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CONCEPT: MOLARITY Molarity (M) can serve as the connection between the interconversion of to and vice versa. For example, a 5.8 M NaCl solution really means per. ( Molarity = MolesSolute ) (LitersSolution) A typical mixture consists of a smaller amount of one substance, the, dissolved in a larger amount of another substance, the. Together they form a. Page 2
PRACTICE: MOLARITY EXAMPLE 1: 2.64 grams of an unknown compound was dissolved in water to yield 150 ml of solution. The concentration of the solution was 0.075 M. What was the molecular weight of the substance? ( M = MolesSolute ) (Liters Solution) EXAMPLE 2: A solution is prepared by dissolving 0.1408 mol calcium nitrate, Ca(NO3)2, in enough water to make 100.0 ml of stock solution. If 20.0 ml of this solution is then mix with an additional 90 ml of deionized water, calculate the concentration of the calcium nitrate solution. PRACTICE 1: What is the molarity of calcium ions of a 650 ml solution containing 42.7 g of calcium phosphate? PRACTICE 2: A solution with a final volume of 750.0 ml was prepared by dissolving 30.00 ml of benzene (C6H6, density = 0.8787 g ) in dichloromethane. Calculate the molarity of benzene in the solution. ml Page 3
CONCEPT: NORMALITY Another measurement for concentration usually encountered is normality (N), which represents the number of equivalents per liter of solution. N = equivalents of solute Liters of solution equivalent (eq) = n moles An equivalent is the mass of a compound that can either donate or accept an or. An equivalent (eq) and n are both determined by the compound being used (acid or base) and if the compound is undergoing a redox reaction. Acids For an acid the number for n is based on the number of present. EXAMPLE 1: Determine the number of equivalents for each of the acids given. a) 2.5 moles CH3COOH b) 133.4 g H3PO4 Bases For a base the number for n is based on the number of present. EXAMPLE 2: Determine the number of equivalents for the following base. a) 50.0 ml of 0.165 M Ca(OH)2 Redox Reactions For a redox reaction the number for n is based on the number of transferred. EXAMPLE 3: Based on the given redox reaction determine the value for n. MnO4 (aq) + H + (aq) Mn 2+ (aq) + H2O (l) Page 4
PRACTICE: NORMALITY EXAMPLE 1: What is the normality of a solution made by dissolving 325.1 g HNO3 in enough water to create a 750.0 ml solution? EXAMPLE 2: Determine the equivalent weight of the following compounds. a) Al(OH)3 b) H2CO3 EXAMPLE 3: What volume, in ml, of 50.0 g H2SO4 is needed to create a 0.300 N H2SO4 solution? EXAMPLE 4: If a concentrated 3.25 M H3PO4 solution possesses a density of 1.350 g/ml, what is its normality? Page 5
CONCEPT: MOLARITY & CHEMICAL REACTIONS Whenever we are provided given information in a reaction we use to find any unknown information. In aqueous reactions, this given information is typically in units of or. Entities means, or. Entities of Unknown Volume of Given Moles of Given Moles of Unknown Volume of Unknown Grams of Unknown Use this chart when given a chemical equation with the known quantity in either or of a compound or element and asked to find the unknown quantity of another compound or element. EXAMPLE: How many grams of sodium metal are needed to react with 38.74 ml of 0.275 M NaOH? 2 Na (s) + 2 H2O (l) H2 (g) + 2 NaOH (aq) Page 6
PRACTICE: MOLARITY & CHEMICAL REACTIONS PRACTICE 1: How many milliliters of 0.325 M HCl are needed to react with 16.2 g of magnesium metal? ( M = MolesSolute ) (Liters Solution) 2 HCl (aq) + Mg (s) MgCl2 + H2 (g) PRACTICE 2: What is the molarity of a hydrobromic acid solution if it takes 34.12 ml of HBr to completely neutralize 82.56 ml of 0.156 M Ca(OH)2? 2 HBr (aq) + Ca(OH)2 (aq) CaBr2 (aq) + 2 H2O (l) PRACTICE 3 (CHALLENGE): Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation: Cr2O7 2- + 6 Fe 2+ + 14 H + 2 Cr 3+ + 6 Fe 3+ + 7 H2O If it takes 30.0 ml of 0.100 M K2Cr2O7 to titrate a 25 ml Fe 2+ solution, what is the molar concentration of Fe 2+? Page 7
CONCEPT: AQUEOUS SOLUTIONS The of a compound represents the maximum amount of solute that dissolves in a solvent. SOLUBILITY RULES SOLUBLE IONIC COMPOUNDS 1. Group 1A ions (Li +, Na +, K +, etc.) and ammonium ion (NH4 + ) are soluble. 2. (Nitrates) NO3 -, (acetates) CH3COO - or C2H3O2 -, and most perchlorates (ClO4 - ) are soluble. INSOLUBLE IONIC COMPOUNDS 1. (Hydroxides) OH - and (Sulfides) S 2-, are insoluble except when with Group 1A ions (Li +, Na +, K +, etc.), ammonium ion (NH4 + ) and Ca 2+, Sr 2+, Ba 2+. 2. (Carbonates) CO3 2- and (Phosphates) PO4 3- are insoluble except when with Group 1A ions (Li +, Na +, K +, etc.), ammonium ion (NH4 + ). 3. Cl -, Br -, and I - are soluble, except when paired with Ag +, Pb 2+, Cu + and Hg2 2+. 4. (Sulfates) SO4 2- are soluble, except those of Ca 2+, Sr 2+, Ba 2+, Ag +, and Pb 2+. When we classify a compound as soluble it means that the compound is, it is also known as a(n) because it conducts electricity. NaNO3 (s) H 2 O Na + (aq) + NO3 (aq) When we classify a compound as insoluble it means that the compound is a, it is also known as a(n) because it doesn t conduct electricity. CH3OH (l) BaSO4 (s) H 2 O H 2 O CH3OH (aq) BaSO4 (aq) Page 8
CONCEPT: WRITING CHEMICAL REACTIONS EXAMPLE: Predict whether a reaction occurs, and write the balanced molecular equation. a. LiOH (aq) + MgSO4 (aq) EXAMPLE: Predict whether a reaction occurs, and write the balanced molecular equation, the total and net ionic equations. Molecular: Na 2 CO 3 (aq) + HBr (aq) Total Ionic: Net Ionic: Page 9
PRACTICE: Predict whether a reaction occurs, and write the balanced total and net ionic equations. Molecular: Ag 2 SO 4 (aq) + KCl (aq) Total Ionic: Net Ionic: PRACTICE: Predict whether a reaction occurs, and write the balanced total and net ionic equations. Molecular: MgBr 2 (aq) + NaC 2 H 3 O 2 (aq) Total Ionic: Net Ionic: Page 10
CONCEPT: ELECTROLYTES Whenever we add a solute into a solvent three outcomes are possible: the solute will dissolve ( STRONG electrolytes). the solute will dissolve ( WEAK electrolytes). the solute will dissolve ( NON electrolytes). Classification of Solutes in Aqueous Solution STRONG ELECTROLYTES WEAK ELECTROLYTES NONELECTROLYTES 1. STRONG ACIDS: HCl,, HI, HNO3,,,. 2. STRONG BASES: Group 1A Metal with OH -, H -, O 2- or NH2-1. WEAK ACIDS: HF,,,,. 2. WEAK BASES: Be(OH)2, Mg(OH)2,,. 1. MOLECULAR COMPOUNDS: C6H12O6 (glucose) C12H22O11 (sucrose) Groups 2A Metal, Calcium or Lower, with OH -, H -, O 2- or NH2-3) SOLUBLE IONIC COMPOUNDS: Page 11
PRACTICE: ELECTROLYTES EXAMPLE: Each of the following reactions depicts a solute dissolving in water. Classify each solute as a strong electrolyte, a weak electrolyte or a non-electrolyte. a. PbSO4 (s) PbSO4 (aq) b. HC2H3O2 (aq) H + (aq) + C2H3O2 (aq) c. CaS (s) Ca 2+ (aq) + S 2- (aq) d. Hg (l) Hg (aq) PRACTICE: Classify each of the following solutes as either a strong electrolyte, a weak electrolyte or a non-electrolyte. a. Perbromic acid, HBrO4 b. Lithium chloride, LiCl c. Formic Acid, HCO2H d. Methylamine, CH3NH2 e. Zinc bromide, ZnBr2 f. Propanol, C3H8OH Page 12
CONCEPT: OXIDATION-REDUCTION REACTIONS Chemists use some important terminology to describe the movement of electrons. In reactions we have the movement of electrons from one reactant to another. L E O Agent G E R Agent Rules for Assigning an Oxidation Number (O.N.) A. General Rules 1. For an atom in its elemental form (Na, O2, S8, etc.): O.N. = 0 2. For an ion the O.N. equals the charge: Na +, Ca 2+, NO3 B. Specific Rules 1. Group 1A: O.N. = +1 2. Group 2A: O.N. = +2 3. For hydrogen: O.N. = +1 with nonmetals O.N. = -1 with metals and boron 4. For Fluorine: O.N. = -1 5. For oxygen: O.N. = -1 in peroxides (X2O2, X = Group 1(A) element) O.N. = 1 in superoxides (XO2, X = Group 1(A) element) 2 O.N. = - 2 in all other compounds 6. Group 7A O.N. = -1 (except when connected to O) Page 13
CONCEPT: OXIDATION-REDUCTION REACTIONS (PRACTICE) EXAMPLE: In the following reaction identify the oxidizing agent and the reducing agent: a. 2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (g) PRACTICE: What is the oxidation number of each underlined element? a. P4 b. BO3 3- c. AsO4 2- d. HSO4 PRACTICE: In the following reaction identify the oxidizing agent and the reducing agent: a. Cr2O7 2- + 6 Fe 2+ + 14 H + 2 Cr 3+ + 6 Fe 3+ + 7 H2O Page 14
} } CHEMISTRY - CLUTCH CONCEPT: BASIC REDOX CONCEPTS OXIDATION-REDUCTION (REDOX) reactions deal with the transfer of electrons from one reactant to another. Lose Electrons Oxidation } Element becomes more positive Li (s) + Cl 2 (g) Li + (aq) + 2 Cl aq) Oxidation Gain } Element Number Electrons becomes Increases more negative Reduction } } Oxidation Number Decreases Reducing Agent (Reductant) Oxidizing Agent (Oxidant) Li (s) Li + (aq) + e Cl 2 (g) + 2 e 2 Cl (aq) Electrical Charge The units for electrical charge are measured in (C). (1.602 10 19 C) (6.022 10 23 mol 1 ) = 9.647 104 C 1 mole e charge mole e Charge of 1 electron Faraday Constant q = n F Faraday Constant Electrical Current The units for electrical current are in (A). I = q t Current Charge Time Electrical Voltage The relationship between work and voltage can be expressed as: w = E q Work Voltage Charge The relationship between Gibbs Free Energy and electric potential can be expressed as: ΔG = n F E Gibbs Free Energy mole e Faraday Constant Voltage Ohm's Law The units for resistance are in (Ω). I = E R Current Voltage Resistance Power Power represents work done per unit of time. The units for power are in (W). P = E I Power Voltage Current Page 15
CONCEPT: BALANCING REDUCTION-OXIDATION REACTIONS When balancing a redox reaction we balance them in terms of the number of electrons transferred between reactants. Balancing A Redox Reaction in Acidic Reactions: STEP 1: Write the equation into 2 half-reactions. STEP 2: Balance elements that are not oxygen or hydrogen. STEP 3: Balance Oxygens by adding. STEP 4: Balance Hydrogens by adding. STEP 5: Balance overall charge by adding e to the more side. Both reactions must have an equal number of e. STEP 6: Combine the half-reactions and cross out reaction intermediates. Balancing A Redox Reaction in Basic Reactions: Follow Steps 1-6 from above. STEP 7: Balance remaining H + by adding an equal amount ions to both sides of the chemical reaction. EXAMPLE: Balance the following reaction in an acidic solution. O 2 + F2 O2 + F Page 16
PRACTICE: BALANCING REDUCTION-OXIDATION REACTIONS EXAMPLE 1: Balance the following reaction in an acidic solution. NO2 NO3 + NO EXAMPLE 2: Balance the following reaction in a basic solution. Cr2O7 2 + Hg Hg 2+ + Cr 3+ Page 17
CONCEPT: SINGLE REPLACEMENT REACTIONS In a single replacement or displacement reaction one element displaces another element from a compound. More reactive or active metals displace less reactive metals or hydrogen from compounds. Metals Lithium (Li) > Potassium (K) > Barium (Ba) > Strontium (Sr) > Calcium (Ca) > Sodium (Na) Activity Metals in this category can displace hydrogen from liquid water, steam and acids: Metals Magnesium (Mg) > Aluminum (Al) > Zinc (Zn 2+ ) > Chromium (Cr 2+, Cr 3+ ) > Iron (Fe 2+, Fe 3+ ) Activity Metals in this category can displace hydrogen from steam and acids: Activity Increases Metals Cadmium (Cd 2+ ) > Cobalt (Co 2+, Co 3+ ) > Nickel (Ni 2+ ) > Tin (Sn 2+, Sn 4+ ) > Lead (Pb 2+, Pb 4+ ) Activity Metals in this category can displace hydrogen from acids: Hydrogen and Metals Hydrogen (H) > Antimony (Sb 3+ ) > Arsenic (As 3+, As 5+ ) > Bismuth (Bi 3+ ) > Copper (Cu +, Cu 2+ ) > Mercury (Hg2 2+, Hg 2+ ) > Silver (Ag + ) > Palladium (Pd 3+ ) > Platinum (Pt 2+, Pt 3+ ) > Gold (Au +, Au 3+ ) Page 18
PRACTICE: SINGLE REPLACEMENT REACTIONS EXAMPLE 1: Based on your understanding of activities determine if a reaction occurs and if so provide the products formed. Ba (s) + H2O (g) EXAMPLE 2: Based on your understanding of activities determine if a reaction occurs and if so provide the products formed. Zn (s) + NiCl2 (aq) EXAMPLE 3: If the activity of halogens is stated as: Fluorine > Chlorine > Bromine > Iodine, determine if a reaction occurs and if so provide the products formed. Cl2 (g) + AlBr3 (aq) Page 19
16. How many milligrams of NaCN are required to prepare 712 ml of 0.250 M NaCN? Page 20
17. What volume (in µl) of 0.100 M HBr contains 0.170 moles of HBr? Page 21
18. How many moles of Ca 2+ ions are in 0.100 L of a 0.450 M solution of Ca3(PO4)2? Page 22
19. How many chloride ions are present in 65.5 ml of 0.210 M AlCl3 solution? a) 4.02 1023 chloride ions b) 5.79 1024 chloride ions c) 2.48 1022 chloride ions d) 8.28 1021 chloride ions e) 1.21 1022 chloride ions Page 23
22. To what final volume would 100 ml of 5.0 M KCl have to be diluted in order to make a solution that is 0.54 M KCl? Page 24
23. If 880 ml of water is added to 125.0 ml of a 0.770 M HBrO4 solution what is the resulting molarity? Page 25
26. Consider the following balanced redox equation: H2O + 2 MnO4 + 3 SO3 2-2 MnO2 + 3 SO4 2- + 2 OH How many grams of MnO2 (MW: 86.94 g/mol) are produced when 32.0 ml of 0.615 M MnO4 - (MW: 118.90 g/mol) reacts with excess water and sulfite? Page 26
27. Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation: Cr2O7 2- + 6 Fe 2+ + 14 H + 2 Cr 3+ + 6 Fe 3+ + 7 H2O If it takes 35.0 ml of 0.250 M FeCl2 to titrate 50 ml of a solution containing Cr2O7 2-, what is the molar concentration of Cr2O7 2-? Page 27
28. Vinegar is a solution of acetic acid, CH3COOH, dissolved in water. A 5.54 g sample of vinegar was neutralized by 30.10 ml of 0.100 M NaOH. What is the percent by weight of acetic acid in the vinegar? Page 28
29. What is the molar mass of a 0.350 g sample of a monoprotic acid if it requires 50.0 ml of 0.440 M Ca(OH)2 to completely neutralize it? Page 29
30. Give the complete ionic equation for the reaction (if any) that occurs when aqueous solutions of sodium sulfide and copper (II) nitrate are mixed. a) Na+ (aq) + SO4 2- (aq) + Cu+(aq) + NO3 - (aq) CuS(s) + Na+(aq) + NO3 - (aq) b) Na+ (aq) + S-(aq) + Cu+(aq) + NO3 - (aq) CuS(s) + NaNO3(aq) c) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3 - (aq) Cu2+(aq) + S2-(aq) + 2 NaNO3(s) d) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3 - (aq) CuS(s) + 2 Na+(aq) + 2 NO3 - (aq) e) No reaction occurs. Page 30
31. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed. a) H+(aq) + OH-(aq) H2O(l) b) 2 K+(aq) + SO4 2- (aq) K2SO4(s) c) H+(aq) + OH-(aq) + 2 K+(aq) + SO4 2- (aq) H2O(l) + K2SO4(s) d) H2 2+ (aq) + OH-(aq) H2(OH)2(l) e) No reaction occurs. Page 31
32. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of Na2CO3 and HCl are mixed. a) 2 H+(aq) + CO3 2- (aq) H2CO3(s) b) 2 Na+(aq) + CO3 2- (aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 NaCl(aq) c) 2 H+(aq) + CO3 2- (aq) H2O(l) + CO2(g) d) 2 Na+(aq) + CO3 2- (aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq) e) No reaction occurs. Page 32