogen  2 H 2 O + 2 e -  H 2 + 2 OH -  Transition metals tend to reduce before water, main group metals tend to reduce after

89. Rules for anode reactions  An anion nonmetal may be oxidized to a nonmetal  2 Cl -  Cl 2 + 2 e -  Water may be oxidized to oxygen  2 H 2 O  O 2 + 4 H + + 4 e -  Chlorine, bromine and iodine will oxidize before oxygen. That is it.

90. Rules for molten binary salts  Molten means melted, with no water.  These are straightforward and easy!  Molten magnesium chloride is electrolyzed  MgCl 2  Mg + Cl 2

91. examples  Aqueous calcium bromide is electrolyzed  Aqueous chromium (III) nitrate is electrolyzed  Aqueous cobalt (II) bromide electrolyzed  Molten sodium chloride is electrolyzed

92. Ch 13 Complex ion reactions  Formation of complex ions  Common complex ions metals  Fe Co Ni Cr Cu Zn Ag Al  Common ligands  NH 3 CN - OH - SCN -  General rule: the number of ligands will be twice the charge of the metal ion

93. Example  Iron (III) chloride reacts with potassium cyanide  Fe 3+ + CN -  Fe(CN) 6 3-  How did I get the charge? Iron is 3+, 6 cyanides at 1-

94. Examples  Zinc (I) fluoride reacts with sodium thiocyanate  Concentrated ammonia is reacted with cobalt (III) iodide  Barium hydroxide reacts with nickel (II) nitrate

95. Solution Stoichiometry Ch 4

96. Homework  Due with test  Pg 171 Chapter 4 review  1-87 odd

97. Water and the Nature of Aqueous Solutions   Water is the foundation of all life on Earth.  oxygen  Each O-H covalent bond is highly polar because the electronegativity of oxygen is so much greater than hydrogen.   The dipole moment creates a slightly negative O and a slightly positive H as the electrons are pulled strongly toward O.

98. Water is a bent molecule with a bond angle of 109.5 o 109.5 o

99. Hydrogen bond   Polar molecules are attracted to one another by dipole attractions.   In water, it is hydrogen bonding.   The molecules can’t easily slide past one another, they are held in place.   The strength of the hydrogen bonds accounts for water’s high surface tension, its low vapor pressure, its high specific heat, its high heat of vaporization, and its high boiling point.

100. Hydrogen bonding in water

101. Solvents and Solutes   Chemically pure water never exists in nature because water dissolves so many substances   * An aqueous solution is one in which water samples contain dissolved substances.   Components of a solution.   A solvent is the dissolving medium. A solute is the dissolved particles.

102. Solutions   Solutions are homogeneous mixtures in which solute particles are usually less than 1 nm in diameter. The solute and solvent are not capable of being separated by filtration.   Substances that dissolve most readily in water include ionic compounds and polar covalent molecules.   Nonpolar molecules do not dissolve in water (a polar molecule).   Rule: Like dissolves like

103. The Solution Process   Water molecules are in constant motion as a result of their kinetic energy.  T  The molecules collide with solute particles.   The solvent molecules attract the solute ions.   The ionic crystal breaks apart by the action of the solvent.   Solvation is the process that occurs when a solute dissolves.  http://mw2.concord.org/public/student/solution/dissolve.cml http://mw2.concord.org/public/student/solution/dissolve.cml

104. Insoluble   In some ionic compounds, the attractions between the ions in the crystal are stronger than the attractions exerted by water.   * These compounds are insoluble.   Nonpolar substances form a solution because there are no repulsive forces between them, not because the solute and solvent are attracted.

105. Electrolytes and Nonelectrolytes   Compounds that conduct an electric current in aqueous solution or the molten state are called electrolytes.   All ionic compounds are electrolytes.   Soluble ionic compounds conduct electricity in both a solution and in the molten state.   Insoluble ionic compounds only conduct electricity in the molten state.

106. Nonelectrolytes   Compounds that do not conduct an electric current in either aqueous solution or the molten state are called nonelectrolytes.   Many molecular compounds are nonelectrolytes because they do not contain ions.   Most compounds of carbon are nonelectrolytes.   Some very polar molecular compounds are nonelectrolytes in the pure state, but are electrolytes when they dissolve in water.   This occurs because such compounds ionize in solution.

107. Weak Electrolytes   Not all electrolytes conduct electricity to the same degree.   A weak electrolyte conducts electricity poorly because only a fraction of the solute exists as ions.   * Most of the compound is in the original form.   * The most common weak electrolytes are weak acids and weak bases.

108. Strong Electrolytes   A strong electrolyte conducts electricity very well because almost all of the solute exists as separated ions.   * Very little of the original compound remains intact.   * Classes of electrolytes include soluble salts, strong acids and strong bases.

109. The Composition of Solutions   To perform stoichiometric calculations when two solutions are mixed, two things must be known.   The nature of the reaction, which depends on the exact forms the chemical takes when dissolved.   The amounts of chemical present in solution, usually expressed as concentrations.

110. Molarity   Molarity (M) is defined as the moles of solute per volume of solution in liters.   M = molarity = moles of solute   liters of solution     Calculate the molarity of a solution prepared by dissolving 11.5 g of solid NaOH in enough water to make 1.50 L of solution.

111. Calculation of Molarity II   Calculate the molarity of a solution prepared by dissolving 1.56 g of gaseous HCl in enough water to make 26.8 mL of solution.

112. Concentrations of Ions   Give the concentration of each type of ion in the following solutions:   a. 0.50 M Co(NO 3 ) 2     b. 1 M Fe(ClO 4 ) 3

113. Concentrations of Ions II   Calculate the number of moles of Cl - ions in 1.75 L of 1.0 x 10 -3 M ZnCl 2.     Concentration and Volume I.   Typical blood serum is about 0.14 M NaCl. What volume of blood contains 1.0 mg NaCl?

114. Solutions of Known Concentration  .To analyze the alcohol content of a certain wine, a chemist needs 1.00 L of an aqueous 0.200 M K 2 Cr 2 O 7 (potassium dichromate) solution. How much solid K 2 Cr 2 O 7 must be weighed out to make this solution?

115. Dilution  money .To save money and space in a laboratory, solutions are often in concentrated form. Water is added to achieve the molarity desired for a particular solution.   This process is called dilution.   Dilution with water does not alter the numbers of moles of solute present.

116. Concentration and Volume   What volume of 16M sulfuric acid must be used to prepare 1.5 L of a 0.10 M H 2 SO 4 solution?     Another way to express the dilution process is:M 1 V 1 = M 2 V 2  What volume of 14.8M ammonia must be used to prepare.75 L of a 1.5 M H 2 SO 4 solution?

117. Stoichiometry of Precipitation Reactions   1. Identify the species present in the combined solution, and determine what reaction occurs.   2. Write the balanced net ionic equation for the reaction.   3. Calculate the moles of reactants.   4. Determine which reactant is limiting.   5. Calculate the moles of reactants or products, as required.   6. Convert to grams or other units, as required.

118. Determining the Mass of Products Formed   Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO 3 solution to precipitate all the Ag + ions in the form of AgCl.

119. Stoichiometry of Precipitation Reactions   1. Identify the species present in the combined solution, and determine what reaction occurs.   2. Write the balanced net ionic equation for the reaction.   3. Calculate the moles of reactants.   4. Determine which reactant is limiting.   5. Calculate the moles of reactants or products, as required.   6. Convert to grams or other units, as required.

120. Determining the Mass of Products Formed   Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO 3 solution to precipitate all the Ag + ions in the form of AgCl.

121. Acid-Base Reactions   The Bronsted-Lowry Definitions:   An acid is a proton donor.   A base is a proton acceptor.   Arrhenius definitions:   Acid generate protons (hydronium)  Bases generate hydroxide  Lewis definitions  Acids are electron pair acceptors  Bases are electron pair donors

122. Performing Calculations for Acid- Base Reactions   1.List the species present in the combined solution before any reaction occurs, and decide what reaction will occur.   2. Write the balance net ionic equation for this reaction.   3. Calculate the moles of reactants. For reactions in solution, use the volumes of the original solutions and their molarities.   4. Determine the limiting reactant.   5. Calculate the moles of the required reactant and product.   6. Convert to grams or volume (of solution), as required.

123. Example Problem   What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL of 0.350 M NaOH?   In a certain experiment, 28.0 mL of 0.250 M HNO 3 and 53.0 mL of 0.320 M KOH are mixed. Calculate the amount of water formed in the resulting reaction. What is the concentration of H + or OH - ions in excess after the reaction goes to completion?

124. Acid-Base Titrations   Volumetric analysis is a technique for determining the amount of a certain substance by doing a titration.   A titration is a process in which a solution of known concentration is used to determine the concentration of another solution through a monitored reaction.   The equivalence point is reached when all of the moles of H + ions present in the original volume of acid solution have reacted with an equivalent number of moles of OH - ions added. This is also known as the stoichiometric point.

125. Titration   An acid-base indicator is a chemical that changes color once the equivalence point is reached.   The endpoint of a titration is the point in a titration where the indicator actually changes color.   Three requirements for a titration   1. The exact reaction between titrant (known solution) and analyte (unknown solution or one you are analyzing) must be known (and rapid).

126. Three requirement (cont.)  2. The stoichiometric (equivalence) point must be marked accurately.  3. The volume of titrant required to reach the stoichiometric point must be known accurately.

127. Problem solving   The first step in the analysis of a complex solution is to write down the components and focus on the chemistry of each one.  Take a big problem and look at the small problems within.

128. Neutralization Titration A student carries out an experiment to standardize (determine the exact concentration of) a sodium hydroxide solution. To do this, the student weighs out a 1.3009-g sample of potassium hydrogen phthalate (KHC 8 H 4 O 4, often abbreviated KHP. KHP (molar mass 204.22 g/mol) has one acidic hydrogen. The student dissolves the KHP in distilled water, adds phenolphthalein as an indicator, and titrates the resulting solution with the sodium hydroxide solution to the phenolphthalein endpoint. The difference between the final and initial buret readings indicates that 41.20 mL of the sodium hydroxide solution is required to react exactly with the 1.3009 g KHP. Calculate the concentration of the sodium hydroxide solution.

129. Neutralization Analysis   An environmental chemist analyzed the effluent (the released waste material) from an industrial process known to produce the compounds carbon tetrachloride (CCl 4 ) and benzoic acid (HC 7 H 5 O 2 ), a weak acid that has one acidic hydrogen atom per molecule. A sample of this effluent weighing 0.3518 g was shaken with water, and the resulting aqueous solution required 10.59 mL of 0.1546 M NaOH for neutralization. Calculate the mass percent of HC 7 H 5 O 2 in the original sample.