Aqueous StuffAqueous Stuff

Reactions Between Ions Ionic compounds, also called salts, consist of both

positive and negative ions When an ionic compound dissolves in water, it

dissociates to aqueous ions

What happens when we mix aqueous solutions of two different ionic compounds? if two of the ions combine to form a water-insoluble

compound, a precipitate will form otherwise no physical change will be observed

NaCl(s) +Na+(aq)H2O

Cl-(aq)

Reactions Between Ions Example:

suppose we prepare these two aqueous solutions

if we then mix the two solutions, we have four ions present; of these, Ag+ and Cl- react to form AgCl(s) which precipitates

+Ag+(aq) Cl-(aq)

AgCl(s)

NO3-(aq) + Na+(aq) +

+ Na+(aq) + NO3-(aq)

AgNO3(s)H2O

Ag+(aq) + NO3-(aq)Solution 1

NaCl(s)H2O

Na+(aq) + Cl-(aq)Solution 2

Reactions Between Ions we can simplify the equation for the formation

of AgCl by omitting all ions that do not participate in the reaction

the simplified equation is called a net ionic net ionic equationequation; it shows only the ions that react

ions that do not participate in a reaction are called spectator ionsspectator ions

Ag+(aq) Cl-(aq) AgCl(s)+Net ionic equation:

Reactions Between Ions In general, ions in solution react with each other

when one of the following can happen two of them form a compound that is insoluble in water two of them react to form a gas that escapes from the

reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid

an acid neutralizes a base one of the materials can oxidize another

HCO3-(aq) + H3O+(aq) +CO2(g) 2H2O(l)

Bicarbonate ion Carbon dioxide

Reactions Between Ions Following are some generalizations about which ionic

solids are soluble in water and which are insoluble all compounds containing Na+, K+, and NH4

+ are soluble in water

all nitrates (NO3-) and acetates (CH3COO-) are soluble in

water most chlorides (Cl-) and sulfates (SO4

2-) are soluble; exceptions are AgCl, BaSO4, and PbSO4

most carbonates (CO32-), phosphates (PO4

3-), sulfides (S2-), and hydroxides (OH-) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH4OH which are soluble in water

Oxidation-Reduction

Oxidation:Oxidation: the loss of electrons Reduction:Reduction: the gain of electrons Oxidation-reduction (redox) reaction:Oxidation-reduction (redox) reaction: any

reaction in which electrons are transferred from one species to another

Oxidation-Reduction Example: if we put a piece of zinc metal in a beaker containing a

solution of copper(II) sulfate some of the zinc metal dissolves some of the copper ions deposit on the zinc metal the blue color of Cu2+ ions gradually disappears

In this oxidation-reduction reaction zinc metal loses electrons to copper ions

copper ions gain electrons from the zincZn(s) +Zn2+(aq) 2e- Zn is oxidized

Cu(s)+ 2e-Cu2+(aq) Cu2+ is reduced

Oxidation-Reduction

we summarize these oxidation-reduction relationships in this way

electrons flowfrom Zn to Cu2+

+Zn(s) Cu2+(aq) + Cu(s)Zn2+(aq)

loses electrons;is oxidized

gains electrons;is reduced

gives electronsto Cu2+; is thereducing agent

takes electronsfrom Zn; is the oxidizing agent

Oxidation-Reduction Although the definitions of oxidation (loss of

electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others for example, the combustion of methane

An alternative definition of oxidation-reduction is oxidation:oxidation: the gain of oxygen or loss of hydrogen reduction:reduction: the loss of oxygen or gain of hydrogen

CH4(g) + O2(g) CO2(g) + H2O(g)Methane

Oxidation-Reduction using these alternative definitions for the

combustion of methane

CH4(g) + O2(g) CO2(g) + H2O(g)

gains O and losesH; is oxidized

gains H; is reduced

is the reducingagent

is the oxidizingagent

electrons are transferred from carbon to oxygen

Oxidation-Reduction Five important types of redox reactions

combustion:combustion: burning in air. The products of complete combustion of carbon compounds are CO2 and H2O.

respiration:respiration: the process by which living organisms use O2 to oxidize carbon-containing compounds to produce CO2 and H2O. The importance of these reaction is not the CO2 produced, but the energy released.

rusting:rusting: the oxidation of iron to a mixture of iron oxides

bleaching:bleaching: the oxidation of colored compounds to products which are colorless

batteries:batteries: in most cases, the reaction taking place in a battery is a redox-reaction

4Fe(s) +3O2(g) 2Fe2O3(s)

Heat of Reaction In almost all chemical reactions, heat is either given

off or absorbed example: the combustion (oxidation) of carbon liberates

94.0 kcal per mole of carbon oxidized

Heat of reaction:Heat of reaction: the heat given off or absorbed in a chemical reaction exothermic reaction:exothermic reaction: one that gives off heat endothermic reaction:endothermic reaction: one that absorbs heat heat of combustion:heat of combustion: the heat given off in a combustion

reaction; all combustion reactions are exothermic

C(s) + O2(g) CO2(g) + 94.0 kcal/mole C

Properties of Acids & Bases

Reaction with metal oxides strong acids react with metal oxides to give water

plus a salt

2H3O+(aq) + CaO(s) 3H2O(l) + Ca2+(aq)Calciumoxide

Properties of Acids & Bases Reaction with carbonates and bicarbonates

strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O

strong acids react similarly with bicarbonates

2H3O+(aq) + CO32-(aq) H2CO3(aq) + 2H2O(l)

H2CO3(aq) CO2(g) + H2O(l)

2H3O+(aq) + CO32-(aq) CO2(g) + 3H2O(l)

H3O+(aq) + HCO3-(aq) H2CO3(aq) + H2O(l)

H2CO3(aq) CO2(g) + H2O(l)

H3O+(aq) + HCO3-(aq) CO2(g) + 2H2O(l)

Properties of Acids & Bases Reaction with ammonia and amines

any acid stronger than NH4+ is strong enough to react

with NH3 to give a salt

HCl(aq) + NH3(aq) NH4+(aq) + Cl-(aq)

Self-Ionization of Water

pure water contains a very small number of H3O+ ions and OH- ions formed by proton transfer from one water molecule to another

the equilibrium expression for this reaction is

we can treat [H2O] as a constant = 55.5 mol/L

H2O+H2O H3O++OH-

BaseAcid Conjugateacid of H2O

Conjugatebase of H2O

[H2O]2

[H3O+][HO-]Keq =

Self-Ionization of Water combining these constants gives a new constant called the ion ion

product of water, Kproduct of water, Kww

in pure water, the value of Kw is 1.0 x 10-14

this means that in pure water

[H3O+][OH-]Kw = Keq[H2O]2 =

Kw = 1.0 x 10-14

[H3O+]

[OH-]

= 1.0 x 10-7 mol/L

= 1.0 x 10-7 mol/Lin pure water

Self-Ionization of Water

the product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14 for solutions as well.

for example, if we add 0.010 mole of HCl to 1 liter of pure water, it reacts completely with water to give 0.010 mole of H3O+

in this solution, [H3O+] is 0.010 or 1.0 x 10-2

this means that the concentration of hydroxide ion is

[OH-] = 1.0 x 10-14

1.0 x 10-2= 1.0 x 10-12

pH and pOH

we commonly express these concentrations as pH, where

pH = -log [H3O+] we can now state the definitions of acidic and basic

solutions in terms of pH acidic solution:acidic solution: one whose pH is less than 7.0 basic solution:basic solution: one whose pH is greater than 7.0 neutral solution:neutral solution: one whose pH is equal to 7.0

pH and pOH

just as pH is a convenient way to designate the concentration of H3O+, pOH is a convenient way to designate the concentration of OH-

pOH = -log[OH-] the ion product of water, Kw, is 1.0 x 10-14

taking the logarithm of this equation gives

pH + pOH = 14 thus, if we know the pH of an aqueous solution, we can

easily calculate its pOH

Kw = [H3O+][OH-] = 1.0 x 10-14

pH of Salt Solutions

When some salts dissolve in pure water, there is no change in pH from that of pure water

Many salts, however, are acidic or basic and cause a change the pH when they dissolve

We are concerned in this section with basic salts and acidic salts

pH of Salt Solutions

Basic salt: Basic salt: raises the pH as an example of a basic salt is sodium acetate when this salt dissolves in water, it ionizes; Na+ ions

do not react with water, but CH3COO- ions do

the position of equilibrium lies to the left nevertheless, there are enough OH- ions present in

0.10 M sodium acetate to raise the pH to 8.88

OH-CH3COOHH2OCH3COO- + +Acetic acid

(stronger acid)Acetate ion

(weaker baseHydroxide ion(stronger base)

Water(weaker base)

pH of Salt Solutions

Acidic salt:Acidic salt: lowers the pH an example of an acidic salt is ammonium chloride chloride ion does not react with water, but the

ammonium ion does

although the position of this equilibrium lies to the left, there are enough H3O+ ions present to make the solution acidic

NH4+ + H2O NH3 + H3O+

Ammonia(stronger base)

Ammonium ion(weaker acid)

Hydronium ion(stronger acid

Water(weaker base)

Acid-Base Titrations

Titration:Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined

Acid-Base Titrations

An acid-base titration must meet these requirement1. we must know the equation for the reaction so that we

can determine the stoichiometric ratio of reactants to use in our calculations

2. the reaction must be rapid and complete

3. there must be a clear-cut change in a measurable property at the end pointend point (when the reagents have combined exactly)

4. we must have precise measurements of the amount of each reactant

Acid-Base Titrations

As an example, let us use 0.108 M H2SO4 to determine the concentration of a NaOH solution requirement 1:requirement 1: we know the balanced equation

requirement 2:requirement 2: the reaction between H3O+ and OH- is rapid and complete

requirement 3:requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration

requirement 4:requirement 4: we use volumetric glassware

2NaOH(aq)+H2SO4(aq) Na2SO4(aq) + 2H2O(l)(concentration

known)(concentrationnot known)

Acid-Base Titrations experimental measurements

doing the calculations

Trial I

Volume of 0.108 M H2SO4

Volumeof NaOH

25.0 mL 33.48 mLTrial II 25.0 mL 33.46 mL

Trial III 25.0 mL 33.50 mL

average = 33.48 mL

2 mol NaOH1 mol H2SO4

= 0.161 mol NaOHL NaOH

= 0.161 M

mol NaOHL NaOH = 0.108 mol H2SO4

1 L H2SO4x x0.0250 L H2SO4

0.03348 L NaOH

pH Buffers

pH buffer:pH buffer: a solution that resists change in pH when limited amounts of acid or base are added to it a pH buffer as an acid or base “shock absorber” a pH buffer is common called simply a buffer the most common buffers consist of approximately equal

molar amounts of a weak acid and a salt of the conjugate base of the weak acid

for example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer

pH Buffers

How an acetate buffer resists changes in pH if we add a strong acid, such as HCl, added H3O+ ions

react with acetate ions and are removed from solution

if we add a strong base, such as NaOH, added OH- ions react with acetic acid and are removed from solution

CH3COO- H3O+ CH3COOH H2O+ +

CH3COOH OH- CH3COO- H2O+ +

CH3COOH H2O CH3COO- H3O++ +

Added asCH3COOH

Added asCH3COO-Na+

pH Buffers

The effect of a buffer can be quite dramatic consider a phosphate buffer prepared by

dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution

waterpH

0.10 M phosphate buffer7.07.21

2.0 12.07.12 7.30

pH afteraddition of

0.010 mole HCl

pH afteraddition of

0.010 mole NaOH

pH Buffers

Buffer pHBuffer pH if we mix equal molar amounts of a weak acid and

a salt of its conjugate base, the pH of the solution will be equal to the pKa of the weak acid

if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base

pH Buffers

Buffer capacity depends both its pH and its concentration

pH The closer the pH of the buffer is to the pKaof the weak acid, the greater the buffer capacity

Concentration The greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity

Henderson-Hasselbalch Eg. Henderson-Hasselbalch equation:Henderson-Hasselbalch equation: a mathematical

relationship between pH, pKa of the weak acid, HA concentrations HA, and its conjugate base, A-

It is derived in the following way

taking the logarithm of this equation gives

HA H2O A- H3O++ +

[HA]

[A-][H3O+]Ka =

[HA]log [H3O+] + log

[A-]log Ka =

Henderson-Hasselbalch Eg. multiplying through by -1 gives

-log Ka is by definition pKa, and -log [H3O+] is by definition pH; making these substitutions gives

rearranging terms gives

[HA]-log [H3O+] - log

[A-]-log Ka =

[HA]

[A-]+ logpH = pKa Henderson-Hasselbalch Equation

[HA][A-]

pKa = pH - log

Henderson-Hasselbalch Eg.

Example:Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 dissolved in enough water to make 1.0 liter of solution

Henderson-Hasselbalch Eg. Example:Example: what is the pH of a phosphate buffer

solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 in enough water to make one liter of solution

SolutionSolution the equilibrium we are dealing with and its pKa are

substituting these values in the H-H equation gives

H2PO4- H2O HPO4

2- H3O++ + pKa = 7.21

1.0 mol/L 0.50 mol/L

= 7.21 - 0.30 = 6.91

+ logpH = 7.21 0.501.0