Chapter 13
40. Methanol can be prepared from carbon monoxide and hydrogen at high temperature and pressure in the presence of a suitable catalyst. (a) Write the expression for the equilibrium constant (Kc) for the reversible reaction2H2(g)+CO(g)⇌CH3OH(g)ΔH=−90.2kJ (b) What will happen to the concentrations of H2, CO, and CH3OH at equilibrium if more H2 is added? (c) What will happen to the concentrations of H2, CO, and CH3OH at equilibrium if CO is removed? (d) What will happen to the concentrations of H2, CO, and CH3OH at equilibrium if CH3OH is added? (e) What will happen to the concentrations of H2, CO, and CH3OH at equilibrium if the temperature of the system is increased?
(a) Kc=[CH3OH]/[H2]2[CO]; (b) [H2] increases, [CO] decreases, [CH3OH] increases; (c), [H2] increases, [CO] decreases, [CH3OH] decreases; (d), [H2] increases, [CO] increases, [CH3OH] increases; (e), [H2] increases, [CO] increases, [CH3OH] decreases
43. Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign. (a) acidity: NaHSO3, NaHSeO3, NaHSO4 (b) basicity: BrO2−,BrO2−, ClO2−,ClO2−, IO2−IO2− (c) acidity: HOCl, HOBr, HOI
(a) NaHSeO3 < NaHSO3 < NaHSO4; in polyoxy acids, the more electronegative central element—S, in this case—forms the stronger acid. The larger number of oxygen atoms on the central atom (giving it a higher oxidation state) also creates a greater release of hydrogen atoms, resulting in a stronger acid. As a salt, the acidity increases in the same manner. (b) ClO2−<BrO2−<IO2−;ClO2−<BrO2−<IO2−; the basicity of the anions in a series of acids will be the opposite of the acidity in their oxyacids. The acidity increases as the electronegativity of the central atom increases. Cl is more electronegative than Br, and I is the least electronegative of the three. (c) HOI < HOBr < HOCl; in a series of the same form of oxyacids, the acidity increases as the electronegativity of the central atom increases. Cl is more electronegative than Br, and I is the least electronegative of the three.
Write the mathematical expression for the reaction quotient, Qc, for each of the following reactions: (a) CH4(g)+Cl2(g)⇌CH3Cl(g)+HCl(g) (b) N2(g)+O2(g)⇌2NO(g) (h) CuSO4⋅5H2O(s)⇌CuSO4(s)+5H2O(g)
(a) Qc=[CH3Cl][HCl]/[CH4][Cl2]; b) Qc=[NO]2/[N2][O2] (h) Qc = [H2O]5
How will an increase in temperature affect each of the following equilibria? How will a decrease in the volume of the reaction vessel affect each? (a) 2NH3(g)⇌N2(g)+3H2(g)ΔH=92kJ (d) CaO(s)+CO2(g)⇌CaCO3(s)ΔH=−176kJ
(a) T increase = shift right, V decrease = shift left; (d) T increase = shift left, V decrease = shift right.
44. Ammonia is a weak base that reacts with water according to this equation:NH3(aq)+H2O(l)⇌NH4+(aq)+OH−(aq) Will any of the following increase the percent of ammonia that is converted to the ammonium ion in water? (a) Addition of NaOH (b) Addition of HCl (c) Addition of NH4Cl
(b)
When x is so small and less than 5% of the initial it is safe to assume it to be?
0
1. Write equations that show NH3 as both a conjugate acid and a conjugate base.
1. One example for NH3 as a conjugate acid: NH2−+H+⟶NH3;NH2−+H+⟶NH3; as a conjugate base: NH4+(aq)+OH−(aq)⟶NH3(aq)+H2O(l)
Calculate the PH of these strong acid/bases: 1.) .10M HCl 2.) 12M HCl 3.).20M Ca(OH)2
1.) 1.00 pH 2.) -1.08 pH 3.) 13.60 pH
Consider a buffer made by adding 42.3 g of (CH₃)₂NH₂I to 250.0 mL of 1.42 M (CH₃)₂NH (Kb = 5.4 x 10⁻⁴) 1.) What is the pH of this buffer? 2.)What is the pH of the buffer after 0.300 mol of H⁺ have been added? 3.)What is the pH of the buffer after 0.120 mol of OH⁻ have been added?
1.) 10.93 2.) 9.73 3.) 11.31
What is roughly the pH for each of the following titrations (greater than 7, 7, less than 7) 1.) Strong Acid-Strong Base 2.) Strong Acid- Weak Base 3.) Strong Base- Weak Acid
1.) 7 2.) Less than 7 3.) Greater than 7
To manipulate K: 1.) Flipped Equation ? 2.) Change coefficients by n? 3.) Add Equations?
1.) Inverse K 2.) Change Exponents by n (K)^n 3.) Multiply K Values
In Exothermic Reactions: 1.) Increasing Temperature causes a shift: 2.) Decreasing temperature causes a shift:
1.) Left 2.) Right
Predict whether the solution will be acidic or basic: 1.) NaCl 2.) Na2CO3 3.) NH4Cl 4.) NaF
1.) Neutral 2.) Basic 3.) Acidic 4.) Basic
In endothermic reactions: 1.) Increasing temperature causes a shift: 2.) Decreasing temperature causes a shift:
1.) Right 2.) Left
1.) What is the pH of a solution where 0.10M HCl is added to a 1.0L buffer? -Contains 0.90 M NaH2PO4 -Contains 0.80 M Na2HPO4 2- -Ka2: 6.3 * 10^-8 2.) What is the pH if 0.90 mol of NaOH is added to the buffer?
1.) pH= 7.04 THIS IS A BUFFER (first calculate regular ph, then draw chemical equation with IRN Table, and then use pH=pka + log (a/b) 2.) pH=10.72 This is a weak base, Firstly set up IRN table then set up ice table. Then find the kb from the given Ka. Then plug this in to find pOH and use to find pH.
In neutral water H30+ and OH- are both equal to what molarity?
1.0x10^7M
The constant KW for the autoronization of water is.......
1.4 x 10^14M
Consider the following reactions: A ⇌ B, K₁=3.76 A ⇌ C, K₂=2.00 What is K for the reaction C ⇌ B?
1.88
What is the pH of a 0.750 M solution of NaCN (Ka of HCN is 4.9 × 10⁻¹⁰)?
11.66
19. Calculate the pH and the pOH of each of the following solutions at 25 °C for which the substances ionize completely: (a) 0.000259 M HClO4 (b) 0.21 M NaOH (c) 0.000071 M Ba(OH)2 (d) 2.5 M KOH
19. (a) pH = 3.587; pOH = 10.413; (b) pOH = 0.68; pH = 13.32; (c) pOH = 3.85; pH = 10.15; (d) pOH = −0.40; pH = 14.4
A diprotic acid, H₂A, has Ka1 = 3.4 × 10⁻⁴ and Ka2 = 6.7 × 10⁻⁹. What is the pH of a 0.18 M solution of H₂A?
2.11
The equilibrium constant for the reaction 2 HF (g) ⇌ H₂ (g) + F₂ (g) is 0.210 at a particular temperature. What is the equilibrium constant for the equation ½ H₂ (g) + ½ F₂ (g) ⇌ HF (g)?
2.18
What is the pH of a 0.420 M solution of C₅H₅NHBr (Kb of C₅H₅N is 1.7 × 10⁻⁹)?
2.80
In the titration of 100.0 mL of 0.4000 M HONH₂ with 0.2000 M HBr, how many mL of HBr are required to reach the equivalence point?
200mL
21. What are the hydronium and hydroxide ion concentrations in a solution whose pH is 6.52?
21. [H3O+] = 3.0 ×× 10−7 M; [OH−] = 3.3 ×× 10−8 M
25. The hydroxide ion concentration in household ammonia is 3.2 ×× 10−3 M at 25 °C. What is the concentration of hydronium ions in the solution?
25. [OH−] = 3.1 ×× 10−12 M
A reaction vessel at room temperature contains an equilibrium mixture of SO2 (0.0018 atm), O2 (0.0032 atm), and SO3 (0.0166 atm). What is Kp?
27,000
3. Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid: (a) H3O+H3O+ (b) HCl (c) NH3 (d) CH3CO2H (e) NH4+NH4+ (f) HSO4−
3. (a) H3O+(aq)⟶H+(aq)+H2O(l) (b) HCl(aq)⟶H+(aq)+Cl−(aq) (c) NH3(aq)⟶H+(aq)+NH2−(aq) (d) CH3CO2H(aq)⟶H+(aq)+CH3CO2−(aq) (e) NH4+(aq)⟶H+(aq)+NH3(aq) (f) HSO4−(aq)⟶H+(aq)+SO42−(aq)
33. Gastric juice, the digestive fluid produced in the stomach, contains hydrochloric acid, HCl. Milk of Magnesia, a suspension of solid Mg(OH)2 in an aqueous medium, is sometimes used to neutralize excess stomach acid. Write a complete balanced equation for the neutralization reaction, and identify the conjugate acid-base pairs.
33. Mg(OH)2(s)BB+2HCl(aq)BA⟶Mg2CB+(aq)+2Cl−(aq)+2H2O(l)CA
An unknown weak base with a concentration of 0.170 M has a pH of 9.42. What is the Kb of this base?
4.06 * 10-9
What is the pH of a solution composed of .25M HC2H3O2 and .35M NaC2H3O2? the Ka of HC2H3O2 = 1.8 * 10^-5
4.88
41. Predict which compound in each of the following pairs of compounds is more acidic and explain your reasoning for each. (a) HSO4−HSO4− or HSeO4−HSeO4− (b) NH3 or H2O (c) PH3 or HI (d) NH3 or PH3 (e) H2S or HBr
41. (a) HSO4−;HSO4−; higher electronegativity of the central ion. (b) H2O; NH3 is a base and water is neutral, or decide on the basis of Ka values. (c) HI; PH3 is weaker than HCl; HCl is weaker than HI. Thus, PH3 is weaker than HI. (d) PH3; in binary compounds of hydrogen with nonmetals, the acidity increases for the element lower in a group. (e) HBr; in a period, the acidity increases from left to right; in a group, it increases from top to bottom. Br is to the left and below S, so HBr is the stronger acid.
66. The pH of a 0.15-M solution of HSO4−HSO4− is 1.43. Determine Ka for HSO4−HSO4− from these data.
66. Ka=1.2×10−2
68. The pH of a solution of household ammonia, a 0.950 M solution of NH3, is 11.612. Determine Kb for NH3 from these data.
68. Kb=1.77×10−5
7. What is the conjugate acid of each of the following? What is the conjugate base of each? (a) OH− (b) H2O (c) HCO3−HCO3− (d) NH3
7. (a) H2O, O2−; (b) H3O+, OH−; (c) H2CO3, CO32−;CO32−; (d) NH4+,NH4+, NH2−;NH2−; (e) H2SO4, SO42−;SO42−; (f) H3O2+,H3O2+, HO2−;HO2−; (g) H2S; S2−; (h) H6N22+,H6N22+, H4N2
Calculate the pH of a buffer solution that is .90M NaH2PO4 and .80M Na2PO4. ka2= 6.3*10^-8
7.15
What is the pH of a buffer made from 0.350 mol of HBrO (Ka = 2.5 × 10⁻⁹) and 0.120 mol of KBrO in 2.0 L of solution?
8.14
Calculate the pH of 0.35M NaNo3 ka of HNO2: 7.1 * 10^-4
8.34
If the base=acid: pH is _____ to pka If the base>acid: pH is _____ to pka If the base<acid: pH is _____ to pka
=, >, <
Method of analyzing a solution of unknown concentration by reacting it with a solution of known concentration.
Acid-Base Titration
If [H3O+] of a solution is greater than [OH-], the solution is
Acidic
In CATIONS, if it is a conjugate acid of a weak base the solution will be....
Acidic
In ANIONS if it is conjugate bases of a weak acid that solution will be....
Basic
if the [OH-] of a solution is greater than [H+], the solution is
Basic
Do catalysts affect equilibrium?
Catalysts do not change the position of the equilibrium; the reaction just arrives at the equilibrium faster
A shift right means that what rate will be temporarily faster?
Forward Rate
Rank these based on pH acidity: 1.) pOH= 3.64 2.) H30+ = 1.9 * 10^ -4 3.) OH-= 7.3 * 10^-7
Highest acidity: H30+= 3.72 pH , OH-= 7.86 pH, pOH=10.36 pH
When you have an acid and water the result is........
Hydronium and Conjugate Base
When you have a weak base and water the result is......
Hydroxide and Conjugate Acid
Which of the following statements are always true for a reaction at equilibrium? I. The rate of the forward and reverse reactions are equal. II. The concentrations of the reactants and the products remain constant. III. The amount of reactants is equal to the amount of products.
I and II
Calculate K given that in CO + 2H2 = CH3OH 1.) Initially CO= 0.50M 2.) Initially H2= 1.00 M 3.) At Equilibrium CO= 0.15M
K=26
58. A sample of ammonium chloride was heated in a closed container. NH4Cl(s)⇌NH3(g)+HCl(g)NH4Cl(s)⇌NH3(g)+HCl(g) At equilibrium, the pressure of NH3(g) was found to be 1.75 atm. What is the value of the equilibrium constant KP for the decomposition at this temperature?
KP = 3.06
In polyprotic acids what proton is the easiest to donate: Ka1, Ka2, Ka3....
Ka1
52.) What is the value of the equilibrium constant at 500 °C for the formation of NH3 according to the following equation? N2(g)+3H2(g)⇌2NH3(g)N2(g)+3H2(g)⇌2NH3(g) An equilibrium mixture of NH3(g), H2(g), and N2(g) at 500 °C was found to contain 1.35 M H2, 1.15 M N2, and 4.12 × 10−1 M NH3.
Kc = 6.00 × 10−2
(a) Write the expression for the equilibrium constant for the reaction represented by the equation AgCl(s)⇌Ag+(aq)+Cl−(aq). Is Kc > 1, < 1, or ≈ 1? Explain your answer.
Kc = [Ag+][Cl−], <1
Write a Kc Expression for: 2NO(g) + O2(g) = 2NO2(g)
Kc= ((NO2)2 / (NO)2 (O2))
Consider the equilibrium system described by the chemical reaction below. If the partial pressures at equilibrium of NO, Cl₂, and NOCl are 0.095 atm, 0.171 atm, and 0.28 atm, respectively, in a reaction vessel of 7.00 L at 500 K, what is the value of Kp for this reaction? 2 NO(g) + Cl₂(g) ⇌ 2 NOCl(g)
Kp=51
Removing reactant causes a shift:
Left
When you add a product to a reversible chemical reaction, the reaction is always pushed in the direction:
Left
19. The following reaction has KP = 4.50 × 10−5 at 720 K.N2(g)+3H2(g)⇌2NH3(g If a reaction vessel is filled with each gas to the partial pressures listed, in which direction will it shift to reach equilibrium? P(NH3) = 93 atm, P(N2) = 48 atm, and P(H2) = 52 atm
Left, reactants
In ANIONS if it is conjugate bases of a strong acid that solution will be....
Neutral
In CATIONS if it is conjugate acids of a strong base the solution will be.....
Neutral
Does adding inert (unreactive gas) cause a change?
No, no change
If Q>K
Not at equilibrium: reaction shifts left
If Q < K
Not at equilibrium: reaction shifts right
Find the equilibrium concentrations given that in PCl3+ Cl2 = PCl5 1.) Kp= 1.245 2.) PCl3= 13.2 atm 3.) PCl2= 32.8 atm 4.) PCl5= 217 atm
PCl3= 6.8 atm PCl2= 26.4 atm PCl5= 223.4 atm
If K is bigger than 1
Products are favored (R leaning)
dynamic equilibrium means
Rate forward is equal to rate reverse
If K is smaller than 1
Reactants are favored (L leaning)
A shift left means that what rate will be temporarily faster?
Reverse Rate
Removing product causes a shift:
Right
When you add reactant to a reversible chemical reaction, the reaction is always pushed in the direction:
Right
What do not appear in K expressions : Solid, Liquid, Aqueous, Gas?
Solid and Liquid do not appear
The Equilibrium Constant tells us about what?
The relative amounts of products and reactants for a specific reaction at a specific temperature
If you decrease volume if causes a shift:
Toward side will less gas molecules
If you increase volume it causes a shift:
Toward side with more gas molecules
64. Analysis of the gases in a sealed reaction vessel containing NH3, N2, and H2 at equilibrium at 400 °C established the concentration of N2 to be 1.2 M and the concentration of H2 to be 0.24 M. N2(g)+3H2(g)⇌2NH3(g)Kc=0.50at400°CN2(g)+3H2(g)⇌2NH3(g)Kc=0.50at400°C Calculate the equilibrium molar concentration of NH3.
[NH3] = 9.1 × 10−2 M
82. What is [OH−] in a solution of 0.125 M CH3NH2 and 0.130 M CH3NH3Cl? CH3NH2(aq)+H2O(l)⇌CH3NH3+(aq)+OH−(aq) Kb=4.4×10−4
[OH−] = 4.2 × 10−4 M
Calculate pH at each of the following points of the titration of 40.0 ml of .100M HC5H5O2 with 0.100M NaOH ka= 1.3 *10-5 a.) O.00ml NaOH added b.) 20.0ml of NaOH added c.)40.0ml NaOH added d.) 50.0ml NaOH added
a.) 2.94 b.) 4.89 c.) 8.79 d.) 12.05
Calculate the pH at the following points for the titration of 40.0 ml of .100MHCl with 0.100 NaOH a.) O.00ml base added b.) 100.0 ml base added c.) equivilance point d.) when .0500L NaOh is added?
a.) PH=1.000 b.) pH= 1.222 c.) pH= 7 d.) pH= 12.045
Amphateric compunds can act as both
acids and bases
The higher the Kb/Ka the stronger the
base/acid
Strong acids or bases cannot be in _______ because.
buffers; because they completely dissociate in water
In bronsted lowry acids and bases..... acids (donate/accept) protons and bases (donate/accept) protons
donate:accept
Nearly all salts are strong
electrolytes
At least one strong compound causes a full ______________ reaction
foreward
At the equivalence point for the titration of HCN with KOH, the pH is expected to be
greater than 7
Buffers are solutions that lessen the impact on __ from the addition of an acid or base
pH
Bronsted-Lowry involve the transfer of
protons
If Q=K
reaction is at equilibrium
There is always a combination of a weak base and a _____ of a conjugate ______
salt; acid
There is always a combination of a weak acid and a ______ of a conjugate _______
salt; base
The more electronegative a compound the _________ the acid
stronger
the stronger the acid, the _____ the conjugate base
weaker
the stronger the base, the ______ the conjugate acid
weaker
In terms of bond strength, a large radius means a __________ bond and a __________ acid
weaker, stronger
What is Le Chatelier's principle?
when a chemical system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance
