Transcript Mg 2+
Chapter 6 Problems
6.6, 6.9, 6.15, 6.16, 6.19, 6.21, 6.24
Comments on “Lab Report & Pop Rocks”
6.6
Flip
From the equations HOCl H + + OCl HOCl + OBr HOBr + OCl HOBr + OCl HOCl + OBr Find K for HOBr H + + OBr — K = 3.0 x 10 -8 K= 15 K’ = 1/K= 1/15 K= ?
-8 /15 K= 0.2 x 10 -8
6.9
The formation of Tetrafluorethlene from its elements is highly exothermic. 2 F 2 (g) + 2 C (s) F 2 C=CF 2 (g) (a) if a mixture of F 2 , graphite, and C 2 F 4 equilibrium in a closed container, will the reaction go to the right or to the left if F 2 is at is added?
(b) Rare bacteria … eat C 2 F 4 and make teflon for cell walls. Will the reaction go to the right or to the left if these bacteria are added?
6.9
The formation of Tetrafluorethlene from its elements is highly exothermic . 2 F 2 (g) + 2 C (s) F 2 C=CF 2 (g) (C) will the reaction go to the right or to the left if graphite is added?
(d) will reaction go left or right if container is crushed to one-eighths of original volume?
(e) Does “Q” get larger or smaller if vessel is Heated?
6-15.
What concentration of Fe(CN) 6 4 is in equilibrium with 1.0 uM Ag + and Ag 4 Fe(CN) 6 (s). Ag 4 Fe(CN) 6 4Ag + + Fe(CN) 6 4 K sp = [Ag + 8.5 x 10 -45 ] 4 [Fe(CN) 6 4 ] = [1.0 x 10 -6 ] 4 [Fe(CN) 6 4 ] [Fe(CN) 6 4 ] = 8.5 x 10 -21 M = 8.5 zM
6-16.
I Cu 4 (OH) 6 (SO 4 ) 4 Cu 2+ + 6OH + SO 4 2 I’d first set up an ICE table: C E Cu 4 (OH) 6 (SO 4 ) Some -x Some –x 4 Cu 2+ + 4x + 4x + 6 OH 1.0 x 10 -6 M
Fixed at 1.0 x 10 -6 M
Fixed at 1.0 x 10 -6 M + SO 4 2-
+
x + x Ksp = [Cu Ksp = [4x] 2+ 4 ] 4 [OH ] 6 [SO 4 2 ] = 2.3 x 10 -69 [1.0 x 10 -6 ] 6 [x] = 2.3 x 10 -69 X = 9.75 x 10 -8 M Cu 2+ = 4x = 3.9
0 x 10 -7 M
Chapter 6 Chemical Equilibrium
Chemical Equilibrium
Equilibrium Constant Solubility product (K sp ) Common Ion Effect Separation by precipitation Complex formation
Separation by Precipitation
Separation by Precipitation
Complete
separation can mean a lot … we should define complete.
Complete means that the concentration of the less soluble material has decreased to 1 X 10 -6 M or lower before the more soluble material begins to precipitate
Separation by Precipitation
EXAMPLE:
Can Fe quantitatively as hydroxides from a solution that is 0.10 M in each cation? If the separation is possible, what range of OH +3 and Mg +2 be separated concentrations is permissible.
Add OH -
Mg Fe 3+ Mg 2+ Fe 2+ 3+ Fe 3+ Fe 3+ Mg 2+ Mg 2+ Mg 2+ Fe 3+ Fe 3+ Mg Fe 3+ Mg 2+ 2+ Mg 2+ Mg 2+ Fe 3+ Mg 2+ Fe 3+ Mg 2+ Fe 3+ Fe Fe 3+ 3+
Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Fe 3+ Mg 2+ Mg 2+ @ equilibrium What is the [OH this happens ] when
Fe(OH) 3 (s)
EXAMPLE: Separate Iron and Magnesium?
K sp = [Fe +3 ][OH ] 3 = 2 X 10 -39 K sp = [Mg Assume [Fe +3 ] eq precipitated. +2 ][OH ] 2 = 7.1 X 10 -12 = 1.0 X 10 -6 M when “completely” What will be the [OH ] @ equilibrium required to reduce the [Fe +3 ] to [Fe +3 ] = 1.0 X 10 -6 M ? K sp = [Fe +3 ][OH ] 3 = 2 X 10 -39
EXAMPLE: Separate Iron and Magnesium?
K sp = [Fe +3 ][OH ] 3 (1.0 X 10 -6 M)*[OH ] 3 = 2 X 10 -39 = 2 X 10 -39 [
OH
] 3 2 10 33 [
OH
] 1 .
3 10 11
Dealing with Mg 2+
Find [OH ] to start precipitating Mg 2+
• Conceptually – •Will assume a minimal amount
of Mg 2+ will precipitate and determine the respective concentration of OH Evaluate Q
• If •Q>K •Q
“Left” “Right” “Equilibrium”
Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Fe 3+ Mg 2+ Mg 2+ @ equilibrium [OH ] ^
= 1.3 x 10 -11
Is this [OH Mg 2+?
] (that is in solution) great enough to start precipitating
Fe(OH) 3 (s)
Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Mg 2+ Fe 3+ Mg 2+ Mg 2+ @ equilibrium [OH ] ^
= 1.3 x 10 -11
Is this [OH Mg 2+?
] (that is in solution) great enough to start precipitating
Fe(OH) 3 (s)
EXAMPLE: Separate Iron and Magnesium?
What [OH ] is required to begin the precipitation of Mg(OH) 2 ?
[Mg +2 ] = 0.10 M
Really, Really close to 0.1 M [Mg 2+ ] eq = 0.09999999999999999 M
K sp = (0.10 M)[OH ] 2 = 7.1 X 10 -12 [OH ] = 8.4 X 10 -6 M
EXAMPLE: Separate Iron and Magnesium?
@ equilibrium
[OH ^ ‘completely = 1.3 X 10 -11 M ’ remove Fe 3+ [OH ] to start = 8.4 X 10 -6 M removing Mg 2+
“All” of the Iron will be precipitated b/f any of the magnesium starts to precipitate!!
EXAMPLE: Separate Iron and Magnesium?
Q vs. K
Mg(OH) 2 (s)
Mg 2+ + 2OH -
K sp = [Mg 2+ ][OH ] 2 = 7.1 X 10 -12 Q = [0.10 M ][
1.3 x 10 -11
] 2 = 1.69 x 10 -23
Q
Dealing with Mg 2+
Find [OH ] to start precipitating Mg 2+
• Conceptually – •Will assume a minimal amount
of Mg 2+ will precipitate and determine the respective concentration of OH NO PPT Evaluate Q
• If •Q>K •Q
“Left” “Right” “Equilibrium”
“Real Example”
Consider a 1 liter solution that contains 0.3 M Ca 2+ and 0.5 M Ba 2+ . Can you separate the ions by adding Sodium Carbonate?
Sodium Chromate ?
Sodium Fluoride?
Sodium Hydroxide?
Sodium Iodate?
Sodium Oxylate?
An example
Consider Lead Iodide PbI 2 (s) Pb 2+ + 2I K sp = 7.9 x 10 -9
What should happen if I solution?
is added to a Should the solubility go up or down?
Complex Ion Formation
Complex Formation
complex ions (also called coordination ions) Lewis Acids and Bases acid => electron pair acceptor (metal) base => electron pair donor (ligand)
I Pb 2+ I Pb 2+
I Pb 2+ I Pb 2+
I I I I I I I I Pb 2+ Pb 2+ Pb 2+ I I I I Pb I Pb 2+ Pb I Pb 2+
I I I I I Pb 2+ I I I I I Pb I 2+ I Pb 2+ 2+ Pb I Pb I Pb Pb 2+ Pb 2+ Pb 2+
I I I I I I I Pb 2+ I I I I I I I I I I I Pb I 2+ I Pb 2+ 2+ Pb I Pb I Pb Pb 2+ Pb 2+ Pb 2+
I I I I I I I Pb 2+ I I I I I I I I I I I Pb I 2+ I Pb 2+ 2+ Pb I Pb I Pb Pb 2+ Pb 2+ Pb 2+
Effects of Complex Ion Formation on Solubility
Consider the addition of I of Pb +2 ions to a solution Pb 2+ + I <=> PbI + PbI + + I <=> PbI 2
K
1 [ [
PbI Pb
2 ][ ]
I
] 1 .
0
x
10 2 K 2 = 1.4 x 10 1 PbI 2 + I <=> PbI 3 PbI 3 + I <=> PbI 4 2 K 3 =5.9
K 4 = 3.6
Effects of Complex Ion Formation on Solubility
Consider the addition of I of Pb +2 ions Pb 2+ + I <=> PbI + to a solution
K
1 [ [
PbI Pb
2 ][ ]
I
] 1 .
0
x
10 2 PbI + + I <=> PbI 2 K 2 = 1.4 x 10 1 Pb 2+ + 2I <=> PbI 2 K’ =?
Overall constants are designated with
b
This one is
b 2
Protic Acids and Bases
Section 6-7
Question
Can you think of a salt that when dissolved in water is not an acid nor a base?
Can you think of a salt that when dissolved in water IS an acid or base?
Protic Acids and Bases Salts
Consider Ammonium chloride Can ‘generally be thought of as the product of an acid-base reaction.
NH 4 + Cl (s) NH 4 + + Cl From general chemistry – single positive and single negative charges are STRONG ELECTROLYTES – they dissolve completely into ions in dilute aqueous solution
Protic Acids and Bases
Conjugate Acids and Bases in the B-L concept CH 3 COOH + H 2 O CH 3 COO + H 3 O + acid + base <=> conjugate base + conjugate acid conjugate base => what remains after a B-L acid donates its proton conjugate acid => what is formed when a B-L base accepts a proton
Question: Calculate the Concentration of H + and OH in Pure water at 25 0 C.
EXAMPLE: Calculate the Concentration of H + and OH in Pure water at 25 0 C.
I nitial C hange E quilibrium H 2 O liquid -x Liquid-x H +x +x + + OH K w = [H + ][OH ] = 1.01 X 10 K W = (X)(X) = 1.01 X 10 -14 -14 (X) = 1.00 X 10 -7 +x +x -
Example
K w Concentration of OH if [H + ] is 1.0 x 10 -3 M @ 25 o C? = [H 1 x 10 1 x 10 -14 + -11 ][OH ] = [1 x 10 = [OH ] -3 ][OH ] “From now on, assume the temperature to be 25 o C unless otherwise stated.”
pH
~ -3 -----> ~ +16 pH + pOH = - log K w = pK w = 14.00
Is there such a thing as Pure Water?
In most labs the answer is NO Why?
CO 2 + H 2 O HCO 3 + H +
A century ago, Kohlrausch and his students found it required to 42 consecutive distillations to reduce the conductivity to a limiting value.
6-9 Strengths of Acids and Bases
Strong Bronsted-Lowry Acid
A strong Bronsted-Lowry Acid is one that donates all of its acidic protons to water molecules in aqueous solution. (Water is base – electron donor or the proton acceptor).
HCl as example
Strong Bronsted-Lowry Base
Accepts protons from water molecules to form an amount of hydroxide ion, OH , equivalent to the amount of base added.
Example: NH 2 (the amide ion)
Weak Bronsted-Lowry acid
One that DOES not donate all of its acidic protons to water molecules in aqueous solution. Example?
Use of double arrows! Said to reach equilibrium.
Weak Bronsted-Lowry base
Does NOT accept an amount of protons equivalent to the amount of base added, so the hydroxide ion in a weak base solution is not equivalent to the concentration of base added.
example: NH 3
Common Classes of Weak Acids and Bases
Weak Acids
carboxylic acids ammonium ions
Weak Bases
amines carboxylate anion
Weak Acids and Bases
K a HA H + + A -
K a
[
H
][
A
] [
HA
] K a ’ s ARE THE SAME HA + H 2 O (l) H 3 O + + A -
K a
[
H
3
O
][
A
] [
HA
]
Weak Acids and Bases
K b B + H 2 O BH + + OH -
K b
[
BH
][
OH
] [
B
]
Relation Between K
a
and K
b
Relation between Ka and Kb
NH 3 Consider Ammonia and its conjugate base.
K a
+ H 2 O NH 4 + + OH -
K a
[
NH
4 ][
OH
] [
NH
3 ] NH 4 + + H 2
K
O NH 3 + H 3 O +
K b
[
NH
[ 3 ][
NH H
3 ]
O
4 ] H 2 O + H 2 O OH + H 3 O +
K K w
[
NH
[ [ 3 ][
H NH H
4 3 3
O O
] ] ] [
NH
[ [ 4 ][
OH OH
3 ] ] ]
Example
The K a for acetic acid is 1.75 x 10 -5 . Find K b for its conjugate base.
K w = K a x K b
K b
K w K a K b
1 .
0 10 14 1 .
75 10 5 5 .
7 10 10
1
st