Download Unit 7 Gas Laws and Thermodynamics

Reversible Reactions
 A chemical reaction in which the products
can react to re-form the reactants is called a
reversible reaction.
 A reversible reaction is written using double
arrows to show that the reaction is
proceeding in both directions.
Example:

 2Hg(l ) + O2 (g )
2HgO(s ) 

Dynamic Equilibrium
 A reversible reaction reaches
dynamic equilibrium when the
rate of its forward reaction equals
the rate of its reverse reaction
and the concentrations of its
products and reactants remain unchanged.
 At equilibrium, both reactions continue, but
there is no net change in the composition of
the system.
Visual Concept
Equilibrium  Equal
 At equilibrium, the rates of the forward and
reverse reactions are equal. But the
concentrations aren’t necessarily equal.
 Some reactions reach equilibrium only after
almost all reactants are consumed (products
are favored.)
 Others reach equilibrium when only a small
percentage of reactants are consumed
(reactants are favored.)
Le Châtelier’s Principle
 Le Châtelier’s Principle: When a system at
equilibrium is disturbed, the system shifts in a
direction that minimizes the disturbance.
 A shift in equilibrium will result from a change
to any of the following:
 Concentration
 Volume/Pressure
 Temperature
Change in Concentration
 If you increase the
concentration by
adding more of a
reactant or product,
the system will shift
to produce less of
that substance.
 If you lower the concentration by removing
some of a reactant or product, the system will
shift to produce more of that substance.
Change in Volume/Pressure
 When you increase the
pressure (usually by
decreasing volume of the
container), the system
shifts so the least number
of gas molecules are
formed (less collisions
= lower pressure.)
 When you decrease the
pressure, the system will
shift so the greatest number
of gas molecules are formed.
Change in Temperature
 For every reversible reaction, one direction is
endothermic and the other is exothermic.
 If the temperature is increased, the endothermic
reaction will be favored (because it takes in some
of the excess heat.)
 If the temperature
is decreased,
the exothermic
reaction will be
favored (produces heat.)
Le Châtelier’s Principle
Sample Problem
2 SO2(g) + O2(g)  2 SO3(g)
DH° = -198 kJ
How will the reversible reaction above shift in response to
each of the following stresses?
• adding more O2 to the container
Shift right
• condensing and removing SO3
Shift right
• compressing the gases
Shift right
• cooling the container
Shift right
• doubling the volume of the container
Shift left
• warming the mixture
Shift left
The Law of Mass Action
 The relationship between the chemical
equation and the concentrations of reactants
and products is called the Law of Mass Action.
 for the general equation aA + bB  cC + dD,
the Law of Mass Action is:
c [D]d
[C]
K=
[A]a [B]b
 Lowercase letters represent coefficients.
 Always products over reactants.
 Pure solids and pure liquids are not included.
The Equilibrium Constant
 The equilibrium constant (K)
reflects how the concentrations
of the reactants and products
compare at equilibrium.
 It can also be a ratio of pressures
(in atmospheres) if a reaction
involves gaseous reactants
and/or products.
 K is unitless.
The Value of K
 K > 1 more product molecules present than
reactant molecules (the position of equilibrium
favors products.)
 K < 1 more reactant molecules present than
product molecules (the position of equilibrium
favors reactants.)
 K = 1 reactant and product particles are
present in exact equal concentrations at
equilibrium.
Equilibrium Constant
Sample Problem
Equilibrium concentrations of [H2] = 0.033 M, [I2] = 0.53 M
and [HI] = 0.934 M were observed at 445oC for the reaction:
H2(g) + I2(g)  2 HI(g)
1.
Write an equilibrium expression for the above reaction.
[HI]2
K=
[H2][I2]
2. Calculate the value of Kc for this reaction at 445oC.
[0.934]2
[HI]2
= 49.9
=
Kc =
[H2][I2] [0.033] [0.53]
The Reaction Quotient
 When a reaction is not at equilibrium, how do
you know in which direction it will proceed?
 the answer is to compare the equilibrium
constant to a ratio of current concentrations
called the reaction quotient (Q).
 for the general equation aA + bB  cC + dD:
c [D]d
[C]
Q=
[A]a [B]b
 The non-equilibrium concentrations (or
pressures) are used.
Q vs. K
 We calculate Q in order to compare it with K.
 Q < K means the reaction will proceed in the
forward direction ([products] increase and
[reactants] decrease.)
 Q > K means the reaction will proceed in the
reverse direction ([products] decrease and
[reactants] increase.)
 Q = K means the reaction is at equilibrium
([products] and [reactants] will not change.)
Q, K, and the Direction of Reaction
Reaction Quotient
Sample Problem
For the reaction below, which direction will it proceed
if PI2 = 0.114 atm, PCl2 = 0.102 atm & PICl = 0.355 atm?
I2(g) + Cl2(g)  2 ICl(g)
Kp = 81.9
First calculate Q:
(0.355)2
(ICl)2
= 10.8
=
Q=
(I2)(Cl2) (0.114)(0.102)
Then, compare it with K:
Q < K
10.8
81.9
Reaction will proceed to the right
Properties of Acids
 Taste sour.
 React with metals to
release H2 gas.
 React with bases to
produce salts and water.
 Change the color of
acid-base indicators.
 Conduct electric current.
Properties of Bases
 Taste bitter.
 Feel slippery.
 React with acids to
produce salts and water.
 Change the color of
acid-base indicators.
 Conduct electric current.
Arrhenius Acids and Bases
 An Arrhenius acid produces hydrogen ions,
H+, in aqueous solution.
 An Arrhenius base produces hydroxide ions,
OH−, in aqueous solution.
 A strong acid (or base) ionizes completely.
 A weak acid (or base) releases only a few ions.
Arrhenius Theory
HCl ionizes in water,
producing H+ and Cl– ions
NaOH dissociates in water,
producing Na+ and OH– ions
Hydronium Ion
 The H+ ions (protons) produced by the acid
are so reactive they cannot exist in water.
 instead, they react with a water molecule to
form a hydronium ion, H3O+.
H+ + H2O  H3O+
 Chemists use H+ and H3O+ interchangeably.
Brønsted-Lowry Acids and Bases
 In a Brønsted-Lowry Acid-Base reaction, an
H+ ion (proton) is transferred.
 Does not have to take place in aqueous
solution.
 Broader definition than Arrhenius.
 A Brønsted-Lowry acid is a molecule or ion
that is a proton donor.
 A Brønsted-Lowry base
is a molecule or ion that is
a proton acceptor.
Conjugate Pairs
 Brønsted-Lowry theory allows for reversible reactions.
 The original base has an extra H+ after the reaction.
It will act as an acid in the reverse process.
 The original acid has a lone pair of electrons after the
reaction. It will act as a base in the reverse process.
 each reactant and the product it becomes is called a
conjugate pair.
Brønsted-Lowry Acid-Base Reactions
Sample Problem
Identify the Brønsted-Lowry Acids and Bases and Their
Conjugates in the Reactions below:
a. H2SO4
+ H2O 
acid
b. HCO3– +
base
base
HSO4– + H3O+
conjugate conjugate
base
acid
H2O 
H2CO3
+
HO–
conjugate
conjugate
acid
acid
base
Amphoteric Compounds
 An amphoteric substance is one that can
react as either an acid or a base.
Example: water
 Water can act as an acid.

–


NH3 (g ) + H2O(l ) 
NH
(
aq
)

OH
(aq )

4
base

acid
Water can act as a base.
H2SO4 (aq ) + H2O(l )  H3O (aq ) + HSO4– (aq )
acid
base
Visual Concept
Polyprotic Acids
 Molecules with more than one ionizable H are
called polyprotic acids.
 1 H = monoprotic, 2 H = diprotic, 3 H = triprotic
(Ex: HCl = monoprotic, H2SO4 = diprotic,
H3PO4 = triprotic)
 Polyprotic acids ionize in steps (each
ionizable H removed sequentially.)
 Removing the first H automatically
makes removing the second H harder.
(Ex: H2SO4 is a stronger acid than HSO4)
Lewis Acids and Bases
 A Lewis acid is an atom, ion, or molecule that
accepts an electron pair to form a covalent bond.
 A Lewis base is an atom, ion, or molecule that
donates an electron pair to form a covalent bond.
• The Lewis definition is the broadest of the
three definitions.
Comparing the Three Definitions
Visual Concept
Strong and Weak Acids and Bases
 Strong acid – fully
dissociates in water (almost
every molecule breaks up
to form H+ ions.
 Weak acid – partially
dissociates in water.
 Strong base – fully
dissociates in water (almost
A Strong Acid
every molecule breaks up
to form OH- ions.
 Weak base – partially
dissociates in water.
A Weak Acid
Ionization Constants for Weak
Acids and Bases
 The acid ionization constant (Ka) is the
equilibrium constant for the ionization reaction
of a weak acid: HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
[H3O+] [A-]
Ka =
[HA]
 A base ionization constant (Kb) can also be
created for the ionization reaction of a weak
base: B(aq) + H2O(l)  BH+(aq) + OH-(aq)
[BH+] [OH-]
Kb =
[B]
Acid Ionization Constant
Sample Problem
Calculate the Ka of a 0.100 M solution of acetic acid
with a measured [H3O+] of 1.34 x 10-3 M.
HC2H3O2(aq) + H2O(l)  H3O+(aq) + C2H3O2-(aq)
[H3O+][C2H3O2-] [1.34 x 10-3M][1.34 x 10-3M]
=
Ka=
= 1.82 x 10-5
[0.0987 M]
[HC2H3O2]
•
•
For every H3O+ produced, there is also a C2H3O2produced, so the concentrations must be the same.
The equilibrium concentration of original acid is the
original concentration decreased by the amount ionized
(0.100M – 1.34 x 10-3 M = 0.0987 M)
Ka & Kb and Strength of Acids/Bases
 The strength of an acid or base is measured by
the size of its equilibrium constant when it
reacts with H2O.
Six Strong Acids and Bases
 Because these acids and bases are known to
dissociate (ionize) to essentially 100% completion,
it is meaningless to connect them to equilibrium:
6 Strong Acids
6 Strong Bases
HCl – Hydrochloric Acid
Ca(OH)2 – Calcium Hydroxide
HBr – Hydrobromic Acid
Sr(OH)2 – Strontium Hydroxide
HI – Hydroiodic Acid
Ba(OH)2 – Barium Hydroxide
H2SO4 – Sulfuric Acid
LiOH – Lithium Hydroxide
HNO3 – Nitric Acid
NaOH – Sodium Hydroxide
HClO4 – Perchloric Acid
KOH – Potassium Hydroxide
is at donating H,
the weaker the
conjugate base is
at accepting H.
(i.e. strong acids
have weak conjugate
bases, and weak acids
have strong conjugate
bases.)
Increasing Acidity
 the stronger an acid
Conjugate Bases
ClO4-1
H2SO4
HI
HBr
HCl
HNO3
H3O+1
HSO4-1
H2SO3
H3PO4
HNO2
HF
HC2H3O2
H2CO3
H2S
NH4+1
HCN
HCO3-1
HS-1
H2O
CH3-C(O)-CH3
NH3
CH4
OH-1
HSO4-1
I-1
Br-1
Cl-1
NO3-1
H2O
SO4-2
HSO3-1
H2PO4-1
NO2-1
F-1
C2H3O2-1
HCO3-1
HS-1
NH3
CN-1
CO3-2
S-2
OH-1
CH3-C(O)-CH2-1
NH2-1
CH3-1
O-2
Increasing Basicity
Strengths of
Conjugate Pairs
Acids
HClO4
Predicting Acid Strength
 The strength of an acid depends
on its tendency to ionize (let go
of its hydrogen.)
 For binary acids, the
strength of the acid depends
on two factors:
1. The stronger the bond,
the weaker the acid.
2. The more polar the bond,
the stronger the acid.
Periodic Trends are a BEAR
All of these increase in the direction toward their letter:
B=Basicity
E=Electronegativity, ionization Energy, & Electron Affinity
A=Acidity
R=Radius
B
E
R
A
Predicting Acid Strength (continued)
 The strength of Oxyacids of the form H-O-Y,
where Y is any atom (besides H) bonded to O,
depends on two factors:
1. The more electronegative
the element Y, the
stronger the acid.
2. The greater the number
of oxygen atoms
bonded to Y, the
stronger the acid.
Predicting Acid Strength
Sample Problem
Predict the relative strengths of the following acids:
• HCl, HBr, and HI
• HClO, HBrO, and HIO
• HNO3 and HNO2
HCl < HBr < HI
HIO < HBrO < HClO
HNO2 < HNO3
Predict the relative strengths of the following bases:
• Cl-, Br- and I-
I- < Br- < Cl-
• H2PO3- and H2PO4-
H2PO4- < H2PO3-
• LiOH and Mg(OH)2
LiOH < Mg(OH)2
Autoionization of Water
 Water is actually an extremely weak electrolyte.
 About 1 out of every 10 million water molecules
form ions through a process called autoionization.
H2O + H2O  H3O+ + OH–
 In pure water at 25°C, [H3O+] = [OH–] = 10-7M.
Ion Product of Water
 The product of the H3O+ and OH–
concentrations is always the same number,
called the ion product of water (Kw).
Kw = [H3O+] [OH-] = 1 x 10-14 at 25°C
 If you measure one of the concentrations, you
can calculate the other.
 As [H3O+] increases the [OH–] must decrease
so the product stays constant.
Acidic and Basic Solutions
 Neutral solutions have equal [H3O+] and [OH–]
 [H3O+] = [OH–] = 1 x 10-7
 Acidic solutions have a larger [H3O+] than [OH–]
 [H3O+] > 1 x 10-7; [OH–] < 1 x 10-7
 Basic solutions have a larger [OH–] than [H3O+]
 [H3O+] < 1 x 10-7; [OH–] > 1 x 10-7
Acidic, Basic or Neutral?
Sample Problem
Calculate [OH] at 25°C when [H3O+] = 1.5 x 10-9 M, and
determine if the solution is acidic, basic, or neutral.
First calculate [OH-]: Kw = [H3O+][OH-]
-14
K
1.0
x
10
w
-6 M
=
6.7
x
10
[OH-] =
=
-9
[H3O+] 1.5 x 10
Then, compare [H3O+] with [OH-]:
[H3O+]
1.5 x 10-9 M
<
[OH-]
6.7 x 10-6 M
The solution is basic
The pH Scale
 The acidity/basicity of a solution is often
expressed as pH.
 pH is defined as the negative of the common
logarithm of the hydronium ion concentration.
pH = −log [H3O+]
 pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral.
pH of Some Common Substances
The pH Scale
Sample Problem
Calculate the pH at 25°C when the [OH] = 1.3 x 10-2 M,
and determine if the solution is acidic, basic, or neutral
First calculate [H3O+]: Kw = [H3O+][OH-]
-14
K
1.0
x
10
w
-13 M
=
7.7
x
10
[H3O+] =
=
-2
[OH-] 1.3 x 10
Then, calculate pH:
pH = −log [H3O+] = -log(7.7 x 10-13 M)
pH = 12.1
The solution is basic
Finding the Ionization Constant from pH
Sample Problem
A 0.100 M weak acid (HA) solution has a pH of 4.25.
Find Ka for the acid.
First calculate [H3O+]: pH = -log [H3O+]
-pH = 10-4.25 = 5.6 x 10-5 M
+
10
=
[H O ]
3
Then, calculate Ka: HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
[H3O+] [A-] (5.6 x 10-5) (5.6 x 10-5)
= 3.1 x 10-8
Ka =
=
[HA]
(0.100 - 5.6 x 10-5)
The pOH Scale
 Another way of expressing the acidity/basicity of
a solution is pOH.
 pOH is defined as the negative of the common
logarithm of the hydroxide ion concentration.
pOH = −log [OH-]
 pOH < 7 is basic; pOH > 7 is acidic, pOH = 7 is neutral
Relationship Between pH and pOH
 The sum of the pH and pOH of a solution is 14.
pH + pOH = 14.0
Neutralization Reactions
• A neutralization reaction is a double
displacement reaction in which an acid and
a base in an aqueous solution react to
produce a salt and water.
• A salt is an ionic compound made up of a
cation from a base and an anion from an
acid.
Acid Base Properties of Salts
 Salts of strong acids and strong bases are neutral.
Ex: HCl(aq) + NaOH (aq)  NaCl (aq) + H2O (l)
 Salts of strong acids and weak bases are acidic.
Ex: NH3(aq) + HCl(aq)  NH4Cl (aq)
 Salts of strong bases and weak acids are basic
Ex: 2NaOH(aq)+ H2CO3 (aq)  Na2CO3 (aq) + 2H2O(l)
 Salts of weak acids and weak bases can be acidic,
basic or neutral depending on the relative
strength of acids and bases.
Buffers
 Buffers are solutions
that resist changes in
pH when an acid or base
is added.
 Buffers contain both a
weak acid and its
conjugate base (or a
weak base and its
conjugate acid.)
 The weak acid neutralizes added base.
 The conjugate base neutralizes added acid.
Buffering Effectiveness
 A good buffer should be able to neutralize
moderate amounts of added acid or base.
 However, there is a limit to how much can be
added before the pH changes significantly.
 The buffering capacity is the amount of acid or
base a buffer can neutralize.
 The buffering range is the pH range over which
the buffer can be effective.
 The effectiveness of a buffer depends on:
The relative amounts of acid and base.
2. The absolute concentrations of acid and base.
1.
Buffers in Human Blood
 Many of the chemical reactions that occur
in the body are pH-dependent.
 Ideally, the pH of the blood should
be maintained at a slightly basic 7.4.
 pH below 6.8 or above 7.8 can be fatal.
 Fortunately, we have buffers in the blood to
protect against large changes in pH.
Acid-Base Indicators
 An acid-base indicator is a chemical dye that
changes colors at definite pH values.
 There are a variety
of indicators that
change color at
different pH levels.
 A properly selected
indicator can be used
to visually "indicate"
the approximate pH of a sample.
Litmus Paper
 A common indicator is found
on litmus paper. It is red below
pH 4.5 and blue above pH 8.2.
Phenolphthalein
 Phenolphthalein is an organic compound
often used as an acid-base indicator.
 Phenolphthalein is colorless in acidic
solutions, but turns pink when the pH is
greater than 8.3.
Bromothymol Blue
 Bromothymol Blue (BTB) is a useful indicator
for substances that have a relatively neutral pH
(near 7).
 BTB is yellow in acids,
green in neutral solutions,
and blue in bases.
Acid-Base Titration
 Titration is a method for
determining the concentration
of a solution by reacting a known
volume of that solution with a
solution of known concentration.
 The analyte is a measured
analyte
volume of an acid or base
of unknown concentration.
 The standard solution (titrant)
is an acid or base solution whose
concentration is known.
Standard
solution
(titrant)
Titration Procedure
 In a titration procedure, a measured
volume of an analyte is placed in a
beaker or flask, and initial pH recorded.
 The standard solution (titrant) is
filled in a burette.
 A couple of drops of an acid-base
indicator are added to the flask.
 The standard solution is slowly added
to the unknown solution in the flask.
 As the two solutions are mixed the
acid and the base are neutralized.
Titration Procedure (cont’d)
 As the base is added to the acid,
H+ reacts with OH– to form
water. But there is still excess
acid present so the color
does not change.
 Once enough base has been
added to neutralize all the acid,
the indicator changes color.
 The difficulty is determining
when there has been just enough titrant added to
complete the reaction…without going over!
The Equivalence Point
 End point - The point
at which an indicator
changes color.
 Equivalence point –
The point at which the
moles of acid added
equals the moles of
base that you started
with (should be the
same as the end point.)
 An abrupt change in
pH occurs at the equivalence point.
Acid-Base Titration
Sample Problem
The titration of 10.00 mL of HCl solution of unknown
concentration requires 12.54 mL of 0.100 M NaOH
solution to reach the end point. What is the First, write the
concentration of the unknown HCl solution? neutralization
reaction
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
12.54 mL
NaOH
1 L NaOH
1000 mL
NaOH
M=
moles
L
0.100 mol NaOH 1 mol HCl
-3
1.25
x
10
=
mol
NaOH
1
1 L NaOH
mol HCl
1.25 x 10-3 mol HCl
= 0.125 M HCl
=
0.010000 L HCl
Titration of Strong Acid with Strong Base
 Since the salt produced
is neutral, the solution
at the equivalence point
has a pH of 7.
 the pH starts off low and
increases as you add
more of the base.
 The pH doesn't change
very much until you get
close to the equivalence
point. Then it surges
upwards very steeply
Titration of Strong Base with Strong Acid
 This curve is very similar
to the previous one for
the titration of a strong
acid with strong base.
 The main difference is
that the curve starts
basic and then turns
acidic after the
equivalence point
(rather than vice-versa.)
Titration of Weak Acid with Strong Base
 The salt is basic, so equivalence point is at a pH > 7.
 Before the equivalence point, the solution acts as a
buffer. The start of the graph shows a relatively
rapid rise in pH but this slows down due to the
buffering effect.
Titration of Weak Base with Strong Acid
 Salt formed is acidic,
hence, equivalence
point comes at a pH < 7.
 This curve is very similar
to the titration of a weak
acid with a strong base.
 The main differences are
that the curve starts
basic and has an acidic
equivalence point.
Titration of a Polyprotic Acid
 A polyprotic acid
titration will have more
than one equivalence
point.
 The first equivalence
point represents the
titration of the first
proton, while the second
equivalence point
represents the titration
of the second proton.
Titration Curves
Sample Problem
Two acidic solutions were titrated with a strong base.
Which curve represents a weak acid and which
represents a strong acid?
Strong Acid
Weak Acid
Acid Rain
 Rain is naturally somewhat acidic (pH ~5.6)
due to atmospheric CO2. Carbon dioxide
combines with rainwater to form carbonic acid:
CO2 + H2O → H2CO3.
 Rain water with
a pH < 5.6 is
called acid rain.
 Acid rain is
linked to damage
in ecosystems and
structures.
What Causes Acid Rain?
 Nonmetal oxides such as SO2 and NO2 are acidic:
2 SO2 + O2 + 2 H2O  2 H2SO4
4 NO2 + O2 + 2 H2O  4 HNO3
 Processes that produce nonmetal oxide gases
increase the acidity
of the rain
 natural – volcanoes,
bacterial action.
 man-made –
combustion of fuel
Sources of SO2 from Utilities
 Weather patterns
may cause rain
to be acidic in
regions other
than where the
nonmetal oxide
is produced.
Damage from Acid Rain
 Acids react with metals, so acid rain
damages bridges, cars, and other
metallic structures.
 Acid reacts with carbonates, so acid
rain damages buildings and other
structures made of limestone
or cement.
 Acidified lakes affect aquatic life.
 Acid dissolves and leaches
minerals from soil, which
weakens and kills trees.
Acid Rain Legislation
 The 1990 Clean Air Act was passed to reduce
acid rain.
 It requires industries
to minimize pollutant
gas emissions.
 As a result, the acidity
of rain in the northeast
has stabilized and is
beginning to be reduced.