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huskyct uconn

huskyct uconn

Description

School: University of Connecticut
Department: Engineering
Course: General Chemistry
Professor: Tom wood
Term: Spring 2017
Tags: Chemistry, Catalysis, gas, Equilibrium, acids and bases, Bronsted-Lowry, le, and LeChatelier's Principle
Cost: 50
Name: Exam 2 CHEM1128 UConn Study Guide
Description: The exam will be covering Chapter 13 and Chapter 14 Chapter 13 is mainly about gas equilibrium, catalysis, equilibrium constant Chapter 14 is about acid-base
Uploaded: 03/06/2017
19 Pages 192 Views 0 Unlocks
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5-day Chemistry Exam Study Plan Day 1 Day 2 Day 3 Day 4 Day 5 +​Go to SI study session +Finish review question +Review Chapter 13 Concepts and Videos on Husky CT +Practice the problems in the videos on Husky CT for chapter 13 +​Finish practice exam and re-do the incorrect one +Review Chapter 14 concepts and videos on husky ct +Find some more practice test online if you don’t think you nail that concept just yet. (Go to office hours if there are any confusion) +​Go to SI study session +Review Chapter 14 concepts +Try to finish ALEKs as much as possible +Try to finish Discussion problems as much as possible (Go to office hours if there are any confusion) +​Go to SI study session +Memorize concepts +Re-do exercises that you are still confused about +Write out commonly made mistakes +Get plenty of sleep (Go to office hours if there are any confusion) +​Try to be as relaxed as possible +Review concepts and equations


What is the equilibrium constant, K, at this temperature for the decomposition reaction?



Don't forget about the age old question of uwec biology

Chapter 13 Notes 13.1 Chemical Equilibria: Equilibrium: the forward and reverse reactions occur at equal rate and concentration of reactants and products remain constant -> Chemical Equilibrium is a dynamic process 13.2 Equilibrium Constants Reaction Quotient: Qc= [C]x[D]y/ [A]m[B]n Law of mass action: When a mixture of reactants and products reaches equilibrium, its reaction quotient always has the same value For dilute solutions, a substance's activity, and its molar concentrations are equal. Activities for pure condensed phases = 1Heterogeneous equilibria: solids and liquids do not appear in equilibrium constant 13.3 Le chatelier Principle Disturbance Observed Change as Equilibrium is Resored Direction of Shift reactant added added reactant is consumed, more products will be produced toward product product added more reactants is remained in order to compensate for the products added toward reactants decreases in volume -> increase pressure decrease pressure toward the side with fewer moles of gas increase in volume, decrease in pressure pressure increased toward the side with more gas molecules temperature is increased heat is absorbed endothermic: toward products exothermic: toward reactants temperature is decreased heat is given off endothermic: toward reactants exothermic: toward products


What is the mass percent of acetic acid in the vinegar?




How to choose the appropriate acid-base indicator color over a range of pH values rather than at a specific pH?



If you want to learn more check out matthew lang ucr
If you want to learn more check out micah moore college
If you want to learn more check out arndt schultz principle
Don't forget about the age old question of choda gave a speech about growing up in israel. as support, she discussed her home life, her schooling, and her life in the military. what type of supporting material is this?
If you want to learn more check out acid and base nomenclature

Explanation: In general, if a system at equilibrium is disturbed by any change in concentration, the system will shift to counteract the change Effects of pressure on equilibrium ● Pressure: caused by gas molecules hitting sides of containers ● The more gas molecules there are, the higher pressure there would be ● So in general, increasing the pressure would make the reaction favors the side where it has less molecules to relieve the stress since high pressure means high concentration of molecules. Effects of temperature on equilibrium ● Endothermic reaction: Ex: heat + A + B -> C + D , ∆H>0 Increase temperature -> increase internal energy (heat) → the reaction would shift to the right to relieve the stress -> Form more products, less reactants Decrease temperature -> decrease internal energy (heat)-> the reaction would shift to the left in order to recover the energy lost -> Form less products, more reactants ● Exothermic reactionEx: A+B-> C+D+heat , ∆H<0 Increase temperature -> increase internal energy (heat) -> reaction would shift to the left to reduce heat and relieve the stress-> More reactants would be formed Decrease temperature -> decrease internal energy (heat)-> reaction would shift to the right to increase energy (heat) -> More products, less reactants Catalyst do not change equilibrium Conversion P=MRT • KP = Kc(RT)Δn Commonly made mistakes -Le Chatelier Principle -Forget to determine which side of the reaction is favored Sample problems 1) A sample of solid ammonium chloride was placed in an evacuated flask and then heated to a certain temperature so that it decomposed to ammonia gas and hydrogen fluoride gas. After heating the total pressure inside the container is 4.4 atm. What is the equilibrium constant, K, at this temperature for the decomposition reaction? Chapter 14 14.1 Bronsted Lowry ● Bronsted Lowry Acid: Compound that donates proton to another compound ● Bronsted Lowry Base: Comppound that accepts a proton ● Conjugate base of acid: product that remains after acid ● Conjugate acid: product that results when a base accepts the proton ● Amphoteric: species that may either act as an acid or base by any definition. Ex: bicarbonate ion 14.2. pH and pOH: ● pH=-log[H3O+] ● [H3O+]=10-pH ● pOH=-log[OH-] ● KW=[H3O+][OH-] ● Kw=pH+pOH=14 14.3 Relative Strengths of Acids and Base: ● 6 strong acids: HClO4, HCl, HBr, HI,HNO3, H2SO4 ● 6 strong base: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 ● Weak acids and bases do not ionize completely in waterEffects of molecular structure on acid-base strength ● H-A acid: increases as bond energy decreases ● Oxyacids: increasing electronegativity, increasing acidity ● Amphoteric: act as acids when react with strong bases, act as bases when react with strong acids ● Base strength increases up a group and decrease across the table 14.4 Hydrolysis of Salt Solutions: ● Strong acid + strong base= neutral ● Strong acid + weak base= weak acid ● Weak acid + strong base= weak base ● Weak acid + weak base = acid, base, neutral 14.5 Polyprotic acids: ● Monoprotic acid: HCl, HNO3, HCN ● Diprotic acid: H2SO4 ● Triprotic acid: H3PO4 14.6 Buffers: ● Buffer capacity: amount of an acid or base that can be added to a volume of a buffer solution before its pH changes significantly (usually by one pH unit) Selection of suitable buffer mixtures ● Good buffer- equal concentrations of both of its components ● Weak acids need buffers that have pH <7, weak bases need buffers that have pH>7 14.7 Acid-Base Titration: Titration of a strong acid with a strong base: the equivalence point occurs at a pH of 7.00 Titration of a weak acid with a strong base: pH > 7 due to a weak acid Titration ● At half equivalence point, pKa=pH How to choose the appropriate acid-base indicator color over a range of pH values rather than at a specific pH? At the equivalence point in the titration of a weak base with a strong acid, the resulting solution is slightly acidic due to the presence of the conjugate acid. Thus, pick an indicator that changes color in the acidic range and bracket the pH at the equivalence point. Equations Summary Kw = [H3O+][OH−] = 1.0 × 10−14 (at 25 °C) pH = −log[H3 O+] pOH = −log[OH−][H3O+] = 10−pH [OH−] = 10−pOH pH + pOH = pKw = 14.00 at 25 °C Ka=[H3O+][A−] [HA] Kb = [HB+][OH−] [B] Ka × Kb = 1.0 × 10−14 = Kw Percent ionization = ([H3 O+]eq/[HA] 0 ) × 100 pKa = −log Ka pKb = −log Kb pH=pKa+log[A−] Practice Exam What is the molarity of H+in a solution formed by mixing 25.00 mL of 0.2345 M Ba(OH)2 with 100.00 mL of 0.1542 M HCl? ○ 7.646 x 10-2 M ○ -7.646 x 10-2 M ○ 8.03 x 10-2 M ○ 2.956 x 10-2 M ○ 3.695 x 10-3 M Vinegar is a solution of acetic acid, CH3COOH, dissolved in water. A 5.54 g sample of vinegar is neutralized by 30.10 mL of 0.100 M NaOH. What is the mass percent of acetic acid in the vinegar? ○ 0.184% ○ 1.63% ○ 3.26% ○ 5.43% ○ 9.23% Lactic acid, C3H6O3, is the acid present in sour milk. A 0.100 g sample of pure lactic acid requires 12.95 mL of 0.0857 M sodium hydroxide for complete reaction. How many moles of hydroxide ion(s) is (are) required to neutralize one mole of lactic acid? ○ 1 ○ 2 ○ 3 ○ 6 ○ none At 25°C, the equilibrium constant for the following reaction is 5.0 x 104: Ni (s) + 4 CO (g) Ni(CO)4 (g) What is K for the following reaction?2 Ni(CO)4 (g) 2 Ni (s) + 8 CO (g) ● 4.0 x 10-10 ● 2.5 x 104 ● 2.0 x 10-5 ● 1.9 x 1019● 2.3 x 10-10 Consider the following reaction Sb2S3(s) + 3H2(g) 2Sb (s) + 3H2S (g) Where the equilibrium constant, K, is 0.429 at a certain temperature. If the equilibrium partial pressure of H2S is 0.200 atm, what is the partial pressure of H2(g)? ○0.466 atm ○0.151 atm ○0.265 atm ○0.200 atm ○0.429 atm Assuming the reaction below is at equilibrium, which of the following changes will drive the reaction to the left? 2NOI (g) 2NO (g) + I2(g) ΔH° = 45.3 kJ ●Increasing the temperature ●Compressing the container ●Removing NO ●i only ●ii only ●iii only ●i and iii ●ii and iii The equilibrium constant for the reaction below is 6.0 x 105 at 25°C. N2(g) + 3H2(g) 2NH3(g) ΔH° = -92.2 kJ At what temperature is the equilibrium constant equal to 1.0 x 103? ●-188 °C ●163 °C ●37 °C ●87 °C●2527 °C In Chapter 13, the Brønsted-Lowry acid-base model was introduced. Please choose which of the following statements fits best the Brønsted-Lowry model: ●A base is a substance that produces an excess of H+ ●A base is an H+ acceptor ●An acid is an H+ acceptor ●An acid is a substance that donates an electron pair The ion product constant of water is Kw= 1.0 x 10-14. Please indicate which of the following is incorrect: ●Water is amphoteric in nature ●Pure water has [H+] = 1.0 x 10-7 ●The Kwof water is higher than that of a typical weak acid ●A solution of water with [H+] = 2.0 x 10-9 is basic ● What is the pOH of a 1.00 x 10-2 M Ba(OH)2solution? ● pOH = 1.70 ● pOH = 2.01 ● pOH = 11.9 ● pOH = 12.3 There are many metal cations that act as weak acids in water solution. Please indicate which of the following metal ions behave as weak acids: ● Na+ ● Ba2+ ● K+ ● Al3+ Please identify the correct formula of percent ionization, where the subscripts “eq” and “o” indicate equilibrium and starting concentrations, respectively: ● % ionization = ([H+]o/[HB]o) x 100%● % ionization = ([H+]eq/[HB]o) x 100% ● % ionization = ([H+]eq/[HB]eq) x 100% ● % ionization = ([H+]o/[HB]eq) x 100% Which of the following is not true​ for polyprotic weak acids? ● The anion formed in one ionization step produces another H+ in the successive ionization step. ● The acid equilibrium constant becomes larger with each successive ionization step. ● The acid equilibrium constant becomes smaller with each successive ionization step. ● Ka1> Ka2> … Salts, dissolve in water, and ionize completely. After that, the liberated ions participate in acid/base equilibrium with water and alter the pH of pure water from its original pH value of 7. Please indicate which of the following is correct: ● Dissolving NaCl in pure water produces an acidic pH ● Dissolving NaCl in pure water produces a basic pH ● Dissolving K2CO3in pure water produces an acidic pH ● Dissolving K2CO3in pure water produces a basic pH Which of the following chloric acids, when dissolved individually at the same concentration in pure water, produce a less acidic solution? ● HClO ● HClO2 ● HClO3 ● HClO4 • You have a weak acid buffer. For every 1 mole of conjugate base there are 2 moles of conjugate acid. You measure the pH of the solution and found it to be 7.25 at equilibrium. Which of these is the unknown weak acid? ● Lactic acid pKa= 3.85 ● Carbonic acid pKa= 6.36 ● Dihydrogen phosphate ion pKa= 7.21 ● Hypochlorous acid pKa= 7.55● Ammonium ion pKa= 9.25 Which of the following weak acids would be the best choice to use if you needed to make a buffer at pH = 4? ● Lactic acid/lactate ion pKa= 3.85 ● Acetic acid/acetate ion pKa= 4.74 ● Carbonic/hydrogen carbonate ion acid pKa= 6.36 ● Dihydrogen phosphate ion/hydrogen phosphate ion pKa= 7.21 A buffer is made by mixing 300 mL of 0.500 M KH2PO4and 300 mL of 0.317 M K2HPO4. Assuming the volumes are additive, calculate the pH of the resulting buffer. The KH2PO4has a Ka2of 6.2 x 10-8. ● 7.41 ● 7.21 ● 7.01 ● 6.71 ● 6.99 Which of the following has the highest buffer capacity? ● 0.10 M H2PO4-/0.10 M HPO42- ● 0.50 M H2PO4-/0.10 M HPO42- ● 0.10 M H2PO4-/0.50 M HPO42- ● 0.50 M H2PO4-/0.50 M HPO42- ● They all have the same buffer capacity. A buffer solution is made by mixing 200 mL of 0.25 M acetic acid and 200 mL of 0.11 M acetate ion. Which of the following actions can be carried out without​ changing the pH of the solution? ● Dilute the acetate ion to 0.50 M before mixing ● Dilute the acetic acid to 0.11 M before mixing ● Add acetic acid after mixing ● Add 10 mL of water after mixing Which of the following statement(s) concerning acid-base indicators is (are) true? ● Acid-base indicators are derived from weak acids and weak bases.● The acid (HIn) and base (In-) forms of the indicator have different colors. ● If ([HIn]/[In-]) 10, the base (In-) color will be seen. ● i only ● ii only ● iii only ● i and ii ● i and iii A 30.00 mL sample of vinegar which contains acetic acid is titrated with 0.4190 M NaOH (aq). If the titration requires 27.83 mL of NaOH, what is the concentration of acetic acid in the vinegar? ● 0.1931 M ● 0.2016 M ● 0.2174 M ● 0.3887 M ● 0.4517 M A 50.0 mL sample of 0.50 M HCl is titrated with 0.50 M KOH. What is the pH of the solution after 28.0 mL of KOH have been added to the acid? ● 0.851 ● 1.49 ● 2.85 ● 2.96 ● 13.15 A 20.0 mL sample of 0.30 M monoprotic weak acid was titrated with 0.30 M NaOH. The following data were collected during the titration. mL NaOH added 5.00 10.00 15.00 20.00 pH 6.98 7.46 7.93 10.31 What is the acid dissociation constant, Ka, for the monoprotic weak acid? ● 1.1 x 10-7 ● 3.5 x 10-8● 4.9 x 10-11 ● 1.2 x 10-4 When 20.0 mL of 0.15 M nitric acid is mixed with 20.0 mL of 0.10 M barium hydroxide, what is the pH of the resulting solution? ● 0.00 ● 0.82 ● 1.60 ● 7.00 ● 12.40 Calculate the concentration of Co2+ dissolved in a saturated aqueous solution of Co(OH)2precipitate at pH of 9.0. The Kspof Co(OH)2is 2.5 x 10-15. ● [Co2+] = 2.5 x 10-5 ● [Co2+] = 1.3 x 10-4 ● [Co2+] = 2.7 x 10-3 ● [Co2+] = 1.5 x 10-2 The Kffor Co(NH3)62+ is 1.0 x 10-5 and the Kspfor Co(OH)2is 2.5 x 10-15. Which is the correct equilibrium constant (K) for the following reaction? Co(OH)2(s) + 6 NH3(aq) Co(NH3)62+ (aq) + 2 OH- (aq) ● K = 2.5 x 10-20 ● K = 2.5 x 10-10 ● K = 1.0 x 10-5 ● K = 4.0 x 109 ● K = 4.0 x 1019 The addition of aqueous NH3dissolves the AgCl precipitate(s). Indicate which statement is correct: ● The AgCl (s) solubilizes due to a common ion effect. ● The AgCl (s) solubilizes due to increased acidity. ● The AgCl (s) solubilizes due to increased basicity. ● The AgCl (s) solubilizes due to NH3– assisted complexation.Which one of the following expressions is correct for the representation of Ca2+ (aq) concentration involved in the solubility product (Ksp) of Ca3(PO4)2in the presence of 0.10 M of Na3PO4: ● [Ca2+] = (Ksp/0.010)1/2 ● [Ca2+] = (Ksp/0.010)1/3 ● [Ca2+] = (Ksp/0.0010)1/2 ● [Ca2+] = (Ksp/0.0010)1/3 Which statement is not true​ for the selective precipitation of Ag+ and not of Pb2+ using chloride ions (Kspof AgCl = 1.8 x 10-10 and Kspof PbCl2= 1.7x 10-5) ● The reaction quotient for AgCl must be bigger than Kspof AgCl. ● The reaction quotient for PbCl2must be smaller than Kspof PbCl2. ● The Cl- concentration must be greater than (1.8 x 10-10)/([Ag+]) and smaller than [(1.7 x 10-5)/([Pb2+])]0.5. ● The Cl- concentration must be smaller than (1.8 x 10-10)/([Ag+]) and greater than [(1.7 x 10-5)/([Pb2+])]0.5. Which statement is not true​ for the dissolution of ZnCO3(s) in acid: ● The dissolution of ZnCO3is facilitated by the formation of a weak acid. ● The dissolution of ZnCO3is facilitated by the formation of a complex ion. ● The dissolution of ZnCO3is facilitated by the evolution of a gas. ● The dissolution of ZnCO3is facilitated by the decomposition of carbonic acid to CO2and H2O. Consider that the concentrations of all complexing agents (i.e. CN-, OH-, NH3, etc.) are equal to 1.0 M. Please indicate in which solution the concentration of the uncomplexed metal ion would be lower: ● Zn(OH)42-, Kf= 3.0 x 1015 ● Cu(NH3)42+, Kf= 5.6 x 1011 ● Al(OH)4-, Kf= 1.0 x 1033 ● Ag(CN)2-, Kf= 2.0 x 1020 Open ResponseA. (8 points)​ A 1.832 g sample of a diprotic acid was dissolved in water. It took 20.27 mL of a 0.1578 M NaOH solution to neutralize the acid. What is the molar mass of the diprotic acid? (one mole of acid reacts with two moles of hydroxide ions) __________________________ B. (12 points)​ Consider the following equilibrium C (s) + CO2(g) 2CO (g) When the system is at equilibrium at 973K in a 2.0 L flask, 0.10 mol CO, 0.20 mol CO2, and 0.40 mol C are present. When the system is cooled to 873 K, and additional 0.040 mol of C forms. What is the equilibrium constant K at 873 K? __________________________C. (12 points)​ The two ionization constants of carbonic acid (H2CO3) are: Ka1= 4.4 x 10-7 and Ka2= 4.7 x 10-11. Assuming you start with 1.0 x 10-3 M solution of H2CO3, calculate: (1)​ the pH of the resulting solution, and (2)​ the resulting concentration of H2CO3. (1)​________________________ (2)​________________________ D. (12 points) ​Assume that you have a 0.20 M solution of acetic acid (CH3COOH). Acetic acid is a weak acid (Ka= 1.8 x 10-5) and it is also written as HC2H3O2. 88 mL of that solution is added to 500 mL of H2O. What would be the pH value of the final solution?__________________________ E. (12 points)​ A buffer solution is made up using acetic acid/acetate (Ka= 1.8 x 10-5) with a pH = 4.95. 1)​ If this buffer solution contains 1.0 M acetic acid what concentration of acetate ion must it contain?__________________________ 2)​ If the pH of the buffer is adjusted to 5.25 by adding a strong base, what concentration of acetic acid was converted to acetate ion? _________________________ F. (15 points)​ A 25.0 mL sample of 0.10 M NH3(ammonia), a weak base with a Kbof 1.8 x 10-5 is titrated with 0.15 M HCl. 1. What is the pH of the solution before any HCl is added?__________________________ 2. What volume of HCl is required to reach the equivalence point? __________________________ 3. What is the pH of the solution at the equivalence point? _________________________ G. (9 points)​ Consider an unknown weak acid which is being titrated with a strong base. 1. A 50.0 mL aqueous solution containing 2.500 g of the weak acid mentioned above is titrated with Ba(OH)2. The titration requires 60.00 mL of 0.425 M Ba(OH)2to reach the equivalence point. What is the molar mass of the unknown weak acid?__________________________ 2. A second experiment uses a 50.0 mL solution of the unknown weak acid identical to what was used in the first experiment. To this solution 30.00 mL of 0.425 M Ba(OH)2was added. The pH after the addition of Ba(OH)2is 5.38. What is the Kafor the unknown acid? __________________________
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