# Rates of Reaction (and thermodynamics)

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27 Οκτ 2013 (πριν από 4 χρόνια και 6 μήνες)

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Chemical Equilibrium

Chapter 17

Review

Given a chemical equation…

A + B

C + D

Which letters represent the products?

Which letters represent the reactants?

What is a rate?

What is a concentration?

Reversible Reactions

A

B

B

A

Reversible reactions occur in both the
forward and reverse directions

CH
4

+ 2H
2
S

CS
2

+ 4H
2

Means that CH
4

+ H
2
S can form CS
2

+ H
2
or CS
2

+ H
2
can form CH
4

+ H
2
S

Or A

B

Equilibrium

When

A

B

B

A

The reactions are at
equilibrium
when the
forward and reverse reactions balance
each other because they take place at
equal rates

If A

B

r1

r2

r1 = r2

equilibrium

Equilibrium

Amounts of products and reactants stays
constant

Does not mean that there are equal amounts
of products and reactants: [A] does not have
to equal [B]

Catalysts speed both forward and backward
rates, so they do not affect equilibrium
concentrations

Law of chemical equilibrium: at a given
temperature, reactant and product
concentrations have a constant value

Equilibrium Constants

For a reversible reaction at equilibrium

aA + bB

cC + dD

K
eq
=

[C]
c

x [D]
d

[A]
a

x [B]
b

Exponents same as

coefficients in equation

Brackets indicate

concentration

Products over

reactants

If <1, more reactants
than products

If >1, more products
than reactants

Writing expressions for
Homogeneous Equilibria

All in the same state of matter

Ex. Write the equilibrium constant expression for
CH
4
(g) + 2H
2
S(g)

CS
2
(g) + 4H
2
(g)

K
eq

=

[CH
4
][H
2
S]
2

[CS
2
][H
2
]
4

Writing expressions for
Heterogeneous equilibria

Different states of matter

Because solids and liquids have a
concentration that doesn’t change, you
can omit them from the equilibrium
expression

Ex. What is the equilibrium constant expression
for MgCO
3
(s)

MgO(s) + CO
2
(g) ?

[MgCO
3
]

[MgO][CO
2
]

Keq=

= [CO
2
]

Calculating equilibrium constants

Plug in concentration data (experimental)
into K
eq

expression

Ex. At a certain temperature, the following
concentrations are measured. Calculate Keq
if
CH
4
(g) + 2H
2
S(g)

CS
2
(g) + 4H
2
(g) and

[CH
4
]=0.5 mol/L

[CS
2
]=0.5 mol/L

[H
2
S]=0.7 mol/L

[H
2
]=0.8 mol/L

[CH
4
][H
2
S]
2

[CS
2
][H
2
]
4

Keq=

(0.5)(0.7)
2

(0.5)(0.8)
4

=0.84

Calculating K
eq

Calculate Keq for the following reaction:

2SO
2

+ O
2

2SO
3

If there are 0.1 mol , 0.2 mol, and 0.3 mol

More on equilibrium constants

K
eq

should always equal the same thing for a
given reaction at a given temperature

An infinite number of equilibrium positions:
varies depending on initial concentration

Must be a closed system: no reactant or
product can be added or escape

Temperature must remain constant

Dynamic: forward and reverse reactions don’t
stop

Le Chatelier’s Principle

If a stress is applied to a system in
dynamic equilibrium, the system
changes to relieve the stress

3 main stresses

Changing concentration of reactants or
products

Changing the temperature

Changing the volume or pressure

Le Chatelier’s Principle

Concentration

2SO
2

+ O
2

2SO
3

Increasing the concentration of a reactant
will shift the equilibrium so that more of it
is used up

Decreasing the concentration of a product
will shift the equilibrium so that more of it

Shifts reaction this way

Le Chatelier’s Principle

Pressure

2SO
2

(g)+ O
2
(g)

2SO
3
(g)

If pressure is increased (or volume decreased), it
favors the side with fewer moles of gas

If pressure is decreased (or volume increased), it
favors the side with more moles of gas

Shifts reaction this way

If we increase the pressure…

If we increase the volume…

Le Chatelier’s Principle

Temperature

2SO
2

+ O
2

2SO
3
+ heat

If heat is added, reaction will shift to favor the
endothermic reaction
(you can treat heat like a
chemical) . If the forward reaction is exothermic
and T increases, Keq decreases.

Shifts reaction this way

If we increase the temperature…

If we decrease the temperature…

Using equilibrium constants

If you know an equation and Keq, you
can calculate equilibrium concentrations.

Ex.
If 2ICl

I
2

+ Cl
2

K
eq
=0.110 at a particular temperature,
[I
2
]=0.0330, and [Cl
2
]=0.220, what is the
[ICl]?

[
ICl
]
2

[
I
2
][
Cl
2
]

K
eq

=

[
ICl
]
2

(0.033)(0.22)

0.110

=

[
ICl
]=0.257 mol/L

Solubility Product Constants

Recall that compounds vary in solubility

At saturation…

AgCl(s)

Ag
+
(aq) + Cl
-
(aq)

So we can write a K
eq

for this process:

K
eq
=[Ag+ ][Cl
-
]

[AgCl]

K
sp
=[Ag
+

][Cl
-
]

(multiply both sides by [AgCl])

=1.7 x 10
-
5

Solubility product constant

Solubility Products

K
sp
=[Ag
+

][Cl
-
]

= 1.7 x 10
-
5

What is the concentration of each ion at
saturation?

Since

AgCl(s)

Ag
+
(aq) + Cl
-
(aq)

[Ag
+
]=[Cl
-
]=x

So,

x
2
= 1.7 x 10
-
5

Predicting Precipitates and
Calculating Ion Concentrations

To determine saturation you can do a
trial solubility product (Qsp) and compare
it to Ksp

Ex. Will CaF
2

precipitate if equal volumes
0.020 M Ca(NO
3
)
2

and 0.0064 M NaF
are mixed at 298 K? Ksp for CaF
2
=3.5 x
10
-
11

Qsp=[Ca
2+
][F
-
]
2
=(0.01)(0.0032)
2
=1.0 x10
-
7

Qsp > Ksp

Yes

Note: concentrations are reduced by half because
the volume of solution is now 2x as big

The Common Ion Effect

The solubility of a compound will
decrease if you mix it with a solution with
the same ion.

A precipitate will form

Ex: When mixing BaSO
4

and NaSO
4
,
BaSO
4

will precipitate out

The Common Ion effect

What is the concentration of [Br
-
] if we
add 0.020 mol AgNO3 to 1.00 L of
saturated AgBr?