Transcript Advanced Chemical Reactions

Reaction rates, Equilibrium, Acids/Bases,
Redox Reactions
Measure of disorder or randomness in
a system
 Natural tendency for system to
increase entropy (more random)
 EXAMPLE – Diffusion
◦ As molecules are dispersed, entropy
increases
◦ Continued dispersal leads to a
uniform solution

Remember, things tend towards an
increase in entropy
 Spontaneous reaction favors the
products (exothermic) and releases
free energy
 C + O2  CO2
◦ Exothermic
◦ Solid  gas increases entropy
 Gibbs free energy – max amt of E that
can be used in another process

Entropy never decreases in a system
and instead will increase over time
 UNLESS you change the surroundings
◦ Spraying air freshener
◦ Spray it into a collapsible box

 Study
of reaction rates (rate at
which a chemical reaction takes
place)
 Measured by:
◦ Rate of formation of products
◦ Rate of disappearance of
reactants
◦ Changes in concentration of
reactants or products
 Concentration
 Pressure
 Temperature
 Surface
Area
 All of the above have a DIRECT
relationship
 When
reactants collide
 Normally, molecules bounce off
each other b/c of electron clouds
repulsion
 BUT, if those molecules have a
LARGE amount of energy, they can
overcome the repulsion and react
 Molecules also must collide in the
right orientation

Energy required to start a chemical
reaction
◦ A nudge, a spark
◦ Potential E
 Activated
complex – “speed bump”
of the reaction – point at which it
could go either way
 H2O + CO2  H2CO3  H+ + HCO3-
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Another factor that affects reaction rate
Speeds the reaction by lowering the
activation energy
Not used up by reaction
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Two basic categories for reactions
1. Completion reactions – 1-way
(combustion, decomp, rusting)
2. Reversible reactions – products
can re-form original reactants
Reversible reactions often use 2
arrows b/c reactions occur at the
same time
 Chemical
equilibrium is
DYNAMIC, not STATIC

Chemical equilibrium – reactions in
which the forward and reverse
reaction rates are equal
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Every reaction has a condition of equilibrium
at a given temperature
That means that 2 reactants will react to form
products until a state is reached where the
amounts of products and reactants no longer
change
◦ CO2 in a half-filled, sealed soda bottle
Things will stay that way until the system is
somehow altered
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Equilibrium constant, Keq – a number that
expresses the necessary concentrations of
reactants and products for the reaction to
be at equilibrium
aA + bB  cC + dD
Keq = [C]c [D]d
[A]a [B]b
If Keq >1, the reaction favors the products
If Keq <1, the reaction favors the reactants
Calculate the Keq of the following
equation
CO2 (g) + H2 (g)  CO (g) + H2O (g)
If the [CO2] = 1.5 M, [ H2 ] = 1.5 M,
[ CO ] = 0.6 M, [ H2O] = 0.6 M
 Keq= [CO]1 [H2O]1 = [0.6] [0.6] = 0.16
[CO2]1 [H2]1 [1.5] [1.5]
 So this reaction favors the….
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When a system at equilibrium is
disturbed, the system adjusts in
a way to reduce the change.
Chemical equilibria responds to
3 kinds of stress or change
1. Change in concentration
2. Change in temperature
3. Change in pressure
 Increasing
concentration of
reactant will make the rate of the
forward reaction faster than the
reverse
◦ Called a shift right
◦ Continues until new equilibrium
 H3O+ + HCO3  2H2O + CO2
 Increasing concentration of
product leads to shift left
Remember that endothermic &
exothermic are opposites
 Increasing the temp adds E so the
endothermic will go faster to use it
 If it is exothermic forward, increasing
the temp favors the reactants
 If it is endothermic forward,
increasing the temp favors the
products
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 Only
affects gases
 Imagine volume has been
decreased, increasing the
pressure
 Immediate effect is increase in
concentration of both product &
reactant
 According to principle, system will
adjust to decrease the pressure
A
pressure increase favors the
reaction that produces fewer
molecules (stoichiometry)
 2NOCl   2 NO + Cl2
 H2O + CO   H2 + CO2
Acids – sour taste, conduct electricity
well, react with many metals, generate
hydronium ions (H3O+), turn litmus
paper red
 Bases – bitter taste, slippery feel,
varying solubility, generate hydroxide
ions (OH-), turn litmus paper blue
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Strong acids & bases COMPLETELY
dissociate or ionize in water (one way
reaction)
◦ HNO3 + H2O  H3O+ + NO3◦ NaOH  Na+ + OHWeak acids & bases only partially
dissociate (reversible reaction)
◦ HOCl + H2O  H3O+ + ClO◦ NH3 + H2O  NH4+ + OH-
 Acid
– ionizes to form an H3O+ ion
when added to water
 Base – generate OH- when
dissolved in water
 Acid
– donates a proton (H+) to
another substance
 Base – accepts a proton (H+)
NH3 + H2O  NH4+ + OHH2O is the Bronsted-Lowry acid &
NH3 is the Bronsted-Lowry base
Always reactants
 Conjugate
Acid – Formed when a
base gains a proton (H+)
 Conjugate Base – Formed when an
acid loses a proton (H+)
NH3 + H2O  NH4+ + OH NH4+ is the conjugate acid & OHis the conjugate base
 Always products
 Can
act as an acid or a base
depending on what it is
combined with
 Can
act as a Bronsted-Lowry acid
or base
 H2O + H2O   H3O+ + OH Called the self-ionization of water
 Results in equal concentrations of
H3O+ and OH- in pure water
 [H3O+] = [OH-] = 1.00 x 10-7 M
 [H3O+]
x [OH-] =
1.00 x 10-7 x 1.00 x 10-7 =
1.00 x 10-14
 Found to be true for other
aqueous solutions at equilibrium
 [H3O+] x [OH-] = 1.00 x 10-14
 Also abbreviated as Kw
 Have proportional amounts of
H3O+ & OH [H3O+] x [OH-] = 1.00 x 10-14
H3
O+
OH-
ACID
H3
O+
OH-
NEUTRAL
H3O+
BASE
OH-
 [H3O+]
x [OH-] = 1.00 x 10-14
 If [H3O+] = 1.00 x 10-2, what is
[OH-]?
 [OH-] = 1.00 x 10-12
 If [H3O+] = 1.00 x 10-5, what is
[OH-]?
 [OH-] = 1.00 x 10-9
 1909
– Soren Sorenson – negative
exponents are annoying…
 So let’s just look at the exponents!
 Logarithm – power to which 10
must be raised to equal that number
 log 100 = 2 because 100 = 102
 log 0.001 = -3 because 0.001 =
10-3
 log
 log
 log
 log
 log
10,000 =
0.01 =
10 =
0.000001 =
1=
 Represents
the “power” of
“Hydrogen”
 pH = - log [H3O+]
 What is the pH of a 0.00010 M
solution of HNO3?
 pH = - log [1.0 x 10-4] = -(-4)
=4
 What
is the pH of a 0.2 M solution
of a strong acid?
pH = - log [.2]
 pH = 0.70
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 [H3O+]
 pH
x [OH-] = 1.00 x 10-14
+ pOH = 14
 You can calculate [H3O+] by
1.00 x 10-14 / [OH-]
 Then you can calculate pH
 What
is the pH of a 0.0136 M
solution of KOH, a strong base?
 [H3O+]
 [H3O+]
= 1.00 x 10-14 / 0.0136
= 7.35 x 10-13
 pH = -log [H3O+]
 pH = - log [7.35 x 10-13]
 pH = 12.13
 Lemonade
has a hydronium ion
concentration of 0.0050 moles/L.
What is it’s pH?
 pH = -log [H3O+]
 pH
= 2.3
 What is it’s pOH?
 Reaction
of H3O+ & OH- to form
water molecules and often a salt
 H3O+ & OH-  2H2O
◦ Neutral means [H3O+] = [OH-]
 HCl + NaOH  H2O + NaCl
 Common way to deal with acid &
base spills
 Baking soda = NaHCO3,Ammonia = NH3
Change color at a certain pH level
 Red cabbage juice – changes to blue
between 3 & 4 and to green at 8/9
 Litmus paper – red or blue
 Phenolphthalein – turns bright pink in
the presence of a base

Used to determine the unknown
concentration of a known reactant
 Uses an indicator to show the
equivalence point
 For strong acid/strong base…
 Equivalence point is where [H3O+] =
[OH-] or where moles of acid = moles
of base
 Often uses phenolphthalein
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 Remember
that electronegativity is
a measure of how tightly atoms
hold on to their electrons
 Atoms with large electronegativity
differences form ionic bonds by
electron transfers
 2Na + Cl2  2NaCl
 Can be written as 2Na + Cl2 
2Na+Cl-
 Oxidation =
◦ Na  Na+
 Reduction
Loss of electrons
= Gain of electrons
◦ Cl2  2 Cl These 2 reactions happen
together
 Oxidation-Reduction or REDOX
 OIL RIG
 Use
“oxidation” numbers
 The number of electrons that
must be added or removed to
convert the atom to elemental
or neutral form
 In other words, it’s the charge
the atom would have if it were
an ion
1. Look at the equation
2. Assign known oxidation numbers
3. Calculate unknowns & verify
- Sum of all atoms in a molecule is
zero
- Sum of all atoms in a polyatomic
is equal to the charge on that ion
Uncombined = 0
O2
 Monatomic ion = ion charge Zn 2+
 Flourine = -1 (most electronegative)
 Group 1 = +1
K
 Group 2 = +2
Ca
 Binary compounds – most
electronegative element = ion
charge
CaCl2
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 Hydrogen
usually = +1
◦ If combo with metal, H = -1
 Oxygen usually = -2
◦ If combo with Flourine, O = +2
◦ Can also be -1 in peroxides like H2O2
 Transition
metals have multiple
oxidation states so save them for
last
 S2O72-
 Oxygen
= -2 so O7 = -14
 Entire molecule must = 2 So S2 + (-14) = 2 S2 = +12
 S = +6
 Ca
(OH)2
 Ca = +2
 The entire molecule must = 0
 So (+2) + (OH)2 = 0
 (OH)2 = -2
 O = -2 so O2 = -4
 So -4 + H2 = -2
 H2 = +2
so H = +1
 From
the given, balanced
formulas, assign oxidation
numbers
 2H3O+ + Zn  H2 + 2H2O + Zn2+
 Since
Zn changes from 0 to +2
and some of the H changes from
+1 to 0, it is a redox reaction
If the oxidation number goes UP
during a reaction, it is oxidized
2H3O+ + Zn  H2 + 2H2O + Zn2+
 If the oxidation number goes DOWN
during a reaction, it is a reduction

 1s2
2s2 2p3 – 5 valence electrons,
-3 oxidation number
 1s2 2s2 2p6 3s1 – 1 valence
electron, +1 oxidation number
 1s2 2s2 2p5
 1s2 2s2 2p6 3s2 3p1