CHEM129 Practice Problems

Determine the percent ionization of 0.125 M HCN solution Ka HCN: 4.9 x 10-10

0.0063%

Calculate the molar solubility of AgCl in 1.0 M NH3. Ksp for AgCl is 1.8 x 10-10; Kf for Ag(NH3)2 + ion is 1.7 x 107

0.050 M

If 25.98 mL of 0.1180 M KOH solution reacts with 52.50 mL of CH3COOH solution, what is the molarity of the acid solution?

0.05839 M

What is the component concentration ratio, [BrO-]/[HBrO] of a buffer that has a pH of 7.95? Ka HBrO: 2.3 10-9

0.20

What is the molar solubility of AgBr in 1.0 M Na2S2O3. Silver ion forms the complex ion Ag(S2O3)2 2-. Ksp of AgBr = 5.0 x 10-13, Kf for Ag(S2O3)2 2- is 2.9 x 1013

0.44 M

As an FDA physiologist, you need 0.700 L of formic acid-formate buffer with a pH of 3.74. What is the required buffer component concentration ratio?

0.99

What is the concentration of Cu2+ (aq) in a solution that was originally 0.015 M Cu(NO3)2 and 0.100 M NH3? The Cu2+ ion forms the complex ion Cu(NH3)4 2+ with a Kf of 4.8 x 1012

1.2 x 10-9

What is the pH of 0.150 M KCN? Ka HCN: 6.2 x 10-10

11.19

A 0.552-g sample of ascorbic acid (Vitamin C) was dissolved in water to a total volume of 20.0 mL and titrated with 0.1103 M KOH. The equivalence point occurred at 28.42 mL. The pH of the solution at 10.0 mL of added base was 3.72. From this data, determine the molar mass and Ka for vitamin C.

176 g/mole; 1.0 x 10-4

Use the Ksp value from the Table to calculate the solubility of iron(II) hydroxide in pure water in grams per 100.0 mL of solution. Ksp Fe(OH)2: 4.87 x 10-17

2.07 x 10-5 g/100 mL

If 15.0 mL of glacial acetic acid (pure HC2H3O2) is diluted to 1.50 L with water, what is the pH of the resulting solution? The density of acetic acid is 1.05 g/mL. Ka HC2H3O2: 1.8 x 10-5

2.75

A 30.0 mL sample of 0.165 M propanoic acid is titrated with 0.300 M KOH. Calculate the pH at each volume of added base: 0 mL, 5 mL, 10 mL, equivalence point, one-half equivalence point, 20 mL, 25 mL. Ka Propanioc: 1.3 x 10-5

2.83, 4.52, 5.07, 8.96, 4.89, 12.32, 12.67

A 20.0 mL sample of 0.115 M H2SO3 solution is titrated with 0.1014 M KOH. At what added volume of base does each equivalence point occur? Ka H2SO3: Ka1 = 1.6 x 10-2 Ka2 = 6.4 x 10-8

22.7 mL and 45.4 mL

A 0.5224-g sample of an unknown monoprotic acid was titrated with 0.0988 M NaOH. The equivalence point of the titration occurs at 23.82 mL. Determine the molar mass of the acid.

220 g/mole

What is the molar solubility of CdC2O4 in 0.10 M NH3? Ksp of CdC2O4 = 1.5 x 10-8, Kf of Cd(NH3)4 2+ = 1.0 x 107

3.0 x 10-3

A buffer is created by combining 150.0 mL of 0.25 M HCHO2 with 75.0 mL of 0.20 M NaOH. Determine the pH of the buffer. Ka HCOOH or HCHO2: 1.8 x 10-4

3.57

A 0.148 M solution of a monoprotic acid has a percent ionization of 1.55%. Determine the acid-dissociation constant for the acid.

3.61 x 10-5

What mass of sodium benzoate should you add to 150.0 mL of a 0.15 M benzoic acid solution to obtain a buffer with a pH of 4.25? (Assume no volume change.) Ka Benzoic(HC7H5O2): 6.5 x 10

3.7 g

What is the pH of 0.85 M NH4Br? Kb NH3: 1.76 x 10-5

4.66

A 0.25 mole sample of a weak acid with an unknown pKa was combined with 10.0 mL of 3.00 M KOH, and the resulting solution was diluted to 1.500 L. The measured pH of the solution was 3.85. What is the pKa of the weak acid?

4.73

Sufficient NaCN was added to 0.015 M AgNO3 to give a solution that was initially 0.100 M in CN1-. What is the concentration of Ag+ in this solution after Ag(CN)2 1- forms? Kf for Ag(CN)2 1- = 5.6 x 1018

5.5 x 10-19

What is the pH of 0.40 M triethylammonium chloride, (CH3CH2)3NHCl? Kb (CH3CH2)3N: 8.6 x 10-4

5.56

A metal ion uses d2sp3 orbitals when forming a complex. What is its coordination number and the shape of the complex?

6, octahedral

What is the concentration of Ag+ (aq) ion in 0.010 M AgNO3 that is also 1.00 M NH3? Kf for Ag(NH3)2 + ion is 1.7 x 107

6.1 x 10-10

A 0.150 M solution of morphine has a pH of 10.5. What is Kb for morphine?

6.7 x 10-7

A 0.185 M solution of a weak acid has a pH of 2.95. Calculate the acid-dissociation constant.

6.8 x 10-6

Caffeine (C8H10N4O2) has a pKb of 10.4. Calculate the pH of the solution containing a caffeine concentration of 455 mg/L

7.48

What is the pH of 0.65 M potassium formate, HCOOK? Ka HCOOH: 1.8 x 10-4

8.78

A buffer that contains 0.11 M HY and 0.220 M Y- has a pH of 8.77. What is the pH after 0.0015 mole of Ba(OH)2 is added to 0.350 L of this solution?

8.82

A 25.0 mL sample of 0.125 M pyridine is titrated with 0.100 M HCl. Calculate the pH at each volume of added acid: 0 mL,10 mL, 20 mL, equivalence point, one-half equivalence point, 40 mL, 50 mL Kb pyridine: 1.7 x 10-9

9.16, 5.56, 4.98, 3.24, 5.23, 1.87, 1.60

Which pair is a Bronsted-Lowry conjugate acid-base pair: a. NH3: NH4+ b. H3O+, OH1- c. HCl, HBr d. ClO4 1-, ClO3 1-

A

A complex, ML6 2+, is violet. The same metal forms a complex with another ligand, Q, that creates a weaker field. What color might MQ6 2+ be expected to show? Explain.

A violet complex absorbs yellow-green light. The light absorbed by a complex with a weaker ligand would be at a lower energy and longer wavelength. Light of lower energy than yellow-green light is yellow, orange, or red light. The color observed would be blue or green.

[Cr(H2O)6] 2+ is violet. Another CrL6 complex is green. Can ligand L be CN−? Can it be Cl−? Explain.

A violet complex is absorbing yellow-green light, and a green complex is absorbing red light. The green complex is absorbing lower energy light, so it must contain a weaker ligand than H2O. Thus, L could not be CN- (which is stronger than H2O), but could be Cl-

Two compounds with general formulas AX and AX2 have Ksp = 1.5 x 10-5. Which of the compounds has the higher molar solubility?

AX2

In H2O, HF is a weak acid and the other hydrohalic acids are equally strong. In NH3, however, all the hydrohalic acids are equally strong. Explain.

Ammonia, NH3 is a more basic solvent than H2O (more willing to accept an H+). In a more basic solvent, weak acids like HF act like strong acids and are 100% dissociated.

Consider three solutions I. 0.10 M solution of a weak monoprotic acid II. 0.10 M solution of a strong monoprotic acid III. 0.10 M solution of a weak diprotic acid Each solution is titrated with 0.15 M NaOH. Which quantity is the same for all three solutions? a. the volume required to reach the final equivalence point b. the volume required to reach the first equivalence point c. the pH at the first equivalence point d. the pH at one-half the first equivalence point

B

Methoxide ion, CH3O1-, and amide ion, NH2 1- are very strong bases that are "leveled" by water. What does this mean? Write the reactions that occur in the leveling process. What species do the two leveled solutions have in common.

Both methoxide ion and amide ion produce OH1- in aqueous solution. In water, the strongest base possible is OH1-. Since both bases produce OH1- in water, both bases appear equally strong. CH3O1- (aq) + H2O (l) à OH1- (aq) + CH3OH (aq) NH2 1- (aq) + H2O (l) à OH1- (aq) + NH3 (aq)

Two monoprotic acids (A and B) are titrated with identical NaOH solutions. The volume to reach the equivalence point for solution A is twice the volume required to reach the equivalence point for solution B, and the pH at the equivalence point of solution A is higher than the pH at the equivalence point for solution B. Which statement is true? a. The acid is solution A is more concentrated than in solution B and is also a stronger acid than that in solution B. b. The acid is solution A is less concentrated than in solution B and is also a weaker acid than that in solution B. c. The acid is solution A is more concentrated than in solution B and is also a weaker acid than that in solution B. d. The acid is solution A is less concentrated than in solution B and is also a stronger acid than that in solution B.

C

Use the table above to rank the following in order of increasing acid strength: HIO3, HI, CH3COOH, HF HIO3: 1.6 x 10-1 HI: -- CH3COOH: 1.8 x 10-5 HF: 6.8 x 10-4

CH3COOH < HF < HIO3 < HI

Given the Ka values above, identify the strongest conjugate base among: HNO2, HCHO2, HClO, HCN Ka HNO2: 4.6 x 10-4 HCHO2: 1.8 x 10-4 HClO: 2.9 x 10-8 HCN: 4.9 x 10-10

CN 1-

Which transition metals have a maximum oxidation state of +6?

Cr, Mo, and W

Which ion will form a basic solution when dissolved in water? (Hint: Write out the reaction of this base with water to see what could form.) a. Br 1- b. NO3 1- c. HSO4 1- d. SO3 2-

D: SO3 2- (aq) + H2O (l) ⇄ HSO3 1- (aq) + OH1- (aq)

Rank the following salts in order of increasing pH of their 0.1 M aqueous solutions: KNO3, K2SO3, K2S, Fe(NO3)2

Fe(NO3)2 < KNO3 < K2SO3 < K2S

Octahedral [Ni(NH3)6] 2+ is paramagnetic, whereas planar [Pt(NH3)4] 2+ is diamagnetic, even though both metal ions are d8 species. Explain.

In an octahedral d8 complex, two electrons occupy the two eg orbitals and will be unpaired. In a square planar d8 complex, the highest energy (dx2 - y2) orbital is unoccupied and all other levels are full, making the complex diamagnetic.

Three of the complex ions that are formed by Co3+ are [Co(H2O)6] 3+, [Co(NH3)6] 3+, and [CoF6] 3−. These ions have the observed colors (listed in arbitrary order) yellow, orange, green, and blue. Match each complex with its color. Explain.

NH3 > H2O > F- in ligand field strength, so [Co(NH3)6] 3+ would absorb light of highest energy (shortest wavelength), followed by [Co(H2O)6] 3+, and then [CoF6] 3- Yelloworange, green, and blue complexes would absorb blue-green, red, and orange light, respectively. Thus, [Co(NH3)6] 3+ absorbs blue-green and is yellow-orange in color, [Co(H2O)6] 3+ absorbs orange light and is blue in color, and [CoF6] 3- absorbs red light and is green in color.

A solution containing lead(II) nitrate is mixed with one containing sodium bromide to form a solution that is 0.0150 M Pb(NO3)2 and 0.00350 M in NaBr. Does a precipitate form in the newly mixed solution? Ksp PbBr2: 4.67 x 10-6

No PbBr2 will form. Qsp < Ksp

Some species with two oxygen atoms only are the oxygen molecule, O2, the peroxide ion, O2 2−, the superoxide ion, O2 −, and the dioxygenyl ion, O2 +. Draw an MO diagram for each, rank them in order of increasing bond length, and find the number of unpaired electrons in each.

O2: (σ2s)2(σ*2s)2(σ2p)2(π2p)2(π2p)2(π*2p)1(π*2p) 1 bond order = 2.0 paramagnetic (2 unpaired) O2 + (σ2s)2(σ*2s)2(σ2p)2(π2p)2(π2p)2(π*2p) 1 bond order = 2.5 paramagnetic (1 unpaired) O 2- (σ2s)2(σ*2s)2(σ2p)2(π2p)2(π2p)2(π*2p)2(π*2p) 1 bond order = 1.5 paramagnetic (1 unpaired) O2 2- (σ2s)2(σ*2s)2(σ2p)2(π2p)2(π2p)2(π*2p)2(π*2p) 2 bond order = 1.0 diamagnetic (0 unpaired) Bond length: O2 + < O2 < O2 - < O2 2-

For each of the following, identify the conjugate acid-base pairs (identify the acid and base in each pair). a. HSO4 1- + NO3 1- ⇄ HNO3 + SO4 2- b. HC2H3O2 + CO3 2- ⇄ HCO3 1- + C2H3O2 1- c. H2CO3 + CO3 2- ⇄ HCO3 1- + HCO3 1-

Species #1 Species #2 Species #3 Species #4 a. Acid Base Conjugate acid Conjugate base b. Acid Base Conjugate acid Conjugate base c. Acid Base Conjugate acid Conjugate base

The hexaaqua complex [Ni(H2O)6] 2+ is green, whereas the hexaammonia complex [Ni(NH3)6] 2+ is violet. Explain.

The aqua ligand is weaker than the ammine ligand. The weaker ligand results in a lower splitting energy and absorbs a lower energy of visible light. The green hexaaqua complex appears green because it absorbs red light (opposite side of the color wheel). The hexaammine complex appears violet because it absorbs yellow light, which is higher in energy (shorter λ) than red light.

A solution containing sodium fluoride is mixed with one containing calcium nitrate to form a solution that is 0.015 M in NaF and 0.010 M in Ca(NO3)2. Does a precipitate form in the mixed solution? If so, identify the precipitate. Ksp CaF2: 1.46 x 10-10

Yes, CaF2 will form

Predict whether a precipitate will form if you mix 75.0 mL of a NaOH solution with [OH-] = 2.6 x 10-3 M with 125.0 mL of a 0.018 M MgCl2 solution. Identify the precipitate, if any. Ksp Mg(OH)2: 2.06 x 10-13

Yes, Mg(OH)2 will form

Rank the following in order of increasing Δ and energy of light absorbed: [Cr(NH3)6] 3+, [Cr(H2O)6] 3+, [Cr(NO2)6] 3−.

[Cr(H2O)6] 3+ < [Cr(NH3)6] 3+ < [Cr(NO2)6] 3−

For each of the following, determine [H3O+], [OH1-], pH, pOH and whether or not the solution is acidic, basic, or neutral. a. A rain water sample with a pH of 4.35 b. An ammonia sample with a pH of 11.28 c. A sample of Ca(OH)2 with a concentration of 0.16 g Ca(OH)2/100.0 mL solution

[H3O+] [OH1-] pOH pH a. 4.5 x 10-5 M 2.2 x 10-10 M 9.65 -- Acidic b. 5.2 x 10-12 M 1.9 x 10-3 M 2.72 -- Basic c. 2.3 x 10-13 M 0.044 M 1.36 12.64 basic

Determine [OH 1-], pH, and pOH of a solution that is 0.125 M CO3 2-. Ka H2CO3: Ka1 = 4.3 x 10-7 Ka2 = 5.6 x 10-11

[OH 1-] = 0.00474 M; pH = 11.68; pOH = 2.32

The molecular orbitals depicted below are derived from 2p atomic orbitals in F2 +. (a) Give the orbital designations. (b) Which is occupied by at least one electron in F2 +? (c) Which is occupied by only one electron in F2 +? PICTURE IN WS14

a) A is the π*2p molecular orbital (two p orbitals overlapping side to side with a node between them); B is the σ2p molecular orbital (two p orbitals overlapping end to end with no node); C is the π2p molecular orbital (two p orbitals overlapping side to side with no node); D is the σ*2p molecular orbital (two p orbitals overlapping end to end with a node). b) F2 + has thirteen valence electrons: The MO electron configuration is(σ2s)2(σ*2s)2(σ2p)2(π2p)2(π2p)2(π*2p)2(π*2p)1 The π*2p molecular orbital, A, σ2p molecular orbital, B, and π2p molecular orbital, C, are all occupied by at least one electron. The σ*2p molecular orbital is unoccupied. c) A π*2p molecular orbital, A, has only one electron.

The orbital occupancies for the d orbitals of several complex ions are diagrammed below. a. Which diagram corresponds to the orbital occupancy of the cobalt ion in [Co(CN)6] 3−? b. If diagram D depicts the orbital occupancy of the cobalt ion in [CoF6] n, what is the value of n? c. [NiCl4] 2− is paramagnetic and [Ni(CN)4] 2− is diamagnetic. Which diagrams correspond to the orbital occupancies of the nickel ions in these species? d. Diagram C shows the orbital occupancy of V2+ in the octahedral complex VL6. Can you determine whether L is a strong- or weak-field ligand? Explain. PICTURE IN WS13

a) In [Co(CN)6] 3-, Co has a +3 charge and an electron configuration of [Ar]3d6. Since CN- is a strong-field ligand, the six electrons will pair in the t2g orbitals. Diagram A is the correct diagram. b) Diagram D shows seven electrons. The Co ion that is d7 and has a charge of +2. Since the six F- ligands each have a charge of -1, the complex ion must have a charge of -4. n = -4. c) The Ni ion in both complexes has a +2 charge and configuration of [Ar]3d8 . Since [Ni(CN)4] 2-, is diamagnetic, Diagram E with eight paired electrons is the correct diagram. The paramagnetic complex [Ni(Cl)4] 2-, is tetrahedral. Diagram B is correct. d) It is not possible to determine if the ligands in the VL6 complex are strong-field or weak-field. The V2+ ion is a d3 ion and the arrangement of three electrons in the d orbitals is the same in both the strong-field and weak-field cases.

Which of the following are Arrhenius acids? a. H2O b. Ca(OH)2 c. H3PO3 d. HI

a, c, and d

Give the charge and coordination number on the central metal ion(s) in each compound and name the following: a. [Ni(H2O)6]Cl2 b. [Cr(en)3](ClO4)3 c. K4[Mn(CN)6] d. K[Ag(CN)2] e. Na2[CdCl4] f. [Co(NH3)4(H2O)Br]Br2

a. +2, 6; hexaaquanickel(II) chloride b. +3, 6; tris(ethylenediamine)chromium(III) perchlorate c. +2, 6; potassium hexacyanomanganate(II) d. +1, 2; potassium dicyanoargentate(I) e. +2, 4; sodium tetrachlorocadmate(II) f. +3, 6; tetraammineaquabromocobalt (III) bromide

Give the number of d electrons (n of dn) for the central metal ion in a. [TiCl6] 2− b. K[AuCl4] c. [RhCl6] 3−

a. 0 b. 8 c. 6

Potassium hydroxide is used to precipitate each of the cations from their respective solution. Determine the minimum concentration of KOH required for precipitation to begin in each case. a. 0.015 M CaCl2 b. 0.0025 M Fe(NO3)2 c. 0.0018 M MgBr2 Ksp Mg(OH)2: 2.06 x 10-13 Ca(OH)2: 4.68 x 10-6 Fe(OH)2: 4.87 x 10-17

a. 0.018 M b. 1.4 10-7 M c. 1.1 x 10-5 M

Calculate the molar solubility of barium fluoride in each liquid or solution. a. Pure water b. 0.10 M Ba(NO3)2 c. 0.15 M NaF Ksp BaF2: 2.45 x 10-5

a. 0.0183 M b. 0.00755 M c. 0.00109 M

Consider the titration of 35.0 mL sample of 0.175 M HBr with 0.200 M KOH. Determine each quantity. a. The initial pH b. The volume of added base required to reach the equivalence point c. The pH at 10.0 mL of added base d. The pH at the equivalence point e. The pH after adding 5.0 mL of base beyond the equivalence point.

a. 0.757 b. 30.6 mL c. 1.038 d. 7 e. 12.15

Use the given molar solubilities in pure water to calculate Ksp for each compound a. NiS; molar solubility = 3.27 x 10-11 M b. PbF2; molar solubility = 5.63 x 10-3 M c. MgF2; molar solubility = 2.65 x 10-4 M

a. 1.07 x 10-21 b. 7.14 x 10-7 c. 7.44 x 10-11

A 350.0 mL buffer solution is 0.150 M in HF and 0.150 M in NaF. a. What mass of NaOH can this buffer neutralize before the pH rises above 4.00? b. If the same volume of the buffer was 0.350 M in HF and 0.350 M NaF, what mass of NaOH could be handled before the pH rises above 4.00? Ka HF: 3.5 x 10-4

a. 1.2 g b. 2.7 g

Determine the pH of an HNO2 solution of each concentration. In which cases can you not make the simplifying assumption? a. 0.500 M b. 0.100 M c. 0.0100 Ka HNO2: 4.6 x 10-4

a. 1.82 b. 2.18 (assumption doesn't work) c. 2.72 (assumption doesn't work)

Calculate the pH of a buffer consisting of 0.25 M (CH3)2NH2Cl and 0.30 M (CH3)2NH: a. Initially b. After addition of 0.73 g of HCl to 1.0 L of the buffer Kb (CH3)2NH: 5.9 x 10-4

a. 10.85 b. 10.80

Consider the titration of 25.0 mL sample of 0.175 M CH3NH2 with 0.150 M HBr. Determine each quantity. a. The initial pH b. The volume of added acid required to reach the equivalence point c. The pH at 5.0 mL of added acid d. The pH at one-half of the equivalence point e. The pH at the equivalence point f. The pH after adding 5.0 mL of acid beyond the equivalence point. Kb CH3NH2: 4.4 x 10-4

a. 11.94 b. 29.2 mL c. 11.33 d. 10.64 e. 5.87 f. 1.90

Consider the titration of 25.0 mL sample of 0.115 M RbOH with 0.100 M HCl. Determine each quantity. a. The initial pH b. The volume of added acid required to reach the equivalence point c. The pH at 5.0 mL of added acid d. The pH at the equivalence point e. The pH after adding 5.0 mL of acid beyond the equivalence point

a. 13.06 b. 28.8 mL c. 12.90 d. 7 e. 2.07

Find the pH at the equivalence point(s) and the volume (mL) of 0.0588 M KOH needed to reach the point(s) in titrations of: a. 23.4 mL of 0.0390 M HNO2 b. 17.3 mL of 0.130 M H2CO3 (two equivalence points)

a. 15.5 mL, 7.76 b. 38.2 (76.4 mL), 9.48, 11.35

Consider the titration of 20.0 mL sample of 0.105 M HC2H3O2 with 0.125 M KOH. Determine each quantity. a. The initial pH b. The volume of added base required to reach the equivalence point c. The pH at 5.0 mL of added base d. The pH at one-half of the equivalence point e. The pH at the equivalence point f. The pH after adding 5.0 mL of base beyond the equivalence point.

a. 2.86 b. 16.8 mL c. 4.37 d. 4.74 e. 8.75 f. 12.17

Calculate the pH of a buffer consisting of 0.50 M HF and 0.45 M KF: a. Initially b. After the addition of 0.40 g of NaOH to 1.0 L of the buffer Ka HF: 3.5 x 10-4

a. 3.12 b. 3.14

What is the coordination number of each metal: a. [Zn(NH3)4]SO4 b. [Cr(NH3)5Cl]Cl2 c. Na3[Ag(S2O3)2] d. [Cr(H2O)6]2(SO4)3 e. Ba[FeBr4]2 f. [Pt(en)2]CO3

a. 4 b. 6 c. 2 d. 6 e. 4 f. 4

A buffer is prepared by mixing 204 mL of 0.452 M HCl and 0.500 L of 0.400 M sodium acetate. a. What is the pH? b. How many grams of KOH must be added to 0.500 L of the buffer to change the pH by 0.15 units?

a. 4.81 b. 0.66 g KOH

Two 20.0 mL samples, one 0.200 M KOH and the other 0.200 M CH3NH2 are titrated with 0.100 M HI. a. What is the volume of added acid at the equivalence point for each titration? b. Is the pH at the equivalence point for each titration acidic, basic, or neutral? c. Which titration has the lower initial pH? d. Sketch each titration curve.

a. 40.0 mL HI for both b. KOH: neutral; CH3NH2 acidic c. CH3NH2

Find the pH at the equivalence point(s) and the volume (mL) of 0.0372 M NaOH needed to reach the point(s) in titrations of: a. 42.2 mL of 0.0520 M CH3COOH b. 28.9 mL of 0.0850 M H2SO3 (two equivalence points) Ka CH3COOH: 1.8 x10-5 H2SO3: Ka1 = 1.6 x 10-2 Ka2 = 6.4 x 10-8

a. 59.0 mL, 8.54 b. 66.0 mL, 132.0 mL, 7.10, 9.69

Use the Ksp values in table above to calculate the molar solubilities of each compound in pure water. a. AgBr b. Mg(OH)2 c. CaF2 Ksps AgBr: 5.35 x 10-13 Mg(OH)2: 2.06 x 10-13 CaF2: 1.46 x 10-10

a. 7.31 x 10-7 M b. 3.72 x 10-5 M c. 3.32 x 10-4 M

Blood is buffered by carbonic acid and the bicarbonate ion. Normal blood plasma is 0.024 M in HCO3 1- and 0.0012 M H2CO3 (pKa1 for H2CO3 at body temperature is 6.1). a. What is the pH of blood plasma? b. If the volume of blood in a normal adult is 5.0 L, what mass of HCl can be neutralized by the buffering system in blood before the pH falls below 7.0 (which would result in death)? c. Given the volume from part (b), what mass of NaOH can be neutralized before the pH rises above 7.8?

a. 7.4 b. 0.3 g c. 0.14 g

A solution is 0.085 M in Pb2+ and 0.025 M in Ag+. a. If selective precipitation is to be achieved using NaCl, what minimum concentration of NaCl do you need to begin to precipitate the ion that precipitates first? b. What is the concentration of each ion left in solution at the point where the second ion begins to precipitate? Ksp PbCl2: 1.17 x 10-5 AgCl: 1.77 x 10-10

a. AgCl precipitates first; [NaCl] = 7.1 x 10-9 M b. [Ag+] = 1.5 x 10-8 M and Pb2+ is 0.085 M

A solution is 0.010 M in Ba2+ and 0.020 M in Ca2+. a. If sodium sulfate is used to selectively precipitate one of the cations while leaving the other cation in solution, which cation will precipitate first? What minimum concentration of sodium sulfate will trigger the precipitation of the cation that precipitates first? b. What is the remaining concentration of the cation that precipitates first when the other cation begins to precipitate? Ksp BaSO4: 1.07 x 10-10 CaSO4: 7.10 x 10-5

a. Ba2+; 1.1 x 10-8 M b. 3.0 x 10-8 M

Use MO diagrams to place C2 −, C2, and C2 + in order of (a) increasing bond energy; (b) increasing bond length.

a. Bond energy increases as bond order increases: C2 + < C2 < C2 - b. Bond length decreases as bond energy increases, so the order of increasing bond length will be opposite that of increasing bond energy. Increasing bond length: C2 - < C2 < C2 +

Use MO diagrams and the bond orders you obtain from them to answer: (a) Is O2 − stable? (b) Is O2 − paramagnetic? (c) What is the outer (valence) electron configuration of O2 −?

a. Bond order = 3/2 = 1.5. O2 - is stable. b. O2 - is paramagnetic with an unpaired electron in the π*2p MO. c. (σ 2s)2(σ*2s)2(σ2p)2(π2p)2(π2p)2(π*2p)2(π*2p)

A solution is 0.022 M in Fe2+ and 0.014 M in Mg2+. a. If potassium carbonate is used to selectively precipitate one of the cations while leaving the other cation in solution, which cation will precipitate first? What minimum concentration of K2CO3 will trigger the precipitation of the cation that precipitates first? b. What is the remaining concentration of the cation that precipitates first, when the other cation begins to precipitate? Ksp FeCO3: 3.07 x 10-11 MgCO3: 6.82 x 10-6

a. Fe2+ will precipitate first, [CO32-] = 1.4 x 10-9 M b. [Fe2+] = 6.3 x 10-8 M

A 50.0-mL volume of 0.50 M Fe(NO3)3 is mixed with 125 mL of 0.25 M Cd(NO3)2. a. If aqueous NaOH is added, which ion precipitates first? b. Calculate the [OH-] that will accomplish the separation. Ksp Fe(OH)3: 1.6 x 10-39 Cd(OH)2: 7.2 x 10-15

a. Fe3+ will precipitate first b. [OH-] = 2.0 x 10-7 M

Choose the stronger acid in each of the following pairs: a. H2SeO3 or H2SeO4 b. H3PO4 or H3AsO4 c. H2S or H2Te d. HBr or H2Se e. H2Se or H3As f. H3AsO4 or H2SeO4 g. HNO3 or HNO2

a. H2SeO4 b. H3PO4 c. H2Te d. HBr e. H2Se f. H2SeO4 g. HNO3

Which are Lewis acids and which are Lewis bases? a. Cu2+ b. Cl− c. SnCl2 d. OF2

a. Lewis acid b. Lewis base c. Lewis acid d. Lewis base

Classify the following as Arrhenius, Brønsted-Lowry, or Lewis acid-base reactions. A reaction may fit all, two, one, or none of the categories: a. Ag+ + 2NH3 ⇄ Ag(NH3)2+ b. H2SO4 + NH3⇄ HSO4− + NH4+ c. 2HCl ⇄ H2 + Cl2 d. AlCl3 + Cl− ⇄ AlCl4

a. Lewis acid-base reaction b. Bronsted-Lowry reaction, Lewis acid-base reaction c. none d. Lewis acid-base reaction

Identify the Lewis acid and Lewis base in each equation: a. Na+ + 6H2O ⇌ Na(H2O)6+ b. CO2 + H2O ⇌ H2CO3 c. F− + BF3 ⇌ BF4 −

a. Na+ (Lewis acid) H2O (Lewis base) b. CO2 (Lewis acd) H2O (Lewis base) c. F− (Lewis base) BF3 (Lewis acid)

Which of these ions cannot form both high- and low-spin octahedral complexes: a. Ti3+ b. Co2+ c. Fe2+ d. Cu2+

a. No b. Yes c. Yes d. No

Use MO diagrams and the bond orders you obtain from them to answer: (a) Is Be2 + stable? (b) Is Be2 + diamagnetic? (c) What is the outer (valence) electron configuration of Be2 +?

a. With a bond order of ½ the Be2 + ion will be stable. b. No, the ion has one unpaired electron in the σ*2s MO, so it is paramagnetic c. (σ2s)2 (σ*2s)1

Which of these ligands can participate in linkage isomerism: a. NO2 − b. SO2 c. NO3 −? Explain with Lewis structures

a. Yes b. Yes c. No

Give the electron configuration and the number of unpaired electrons for: a. Sc3+ b. Cu2+ c. Fe3+ d. Nb3+

a. [Ar]; 0 b. [Ar] 3d(9); 1 c. [Ar]3d(5); 5 d. [Kr] 4d(2); 2

Write the electron configuration for: a. Os b. Co c. Ag

a. [Xe] 6s(2)4f(14)5d(6) b. [Ar]4s(2)3d(7) c. [Kr]5s(1)4d(10)

Give formulas corresponding to the following names: a. Tetraamminezinc sulfate b. Pentaamminechlorochromium(III) chloride c. Sodium di(thiosulfate)argentate(I) d. Hexaaquachromium(III) sulfate e. Barium tetrabromoferrate(III) f. Bis(ethylenediamine)platinum(II) carbonate

a. [Zn(NH3)4]SO4 b. [Cr(NH3)5Cl]Cl2 c. Na3[Ag(S2O3)2] d. [Cr(H2O)6]2(SO4)3 e. Ba[FeBr4]2 f. [Pt(en)2]CO3

For any of the following that can exist as isomers, state the type of isomerism and draw the structures: a. [Co(NH3)5Cl]Br2 b. [Pt(CH3NH2)3Cl]Br c. [Fe(H2O)4(NH3)2] 2+

a. coordination isomers b. coordination isomers c. geometric isomers

According to valence bond theory, what set of orbitals is used by a Period 4 metal ion in forming a. a square planar complex; b. a tetrahedral complex?

a. dsp2 b. sp3

For any of the following that can exist as isomers, state the type of isomerism and draw the structures: a. [Pt(CH3NH2)2Br2] b. [Pt(NH3)2FCl] c. [Pt(H2O)(NH3)FCl]

a. geometric isomers b. geometric isomers c. geometric isomers

For any of the following that can exist as isomers, state the type of isomerism and draw the structures: a. [PtCl2Br2] 2− b. [Cr(NH3)5(NO2)] 2+ c. [Pt(NH3)4I2] 2+

a. geometric isomers b. linkage isomers c. geometric isomers

Write balanced equations to show how the addition of HNO3 (aq) may affect the solubility of: a. Calcium fluoride b. Zinc sulfide c. Silver iodide d. Silver cyanide e. Copper(I) chloride f. Magnesium phosphate

a. increases: F1- + H3O+ ⇄ HF + H2O b. increases: S2- + H3O+ ⇄ HS1- + H2O c. no effect: I1- conjugate base of strong acid d. increases: CN1- + H3O+ ⇄ HCN + H2O e. no effect: Cl1- conjugate base of strong acid f. increases: PO4 3-- + H3O+ ⇄ HPO4 2- + H2O

Explain with equations and calculations, when necessary, whether an aqueous solution of each of these salts is acidic, basic, or neutral: a. KBr b. NH4I c. KCN d. Na2CO3 e. CaCl2 f. Cu(NO3)2 g. SrBr2 h. Ba(CH3COO)2 i. (CH3)2NHBr Ka H2SO3: Ka1 = 1.4 x 10-2 Ka2 = 6.5 x 10-8 H2S: Ka1 = 9 x 10-8 Ka2 = 1 x 10-17 HCN: 6.2 x 10-10 HCOOH: 1.8 x 10-4 Kb (CH3CH2)3N: 8.6 x 10-4 NH3: 1.76 x 10-5

a. neutral (both are conjugates of strong acid and base b. acidic: NH4 + + H2O à NH3 + H3O + c. basic: CN - + H2O à HCN + OH 1- d. basic: CO3 2- + H2O à HCO3 1- + OH 1- e. neutral f. acidic: Cu(H2O)6 2+ + H2O à Cu(H2O)5OH 1+ + H3O + g. neutral h. basic: CH3COO 1- + H2O à HCH3COO + OH 1- i. acidic: (CH3)2NH2 + + H2O à (CH3)2NH + H3O +

Draw orbital-energy splitting diagrams and use the spectrochemical series to show the orbital occupancy for each of the following (assuming that H2O is a weak-field ligand): a. [Cr(H2O)6] 3+ b. [Cu(H2O)4] 2+ c. [FeF6] 3− d. [MoCl6] 3− e. [Ni(H2O)6] 2+ f. [Ni(CN)4] 2−

a. octahedral splitting, 3 unpaired electrons in t2g b. square planar splitting, all paired except for 1 electron in dx2-y2 c. octahedral splitting, 3 unpaired electrons in t2g, 2 unpaired in eg d. octahedral splitting, 3 unpaired electrons in t2g e. octahedral splitting, 6 paired electrons in t2g, 2 unpaired in eg f. square planar splitting, all paired except last dx2-y2

Classify each as a strong or weak, acid or base: a. H3AsO4 b. Sr(OH)2 c. HIO d. HClO4 e. RbOH f. HBr g. H2Te h. HClO

a. weak acid b. strong base c. weak acid d. strong acid e. strong base f. strong acid g. weak acid h. weak acid

Write balanced equations and Kb expressions for these Brønsted-Lowry bases in water: a. Pyridine, C5H5N b. CO32−

answer on ws3

Write the Ka expression for each of the following in water: a. HCN b. HCO3 1- c. HCOOH

answer on ws3

Which of the following are Arrhenius bases? a. H3AsO4 b. Ba(OH)2 c. HClO d. KOH

b and d

Identify the Lewis bases in the following reaction: Cu(H2O)4 2+ (aq) + 4 CN 1- (aq) ⇄ Cu(CN)4 2- (aq) + 4 H2O (l) Given that Kc > 1 for the reaction above, which Lewis base is stronger?

forward direction: CN1- reverse direction: H2O CN1- is stronger base