Reversible ReactionsA 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: 22HgO( ) 2Hg( ) O ( )s l + g

Dynamic EquilibriumA reversible reaction reaches dynamic equilibrium when therate of its forward reaction equalsthe 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 EqualAt 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 (productsare favored.)

Others reach equilibrium when only a small percentage of reactants are consumed (reactants are favored.)

Le Châtelier’s PrincipleLe 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:ConcentrationVolume/PressureTemperature

If you increase the concentration byadding more of areactant 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.

When you increase the pressure (usually by decreasing volume of the container), the system shifts so the least numberof gas molecules are formed (less collisions= lower pressure.)

When you decrease thepressure, the system willshift so the greatest numberof gas molecules are formed.

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 temperatureis decreased,the exothermicreaction will befavored (produces heat.)

2 SO2(g) + O2(g) 2 SO3(g) H° = -198 kJ How will the reversible reaction above shift in response to each of the following stresses?

• adding more O2 to the container

• condensing and removing SO3

• compressing the gases

• cooling the container

• doubling the volume of the container

• warming the mixture

Shift right Shift right Shift right Shift right Shift left

Shift left

The Law of Mass ActionThe 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:

Lowercase letters represent coefficients.Always products over reactants.Pure solids and pure liquids are not included.

K =

[C]c [D]d[A]a [B]b

The Equilibrium ConstantThe 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 KK > 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 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.

2. Calculate the value of Kc for this reaction

at 445oC.

K =

[HI]2[H2][I2]

Kc =

[HI]2[H2][I2]

=

[0.934]2[0.03

3]

=

49.9 [0.5

3]

The Reaction QuotientWhen 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:

The non-equilibrium concentrations (or pressures) are used.

Q =

[C]c [D]d[A]a [B]b

Q vs. KWe 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.)

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:

Then, compare it with K:

Q =

(ICl)2(I2)(Cl2)

=

(0.355)2(0.11

4)

=

10.8 (0.10

2)

10.8

81.9

Q K < Reaction will proceed to the right

Properties of AcidsTaste 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 BasesAn 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.

HCl ionizes in water,producing H+ and Cl– ions

NaOH dissociates in water,producing Na+ and OH– ions

Hydronium IonThe 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 BasesIn 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 PairsBrønsted-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.

Identify the Brønsted-Lowry Acids and Bases and Their Conjugates in the Reactions below:

a.

b.

H2SO4 + H2O HSO4– + H3O+

HCO3– + H2O H2CO3 + HO–

acid

conjugate base

base

conjugate acid

base

conjugate acid

acid

conjugate base

Amphoteric CompoundsAn amphoteric substance is one that can react as either an acid or a base.

Example: waterWater can act as an acid.

Water can act as a base.

g + l aq aq–3 2 4NH ( ) H O( ) NH ( ) OH ( )

base acid

aq + l aq + aq–2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )

acid base

Visual Concept

Polyprotic AcidsMolecules 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 automaticallymakes removing the second H harder.(Ex: H2SO4 is a stronger acid than HSO4

)

Lewis Acids and BasesA 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 BasesStrong acid – fully

dissociates in water (almostevery molecule breaks up to form H+ ions.

Weak acid – partially dissociates in water.

Strong base – fully dissociates in water (almostevery molecule breaks up to form OH- ions.

Weak base – partially dissociates in water.

A Strong Acid

A Weak Acid

Ionization Constants for Weak Acids and BasesThe acid ionization constant (Ka) is

the equilibrium constant for the ionization reaction of a weak acid:

A base ionization constant (Kb) can also be created for the ionization reaction of a weak base:

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

Ka = [H3O+] [A-]

[HA]

B(aq) + H2O(l) BH+(aq) + OH-

(aq)Kb =

[BH+] [OH-][B]

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)

•For every H3O+ produced, there is also a C2H3O2-

produced, 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

=

[H3O+][C2H3O2

-][HC2H3O2]=

[1.34 x 10-

3M][1.34 x 10-

3M] =

1.82 x 10-5[0.0987

M]

Ka & Kb and Strength of Acids/BasesThe strength of an acid or base is

measured by the size of its equilibrium constant when it reacts with H2O.

Six Strong Acids and BasesBecause 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

Strengths ofConjugate Pairs

the stronger an acid is at donating H, the weaker the conjugate base is at accepting H. (i.e. strong acidshave weak conjugatebases, and weak acidshave strong conjugatebases.)

Acids Conjugate Bases HClO4 ClO4

-1 H2SO4 HSO4

-1

HI I-1

HBr Br-1

HCl Cl-1

HNO3 NO3-1

H3O+1 H2O

HSO4-1 SO4

-2

H2SO3 HSO3-1

H3PO4 H2PO4-1

HNO2 NO2-1

HF F-1

HC2H3O2 C2H3O2-1

H2CO3 HCO3-1

H2S HS-1

NH4+1 NH3

HCN CN-1

HCO3-1 CO3

-2

HS-1 S-2

H2O OH-1

CH3-C(O)-CH3 CH3-C(O)-CH2-1

NH3 NH2-1

CH4 CH3-1

OH-1 O-2

Incr

easi

ng A

cidi

ty

Increasing Basicity

Predicting Acid StrengthThe strength of an acid depends

on its tendency to ionize (let go of its hydrogen.)

For binary acids, the strength of the acid dependson 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 BEARAll of these increase in the direction toward their letter:

B=BasicityE=Electronegativity, ionization Energy, & Electron AffinityA=AcidityR=Radius

B E

AR

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 atomsbonded to Y, thestronger the acid.

Predict the relative strengths of the following acids:

• HCl, HBr, and HI

• HClO, HBrO, and HIO

• HNO3 and HNO2

Predict the relative strengths of the following bases:

• Cl-, Br- and I-

• H2PO3- and H2PO4

-

• LiOH and Mg(OH)2

HCl < HBr < HIHIO < HBrO <

HClOHNO2 < HNO3

I- < Br- < Cl-

H2PO4- <

H2PO3-

LiOH < Mg(OH)2

Autoionization of WaterWater 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 WaterThe product of the H3O+ and OH–

concentrations is always the same number, called the ion product of water (Kw).

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.

Kw =

[H3O+] [OH-]

= 1 x 10-14

Acidic and Basic SolutionsNeutral 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

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-]:

Then, compare [H3O+] with [OH-]:

Kw =

[H3O+][OH-]Kw =

1.0 x 10-141.5 x 10-9

=

6.7 x 10-6 M

1.5 x 10-9 M

6.7 x 10-6 M

[H3O+] [OH-] <

The solution is basic

[OH-] = [H3O+

]

The pH ScaleThe 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

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+]:

Then, calculate pH:

Kw =

[H3O+][OH-]Kw =

1.0 x 10-141.3 x 10-2

=

7.7 x 10-

13 M

= -log(7.7 x 10-13 M)

12.1

pH = −log

[H3O+]pH = The solution is basic

[H3O+] = [OH-

]

A 0.100 M weak acid (HA) solution has a pH

of 4.25. Find Ka for the acid.

First calculate [H3O+]:

Then, calculate Ka:

pH =

-log [H3O+]

=

[H3O+] 10-pH=

10-

4.25=

5.6 x 10-5 M

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

Ka = [H3O+] [A-][HA]

=

(5.6 x 10-5)(0.100 - 5.6 x

10-5)

(5.6 x 10-5) =

3.1 x 10-8

The pOH ScaleAnother 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 pOHThe 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 SaltsSalts 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 basicEx: 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.

BuffersBuffers are solutions

that resist changes in pH when an acid or baseis added.

Buffers contain both aweak acid and itsconjugate base (or aweak base and itsconjugate acid.)

The weak acid neutralizes added base.The conjugate base neutralizes added acid.

Buffering EffectivenessA 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. The absolute concentrations of acid and base.

Buffers in Human BloodMany of the chemical reactions that

occur in the body are pH-dependent.

Ideally, the pH of the blood shouldbe 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 IndicatorsAn 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 usedto visually "indicate" the approximate pH of a sample.

Litmus PaperA common indicator is found

on litmus paper. It is red below pH 4.5 and blue above pH 8.2.

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 BlueBromothymol 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.

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 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)

analyte

Titration ProcedureIn 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 beenadded to neutralize all the acid,the indicator changes color.

The difficulty is determiningwhen there has been just enough titrant added to complete the reaction…without going over!

The Equivalence PointEnd point - The point

at which an indicator changes color.

Equivalence point – The point at which themoles 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 TitrationSample ProblemThe 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 concentration of the unknown HCl solution? HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) 12.54 mL

NaOHmL NaOH

L NaOH

1000

1

L NaOH

mol NaOH

1

0.100mol NaOH

mol HCl

11

=1.25 x 10-3 mol HCl

First, write the neutralization reaction

L

moles=M 0.010000 L

HCl

=1.25 x 10-3 mol HCl

= 0.125 M HCl

Titration of Strong Acid with Strong BaseSince 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 AcidThis 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 BaseThe 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 AcidSalt 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 AcidA 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 CurvesSample ProblemTwo 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 RainRain 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 witha pH < 5.6 is called acid rain.

Acid rain is linked to damagein ecosystems andstructures.

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 rainnatural – volcanoes,

bacterial action.man-made –

combustion of fuel

Weather patternsmay cause rain to be acidic in regions other than where thenonmetal oxideis produced.

Damage from Acid RainAcids 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 LegislationThe 1990 Clean Air Act was passed to

reduce acid rain.It requires industries

to minimize pollutantgas emissions.

As a result, the acidityof rain in the northeasthas stabilized and is beginning to be reduced.