1) Which one of the following is a weak acid?

1. HBr
2. HI
3. H2SO3
4. HClO4
5. HNO3

2) Acid rain has a pH value of 3.0, whereas
normal rain has a pH value of 5.6.
Calculate the ratio of the hydronium ion in
acid rain to that in normal rain.

<Numerical Answer>

3) In a sample of pure water, only one of the
following statements is always true at all 
conditions of temperature and pressure.
Which one is always true?

1. [H3O+] = [OH-]
2. pH = 7.0
3. [OH-] = 1.0 * 10^-7 M
4. [H3O+] = 1.0 * 10^-7 M
5. pOH = 7.0

4) Hydroxylamine is a weak molecular base with
Kb = 6.6 * 10^-9.
What is the pH of a 0.0500 M solution of
hydroxylamine?

1. pH = 3.63
2. pH = 4.74
3. pH = 7.12
4. pH = 8.93
5. pH = 9.26
6. pH = 9.48
7. pH = 10.37

5) The [(CO2)-2] in 0.050 M H2CO3 solution 
(e.g. Pepsi) is?

1. 1.0 * 10^-7 M
2. 1.3 * 10^-13 M
3. 1.5 * 10^-4 M
4. 4.3 * 10^-7 M
5. 4.8 * 10^-11 M

Hello again, jwheel-ga!

I won’t bore you with my intro to these questions since you’ve already
seen it. I’ll just say that I hope that you and your daughter together
can learn a lot from these exercises, and I hope that you both will
use these answers to gain a better understanding of the concepts being
quizzed. I respect your willingness to assist your daughter at a
difficult time for her right now. Good luck on the assignment.

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Question 1: Answer 3. H2SO3 is a weak acid (Careful! H2SO4 is not!)

The Shodor Foundation’s Chemistry website
(http://www.shodor.org/unchem/basic/ab/) states on its acid-base page
that there are only 7 “strong acids”:

HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO3, HclO4

“Strong acid” means that the Hydrogen ion [H+] completely dissociates
from its conjugate anion.

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Question 2: Answer 400:1.

pH = -log[H3O+]

Acid rain has a pH of 3.0, so the [H3O+] is 10^-3.

Normal rain has a pH of 5.6, so the [H3O+] is 2.5 x 10^-6

The ratio would be [H3O+ acid]/[H3O+ normal] = [1 x 10^-3]/[2.5 x
10^-6] = 400/1

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Question 3: Answer 1. The concentration of OH- will always equal the
concentration of H3O+.

Using the Ideal Gas Equation (it can be found at
http://wine1.sb.fsu.edu/chm1045/notes/Gases/IdealGas/Gases04.htm) we
see that Volume is related to Pressure and Temperature.

PV = nRT, where R is a constant, and n is the quantity of gas measured
in mols. The pH of a substance is related to the concentration of H+,
so if we raise or lower the Temperature or Pressure of pure water, we
will alter the Volume (or concentration) of H3O+ and OH- ions. As we
alter the concentrations we alter the pH and pOH. The only answer that
is true no matter what the volume or concentration may be is #1.


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Question 4: Answer 5. The pH is 9.26.


From The Net Equation at Thinkquest.org
(http://library.thinkquest.org/C004970/acidbase/concepts.htm)
Kb is the dissociation constant for a base in water. For base in
water,

B + H2O < - - > BH+ + OH-
Kb = [BH+][OH-]/[B]

Utilizing a strategy from the Shodor Foundation
(http://www.shodor.org/unchem/basic/ab/) we can effectively solve this
problem.

          Hydroxylamine < - - > BH+ + OH- (it is not important to
define “B”)
Start		0.05 M           0     0
Change	          -x		 +x    +x
Equilibrium    0.05 – x          +x    +x

So Kb = [x][x]/[0.05 – x] we can exclude “- x” from the denominator
since it is such a small number in comparison.

6.6 x 10^-9 = x2/.05
Solving for x, x = 1.82 x 10^-5, so that is our concentration of OH-

pOH = -log[OH-], so pOH = 4.74

pH + pOH = 14.0, therefore

pH = 14.0 –4.74

pH = 9.26

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Question 5: Answer 3. The concentration of CO2 is 1.5 x 10^-4

Using a similar strategy to question 4 above, we will use the
acid-dissociation constant of carbonic acid (Ka) and a few equations
to solve this problem.

Ka = [H2O][CO2]/[H2CO3]

         H2CO3 < - - > H2O + CO2
Start	.05             0     0
Change   –x            +x    +x
Equil. .05 – x         +x    +x

Ka = 4.2 x 10^-7
Ka = [x][x]/[.05 – x] we can exclude “- x” from the denominator since
it is such a small number in comparison.

4.2 x 10^-7 = x2/.05
Solving for x, x = 1.5 x 10^-4 = [CO2]

Acid-dissociation constant (Ka) of carbonic acid was obtained at
Physchem, a South African science site.
http://www.physchem.co.za/Acids/Strength.htm#Value

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Again, I hope that these answers and explanation assist your daughter
in her studies. Should anything be unclear, please let me know so I
can clarify it for you.

Sincerely,
Boquinha-ga

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Google Search Strategy:
acid base chemistry equation
conjugate acid water
ideal gas equation
acid dissociation constant carbonic acid

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Clarification of Answer by boquinha-ga on 21 Oct 2003 00:21 PDT

Sorry about the formatting on #4. Let me put in the equation with the
start, change, and equilibrium formulas again:

Hydroxylamine < - - > BH+ + OH-  (it is not important to define “B”)
0.05 M                0     0    (Start)
  -x                  +x    +x   (Change)
0.05 – x              +x    +x   (Equilibrium)

Hope this turns out better!

Boquinha-ga

Hello boquinha,

You get a 5 again ! Thank you very much, keep up the good job, you are very helpful.

Just a comment that although the answer given to question 3 ("In a
sample of pure water, only one of the following statements is always
true at all conditions of temperature and pressure. Which one is
always true?"), is correct, the rationale boquinha-ga used to support
his answer is not correct.  He tried to use an argument based on the
ideal gas law, a law that is not applicable to liquids, the subject of
this question.  (Liquid water is a particularly bad misapplication of
the ideal gas law because at 1 atmpsohere, the volume of a given
quantity of water actually *decreases* as the temperature increases
between 0 and 4 degrees centigrade, which is the opposite behavior of
all gasses and even of most condensed (liquid or solid) substances! 
At 1 atmosphere pressure, the density of water has a maximum at ~4
degrees centigrade.)

The correct logic involves recognizing that all the possible answers
2involve the water dissociation reaction: 2*H2O(liquid) <->
H3O+(aqueous) + OH-(aqueous).  As with all reactions, one can write an
equilibrium constant for the reaction:

K_w = [H3O+]*[OH-]/[H2O]^2

Where the brackets mean the concentrations of the species in the
brackets.  (Strictly speaking, these really mean thermodynamic
activities, not concentrations, but this is a distinction not usually
taught in high school chemistry.) By convention, the concentrations
(activities) of pure substances (e.g., H2O) are defined as being equal
to unity, so the equilibrium constant only involves the concentrations
(activities) of hydronium and hydroxide ions.

All equilibrium constants are a function of temperature and pressure. 
That means that the concentrations of H3O+ and OH- will vary as a
function of temperature and pressure (and because pH and pOH are
simply other ways of expressing the concentrations of H3O+ and OH-,
they, too will vary as a function of T and P).

Answer 1, the correct one, however, follows from a consideration of
charge conservation.  If one starts with electrically neutral water
molecules, which dissociate into positive hydronium ions and negative
hydroxide ions, the positive and negative charges must add up to zero
(because that's what we started with.)  In any sample of water, the
number of positive ions (H3O+) must be equal to the number of negative
charges (OH-) and [H3O+] = [OH-] (In this case, the square brackets
really *do* mean concentrations, not thermodynamic activities!)

Again, thank you for the kind words, 5-star rating, and generous tip.
I wish you and your daughter the best!

Sincerely,
Boquinha-ga