CHEM 1212 Exam 2

1.What is the percent ionization of ammonia at this concentration? 2. ^^^^^

% ionization = 1.50 % .679%

What is the pH of a 0.28 M solution of ascorbic acid (Vitamin C)? (The values for Ka1 and Ka2 for ascorbic acid are 8.0×10−5 and 1.6×10−12, respectively.) 2.04 2.32 2.82 4.65 6.17

(8.0*10^-5 X .28)^1/2 4.73* 10^-3 pH=-log(4.73*10^-3) =2.32

Which one of the following is a Bronsted-Lowry base? HNO2 CH3COOH HF (CH3)3N none of the above

(CH3)3N

Which one of the following is a Bronsted-Lowry acid? CH3)3NH+ HF CH3COOH HNO2 all of the above

(CH3)3NH+ HF CH3COOH HNO2 *all of the above*

Calculate the percent ionization of hydrazoic acid (HN3) in solutions of each of the following concentrations (Ka is given in Appendix D in the textbook).

(Ka / molarity given )^1/2

Codeine (C18H21NO3) is a weak organic base. A 5.0×10−3M solution of codeine has a pH of 9.95. 1.Calculate the value of Kb for this substance. 2.What is the pKb for this base?

1. Kb= 1.6×10−6 2.pKb= = 5.79

In a certain acidic solution at 25 ∘C, [H+] is 100 times greater than [OH −]. What is the value for [OH −] for the solution? 1.0×10−8 M 1.0×10−7 M 1.0×10−6 M 1.0×10−2 M 1.0×10−9 M

1.0×10−8 M

You add 10.0 grams of solid copper(II) phosphate, Cu3(PO4)2, to a beaker and then add 100.0 mL of water to the beaker at T = 298 K. The solid does not appear to dissolve. You wait a long time, with occasional stirring and eventually measure the equilibrium concentration of Cu2+(aq) in the water to be 5.01×10−8 M. What is the Ksp of copper(II) phosphate? You add 10.0 grams of solid copper(II) phosphate, , to a beaker and then add 100.0 of water to the beaker at = 298 . The solid does not appear to dissolve. You wait a long time, with occasional stirring and eventually measure the equilibrium concentration of in the water to be . What is the of copper(II) phosphate? 5.01×10−8 2.50×10−15 4.20×10−15 3.16×10−37 1.40×10−37

1.40×10−37

A 0.50 M solution of an acid HA has pH = 2.24. What is the value of Ka for the acid? A 0.50 solution of an acid has = 2.24. What is the value of for the acid? 1.2×10−2 6.6×10−5 1.7×10−12 5.8×10−3 3.3×10−5

10^-2.24 .0057=[H+] at equilibrium [H+]=[A-] so [HA]= .50M-.0057M=.49M Ka=[H+][A-] / [HA] Ka= .0057 * .0057 /.49 Ka=6.6 * 10^-5

What is the H+ concentration for an aqueous solution with pOH = 4.23 at 25 ∘C?

10^pOH= OH concentration [H+]= 10^-14/OH Concentration

Bronstead-Lowry Review Con't

A Brønsted-Lowry acid must have at least one removable (acidic) proton (H+) to donate. A Brønsted-Lowry base must have at least one nonbonding pair of electrons to accept a proton (H+).

Lewis Acid-Base Definition

A base is an electron pair donor. An acid is an electron pair acceptor.

Arrhenius base

A base solution contains an excess of OH- ions. Properties: Bitter or caustic taste Turn red litmus blue Slippery, soapy feeling Neutralize acids

Bronstead-Lowry Acids and bases

Ammonia is a weak base that forms the ammonium ion and the hydroxide ion in water Conjugate acid-base pairs: NH4+ - NH3 and H2O - OH-

REVIEW BRONSTEAD LOWRY ACIDS AND BASES

An acid is a proton donor. A base is a proton acceptor

REVIEW ARRHENIUS ACIDS AND BASES

An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions. A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions.

Arrhenius acid

An acid solution contains an excess of H+ ions. Sour taste Turn blue litmus red The ability to react with: 1.Metals to produce H2 gas 2.Bases to produce salt and water 3.Carbonates to produce carbon dioxide

n the forward reaction of this equilibrium, which substance acts as the Brønsted-Lowry base? H2S(aq)+CH3NH2(aq)⇋HS−(aq)+CH3NH+3(aq)

CH3NH2(aq)

The conjugate acid of CH3NH2 is ________. CH3NH2 CH3NH3+ CH3NH+ CH3NH2+ none of the above

CH3NH3+

The conjugate base of HCO−3 (or NH3) ________. H2CO3 . is NH2- CO3(2-) HCO−4 H3O+ OH−

CO3(2-)

Another metal phosphate is cobalt phosphate. It will behave similar to calcium phosphate in an acid solution. What is the net ionic equation including phases for CoPO4(s) dissolving in H3O+(aq)?

CoPO4(s)+H3O+(aq)⇌Co3+(aq)+HPO42−(aq)+H2O(l) or AlPO4(s)+H3O+(aq)⇌Al3+(aq)+HPO42−(aq)+H2O(l)

Ka

For a weak acid, the equation for its dissociation is HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) Since it is an equilibrium, there is an equilibrium constant related to it, called the acid-dissociation constant, Ka:Ka = [H3O+][A-] / [HA] The greater the value of Ka, the stronger is the acid

The conjugate acid of HSO4- is ________. HSO3+ SO42- H2SO4 H+ HSO4+

H2H2SO4

Write a chemical equation that illustrates the autoionization of water.

H2O(l)⇌H+(aq)+OH−(aq)

HSO4−(aq) + OH−(aq) ⇌ SO42−(aq) + H2O(l) which are the acids?

H2SO4 and H2O

Relative strength of acids and bases I. CH3COOH(aq) + HS−(aq) ⇌ CH3COO−(aq) + H2S(aq) II. F−(aq) + NH4+(aq) ⇌ HF(aq) +NH3(aq) III. H2CO3(aq) + Cl−(aq) ⇌ HCO3−(aq) + HCl(aq)

III>II>I

For the generic equilibrium HA(aq) ⇌ H+(aq) + A−(aq), which of these statements is true? If you add the soluble salt KA to a solution of HA that is at equilibrium, the pH would increase. If you add the soluble salt KA to a solution of HA that is at equilibrium, the concentration of A− would decrease. The equilibrium constant for this reaction changes as the pH changes. If you add the soluble salt KA to a solution of HA that is at equilibrium, the concentration of HA would decrease.

If you add the soluble salt KA to a solution of HA that is at equilibrium, the pH would increase.

Why does the addition of acid increase the solubility of calcium phosphate?

It decreases the phosphate ion concentration, forcing the equilibrium to the right.

If the molar solubility of CaF2 at 35 ∘C is 1.24×10−3mol/L, what is Ksp at this temperature? It is found that 1.1×10−2g of SrF2 dissolves per 100 mL of aqueous solution at 25 ∘C. Calculate the solubility product for SrF2. Express your answer using two significant figures. The Ksp of Ba(IO3)2 at 25 ∘C is 6.0×10−10. What is the molar solubility of Ba(IO3)2?

Ksp = 7.63×10−9 Ksp = 2.7×10−9 5.3×10−4 mol/L

Write the expression for the ion-product constant for water, Kw.

Kw=[H+][OH−]

What is the conjugate acid of NH3? NH3+ NH4OH NH4+ NH3 NH2+

NH4+

Why is pH at the equivalence point larger than 7 when you titrate a weak acid with a strong base?

The conjugate base that is formed at the equivalence point reacts with water.

You have an aqueous solution of chromium(III) nitrate that you titrate with an aqueous solution of sodium hydroxide. After a certain amount of titrant has been added, you observe a precipitate forming. You add more sodium hydroxide solution and the precipitate dissolves, leaving a solution again. What has happened?

The precipitate was chromium hydroxide, which then reacted with more hydroxide to produce a soluble complex ion, Cr(OH)4−.

Conjugate Acids and Bases

The term conjugate means "joined together as a pair." Reactions between acids and bases always yield their conjugate bases and acids.

What Is Different about Water?

Water can act as a Brønsted-Lowry base and accept a proton (H+) from an acid It can also donate a proton and act as an acid This makes water amphiprotic.

Calculate the concentration of H+ ions in a 0.010 M aqueous solution of sulfuric acid. Express your answer to three decimal places and include the appropriate units.

[H+] = 0.015 M

f a solution is described as basic, which of the following is true: If a solution is described as basic, which of the following is true: [H+]>[OH−] [H+]<[OH−] [H+]=[OH−]

[H+]<[OH−]

Calculate the concentration of HSO4− ions in a 0.010 M aqueous solution of sulfuric acid.

[HSO4−] = 5.5×10−3 M

Calculate the concentration of SO42− ions in a 0.010 M aqueous solution of sulfuric acid.

[SO42−] = 0.0045 M

Bronsted-Lowry base

a molecule or ion that is a proton acceptor

Bronsted-Lowry acid

a molecule or ion that is a proton donor (H+ donor)

A solution at 25 ∘C has pOH = 10.53. Which of the following statements is or are true? The solution is acidic. The pH of the solution is 14.00 - 10.53. For this solution, [OH−] = 10−10.53 M. Only one of the statements is true. Statements I and II are true. Statements I and III are true. Statements II and III are true. All three statements are true.

all are correct

According to the Arrhenius concept, an acid is a substance that ________. reacts with the solvent to form the cation formed by autoionization of that solvent causes an increase in the concentration of H+ in aqueous solutions is capable of donating one or more H+ can accept a pair of electrons to form a coordinate covalent bond tastes bitter

causes an increase in the concentration of H+ in aqueous solutions

For oxyacids of the same general structure but differing electronegativities of the central atoms, acid strength decreases with increasing electronegativity of the central atom.

false

The strongest acid known is HF because fluorine is the most electronegative element.

false

Would a 1.0×10−8 M solution of HCl have pH < 7, pH=7, or pH > 7?

pH < 7

1. What is the pH of a 8.00×10−2 M ammonia solution? 2. ph of a .39 M ammonia solution

pH = 11.08 pH= 11.42

A 35.0-mL sample of 0.150 M acetic acid (CH3COOH) is titrated with 0.150 MNaOH solution. Calculate the pH after the following volumes of base have been added. 0ml 17.5ml 34.5ml 35ml 35.5ml 50ml

pH = 2.78 pH= 4.74 pH=6.58 pH=8.81 pH=11.03 pH=12.42

Calculate the pH of a buffer that is 0.13 M in lactic acid and 0.11 M in sodium lactate.

pH = 3.78

Calculate the pH of a buffer formed by mixing 85 mL of 0.12 M lactic acid with 95 mL of 0.16 M sodium lactate.

pH = 4.03

What is the pH of the buffer after the addition of 3×10−2 mol of HNO3?

pH = 4.6

What is the pH of this buffer? Express your answer using two significant figures.

pH = 4.8

What is the pH of the buffer after the addition of 3×10−2 mol of KOH?

pH = 5.1

Consider the titration of 50.0 mL of 0.20 M NH3 (Kb=1.8×10−5) with 0.20 M HNO3. Calculate the pH after addition of 50.0 mL of the titrant at 25 ∘C.

pH = 5.13

What is the pH after 0.150 mol of HCl is added to the buffer from Part A? Assume no volume change on the addition of

pH = 5.818 pH = 5.911

What is the pH of a buffer prepared by adding 0.809 mol of the weak acid HA to 0.507 mol of NaA in 2.00 L of solution? The dissociation constant Ka of HA is 5.66×10−7. What is the pH of a buffer prepared by adding 0.405 mol of the weak acid HA to 0.406 mol of NaA in 2.00 L of solution? The dissociation constant Ka of HA is 5.66×10−7.

pH = 6.044 pH = 6.248

What is the pH after 0.195 mol of NaOH is added to the buffer from Part A? Assume no volume change on the addition of the base.

pH = 6.305 pH = 6.704

What is the pH at the equivalence point when 0.10 M HNO3 is used to titrate a volume of solution containing 0.30 g of KOH?

pH = 7

A 84.0 mL volume of 0.25 M HBr is titrated with 0.50 M KOH. Calculate the pH after addition of 42.0 mL of KOH at 25 ∘C.

pH = 7.00

hydrobromic acid (HBr) chlorous acid (HClO2) benzoic acid (C6H5COOH)

pH = 7.00 pH = 7.32 pH = 8.44

A 30.0-mL volume of 0.50 M CH3COOH (Ka=1.8×10−5) was titrated with 0.50 M NaOH. Calculate the pH after addition of 30.0 mL of NaOH at 25 ∘C.

pH =9.07

What is the pH at the equivalence point when 0.10 M HNO3 is used to titrate a volume of solution containing 0.30 g of KOH?

pH= . 7

what is the equation of for pH?

pH=-log[H+]

Calculate the concentration of an aqueous solution of Ca(OH)2 that has a pH of 11.64.

step 1: 14-11.64 2: 10^-2.36 3: .00437 4. .00437 /2

Acid strength in a series of H−A molecules increases with increasing size of A.

true