Chapter 4: Chemical Reactions in Aqueous Solutions

Chapter 4:

Chemical Reactions in Aqueous Solutions

Solute Concentrations, Molarity Solution: homogeneous mixture of two or more substances.

Solute: the substance being dissolved.

Solvent: the substance doing the dissolving.

Concentration of a solution: the quantity of a solute in a given quantity of solution (or solvent). A concentrated solution contains a relatively

large amount of solute vs. the solvent (or solution).

A dilute solution contains a relatively small concentration of solute vs. the solvent (or solution).

“Concentrated” and “dilute” aren’t very quantitative .

Molar Concentration

Molarity (M), or molar concentration, is the amount of solute, in moles, per liter of solution:

Formula: Molarity = moles of solute liters of solution

A solution that is 0.35 M sucrose contains 0.35 moles of sucrose in each liter of solution.

Keep in mind that molarity signifies moles of solute per liter of solution, not liters of solvent.

Volumetric Flasks

Weigh 0.01000 mol (1.580 g)

KMnO4.

Dissolve in water. How much water? Doesn’t matter, as long

as we don’t go over a liter.

Add more water to reach the

1.000 liter mark.

Calculating Molaritya. Given moles of solute and volume of

solution

Example 4.1 What is the molarity of a solution prepared by dissolving 5.0 mol of NaCl in enough water to make 2 L of solution?

Sample Problem

b. Given mass of solute and volume of solution

Example 4.2 What is the molarity of a solution prepared by dissolving 50.75 g of AgNO3 in enough water to make 1500 mL of solution?

Sample Problemc. Calculating Mass from Molarity and Volume

Example 4.3 How many grams of sodium carbonate are needed to prepare 0.250 L of a 0.300 M solution?

Why do some solutions conduct electricity?

Arrhenius’s theory: Certain substances dissociate

into cations and anions when dissolved in water.

These ions allow electricity to flow.

Arrhenius’s Theory of Electrolytic Dissociation

Electrolytes dissociate to produce ions.

Review: Ch 2 Ionic

Compounds

The more the electrolyte dissociates,

the more ions it produces.

Types of Electrolytes

A strong electrolyte dissociates completely. A strong electrolyte is

present in solution almost exclusively as ions.

Strong electrolyte solutions are good conductors.

Ex: most ionic compounds

Types of Electrolytes

A weak electrolyte dissociates partially. Weak electrolyte

solutions are poor conductors.

Different weak electrolytes dissociate to different extents.

Ex: weak acids and bases

Nonelectrolytes

A nonelectrolyte does not dissociate. A nonelectrolyte is

present in solution almost exclusively as molecules.

Nonelectrolyte solutions do not conduct electricity.

Ex: sugar, ethanol

Strong electrolytes include: Strong acids & Strong bases Most water-soluble ionic compounds

Weak electrolytes include: Weak acids and weak bases A few ionic compounds

Nonelectrolytes include: Most molecular compounds Most organic compounds (most are molecular)

Is it a strong electrolyte, a weak electrolyte, or a nonelectrolyte?

How do we tell whether an acid

(or base) is weak? Memorize the

strong acids/bases!

Ion Concentrations in Solution

A. Dissolution equations -For a strong electrolyte write cations and anions as individual particles

NaCl(aq)

MgSO4

Na2SO4

Ion Concentrations in SolutionA. Dissolution equations -For a strong electrolyte write cations and anions as individual particles

NaCl(aq) Na+ + Cl-

MgSO4 Mg+2 + SO42-

Na2SO4 2Na+ + SO42-

Multiply molarity (M) x COEFFICIENT of each ion to find the ion concentration!

EXAMPLE:

Na2SO4 2Na+ + SO42-

.010 M 0.020 M 0.010 M

Calculating Ion Concentrations in Solution

Example Calculate the concentrations of all the ions present in the

following solutions:

a. 0.75M solution of NaCl

b. 1.5M solution of AlF3

c. 0.5M solution of Al2(CO3)3

Example

a. 0.75M solution of NaClNaCl(s) Na+ + Cl-

b. 1.5M solution of AlF3

AlF3 Al +3 + 3F-

c. 0.5M solution of Al2(CO3)3

Al2(CO3) 3 2Al +3 + 3CO3 -2

Chemical Reactions in Water

We will look at DOUBLE REPLACEMENT REACTIONS

AX + BY AY + BX

Pb(NOPb(NO33) ) 22(aq) + 2 KI(aq) ----> PbI(aq) + 2 KI(aq) ----> PbI22(s) + 2 KNO(s) + 2 KNO33 (aq) (aq)

The cations change places.The cations change places.

Definition Precipitate: insoluble product (solid formed in

a DR reaction) Example: PbI2 is a precipitate:

Pb(NOPb(NO33) ) 22(aq) + 2 KI(aq) ---> PbI(aq) + 2 KI(aq) ---> PbI22(s) + 2 KNO(s) + 2 KNO33 (aq) (aq)

The “driving force” is the formation of an insoluble compound — a precipitate.

Pb(NO3)2 (aq) + 2 KI (aq)

-----> 2 KNO3(aq) + PbI2(s)

Precipitation Reactions

Solubility Rules

See handout!!

Solubility RulesAre the following compounds soluble or

insoluble?

NaNO3

Ba(OH)2

NaSO4

CaSO4

Use of Solubility Tables—See Reference Booklet

s = soluble -- dissolves in water, is (aq)

i = ss = insoluble or slightly soluble--forms a precipitate, is (s)

Water Solubility of Ionic Compounds

Common minerals are often formed with anions that lead to insolubility:

sulfide fluoride carbonate oxide

Azurite, a copper carbonate

Iron pyrite, a sulfideOrpiment, arsenic sulfide

Net Ionic EquationsNet Ionic EquationsNet Ionic EquationsNet Ionic Equations

Pb(NO3)2 (aq) + 2 KI (aq)

-----> 2 KNO3(aq) + PbI2(s)

Net ionic equation (See handout!)

Pb2+(aq) + 2 I-(aq) ---> PbI2(s)

Sample ProblemPredict whether or not the following pairs of reactants will form precipitates. Then write the double replacement reaction, the full ionic equation, and the net ionic equation.

a. CuCl2 and (NH4)2SO4

b. Ba(NO3)2 and Na2CO3

c. MgCl2 and AgNO3

Acids and Bases

Properties of Acids Properties of Acids Taste sour

React with indicators: turn blue litmus red turn phenolphthalein colorless

React with certain metals to form H2

Corrosive to metals and skin (strong acids)

Neutralize basesto form water and salts

Properties of BasesProperties of Bases

Taste bitter Feel slipperyReact with indicators:

turn red litmus blue turn phenolphthalein pink

Corrosive to skin (strong bases)Neutralize acids to form water and salts

Arrhenius Definitions of Acids and Bases

Acids produce H+ in water solution Bases produce OH- in water solution

Examples: HCl H+ + Cl-

NaOH Na+ + OH-

Arrhenius Definitions: Acids produce H+ in water solutionBases produce OH- in water solution

Arrhenius definitions are limited!!! Not all bases contain OH-

H+ does not exist by itself in

aqueous systems

Definition: Hydronium Ion

In aqueous solution, H+ does NOT exist!

Note: In problems, [H+] = [H3O+]

H+ + H2O H3O+

(hydronium ion)

An acid ----> HAn acid ----> H++ in water in waterAn acid ----> HAn acid ----> H++ in water in water

SIX strong acids areHCl hydrochloric

HBr hydrobromic

HI hydroiodic

H2SO4 sulfuric

HClO4 perchloric

HNO3 nitric

HNOHNO33

STRONG ACIDS = STRONG ELECTROLYTES

Strong Acids/BasesStrong Acids: THERE ARE ONLY SIX!

HCl, HBr, HI, HNO3, HClO4, H2SO4

Strong acids are 100% ionized in solution! (Use single arrow; no equilibrium established)

HNO3 H+ + NO3-

Strong acids are strong electrolytes; completely ionized in water.

In water: HCl(aq) → H+(aq) + Cl–(aq)

No HCl in solution, only H+

and Cl– ions.

Reactions of Acids and Bases:Strong and Weak Acids

Strong Bases

Strong Bases: Group I metals + OH -; some Group II metals + OH –

Some examples: NaOH, LiOH, Sr(OH)2

Strong bases are 100% ionized in solution! (Use single arrow; no equilibrium established)

Sr(OH)2 Sr2+ + 2 OH -

Common Strong Acidsand Strong Bases- memorize!!!

MEMORIZE the strong acids and know how to recognize the strong bases!

Weak acids are Weak acids are not fully ionizednot fully ionized in water. in water.

Weak organic acids contain the –COOWeak organic acids contain the –COOHH group group One of the best known is acetic acid = One of the best known is acetic acid =

CHCH33COOCOOHH

Weak Acids/BasesWeak Acids/BasesWeak Acids/BasesWeak Acids/Bases

Weak bases are Weak bases are not fully ionizednot fully ionized in water in water.(use .(use double arrow; equilibrium established)double arrow; equilibrium established)

One of the best known weak bases is One of the best known weak bases is ammoniaammonia NHNH33(aq) + H(aq) + H22O(liq) O(liq) NHNH44

++(aq) + OH(aq) + OH--(aq)(aq)

Weak Acids/BasesWeak Acids/BasesWeak Acids/BasesWeak Acids/Bases

ACID-BASE THEORIESACID-BASE THEORIESACID-BASE THEORIESACID-BASE THEORIES

The most general theory for common The most general theory for common aqueous acids and bases is the aqueous acids and bases is the BRØNSTED - LOWRY theoryBRØNSTED - LOWRY theory

DEFINITIONS:DEFINITIONS:

ACIDS DONATE HACIDS DONATE H++ IONS IONS

BASES ACCEPT HBASES ACCEPT H++ IONS IONS

The Brønsted definition means NHThe Brønsted definition means NH33 is a is a

BASEBASE in water — and water is itself an in water — and water is itself an ACIDACID

BaseAcidAcidBaseNH4

+ + OH-NH3 + H2OBaseAcidAcidBase

NH4+ + OH-NH3 + H2O

ACID-BASE THEORIESACID-BASE THEORIESACID-BASE THEORIESACID-BASE THEORIES

BRØNSTED - LOWRY theoryBRØNSTED - LOWRY theory::

ACIDS DONATE HACIDS DONATE H++ IONS IONSBASES ACCEPT HBASES ACCEPT H++ IONS IONS

•NHNH3 3 / NH/ NH44++ is a is a conjugate pairconjugate pair — related by — related by

the gain or loss of Hthe gain or loss of H++

•Every acid has a conjugate base, Every acid has a conjugate base, formed formed when Hwhen H++ is removed from the acid. is removed from the acid.•Every base has a conjugate acid, Every base has a conjugate acid, formed formed when Hwhen H++ is added to the base. is added to the base.

Conjugate Pairs

Conjugate PairsConjugate Pairs

Generalized equation:HB (aq) + A (aq) HA (aq) + B- (aq)

Conjugate Pairs

Sample ProblemExample Write the Bronsted Lowry equations

for the weak acid HNO2 and the weak base NH3, identifying the conjugate acid-base pairs in each equilibrium.

Sample Problem Example Write the Bronsted Lowry equations

for the weak acid HCO3 - and the weak base NH3, identifying the conjugate acid-base pairs in each equilibrium.

Some substances can function as both an Some substances can function as both an ACIDACID OR a OR a BASEBASE, depending , depending

on what they are reacted with. (can on what they are reacted with. (can donatedonate OR OR acceptaccept H H++))

They are called amphiprotic or amphotericThey are called amphiprotic or amphoteric

E.g.,

H2O + H2O H3O+ + OH-

Amphiprotic or Amphoteric Amphiprotic or Amphoteric SubstancesSubstances

For any sample of water molecules:For any sample of water molecules:

HH22O (liq) + HO (liq) + H22O (liq) O (liq) HH33OO+ + (aq) + OH(aq) + OH-- (aq) (aq)

KKww = = [H[H33OO++] [OH] [OH--] = 1.00 x 10] = 1.00 x 10-14-14

(at 25 (at 25 ooC)C)

OH-

H3O+

OH-

H3O+

Water dissociation constant, or the ion Water dissociation constant, or the ion product constant of waterproduct constant of water

Neutral, Acid and Basic SolutionNeutral solution:

[H+] = [OH-] = 1.0 x 10-7 M

Acid solution:[H+] > [OH-]; [H+] > 1.0 x 10-7 M

Basic solution:[H+] < [OH-]; [OH-] > 1.0 x 10-7 M

Sample ProblemExample 15.4 A sample of tap water has a [H+]

= 2.8 x 10-6 M. What is the [OH-]?

The pH ScaleThe pH Scale

A common way to express acidity and basicity A common way to express acidity and basicity is with pH (the “power of hydrogen”)is with pH (the “power of hydrogen”)

pH = - log [HpH = - log [H33OO++]]

In a neutral solution, In a neutral solution,

[H[H3OO++] = [OH] = [OH--] = 1.00 x 10] = 1.00 x 10-7-7 at 25 at 25 ooCC

pH = -log (1.00 x 10pH = -log (1.00 x 10-7-7) = - (-7) = 7) = - (-7) = 7

The pH Scale:The pH Scale: pH = - log [HpH = - log [H33OO++]]

Size of pHSize of pH

Basic solution Basic solution pH > 7pH > 7

Neutral Neutral pH = 7 pH = 7

Acidic solutionAcidic solution pH < 7pH < 7

BBig number = ig number = BBasicasic

pH of Common pH of Common SubstancesSubstances

pOH Definition

pOH = - log [OH- log [OH--]]

pH + pOH = 14.0

The pH ScaleThe pH Scale

Sample ProblemExample Calculate the pH, pOH and [OH-] of

an acid solution whose [H+] is 1.8 x 10-4 M.

Strategy:1. Calculate pH using formula and [H+].2. Calculate pOH from pH.3. Can calculate [OH-] from pOH or by using

Kw and [H+].

If the pH of Coke is 3.12, it is acidic.If the pH of Coke is 3.12, it is acidic.

Because pH = - log [HBecause pH = - log [H3OO++] then] then

log [Hlog [H3OO++] = - pH] = - pH

Take antilog and getTake antilog and get

[H[H3OO++] = 10] = 10-pH-pH

[H[H3OO++] = 10] = 10-3.12-3.12 = =

7.6 x 107.6 x 10-4-4 M M

pH to [HpH to [H++] Calculations] Calculations

Strong AcidsStrong Acids dissociate completely in aqueous

solution:

E.g. HCl H+ + Cl-

In a strong acid, [H+] can be calculated from the molarity of the acid.

2.0 M 2.0 M 2.0 M

Sample ProblemExample Calculate the [H+], pH, and [OH-] of a 0.15M

solution of the strong acid, HNO3.

Strategy:

1. Write the disassociation equation and calculate [H+].

2. Calculate pH from [H+].

3. Use Kw to find [OH-].

Strong BasesStrong Bases dissociate completely in aqueous

solution:

E.g. Sr(OH)2 Sr+2 + 2OH-

In a strong base, [OH-] can be calculated from the molarity of the base.

0.5 M 0.5 M 1.0 M

Sample ProblemExample State the pH, pOH, [H+], and [OH-] of a solution made

by dissolving 0.0500 mol of Ba(OH)2 - a strong base - in 5.00 L of water.

Strategy:

1. Calculate molarity of Ba(OH)2.

2. Write disassociation equation; calculate molarity of OH-.

3. Calculate pOH from [OH-].

4. Calculate pH from pOH.

5. Calculate [OH-] from Kw.

Neutralization Reactions and Titration for Strong Acids and Bases

1. Definitions titration titrant indicator endpoint equivalence point

Acid-Base Indicators and Acid-Base Titrations for Strong Acids and Bases

1. Neutralization reactions Neutralization is the (usually complete)

reaction of an ACID + BASE The products of this neutralization are a

“salt” (ionic compound) + water It is a double replacement reaction. Example: HCl + NaOH H2O + NaCl

DefinitionsA titration is a carefully controlled

neutralization reaction.

A buret is used in a titration.

Why? To determine the concentration of an unknown acid or base.

Definitions

The titrant is the substance of known concentration used to determine the unknown concentration of the other substance.

An indicator--substance that changes color at a certain pH—is added to tell us when the neutralization is complete.

Example: Phenolphthalein undergoes a color change between pH 8 and 10

clear in acid Light pink in neutralDark pink in base

DefinitionsThe equivalence point, is the point in the

titration where the neutralization is complete:

[H3O+] = [OH-]The endpoint is the point where the

indicator changes color.If indicator chosen correctly, two points

are identical!

Acid–Base TitrationAcid–Base Titration

Strong acid and base titration

Chemical reaction:

ACID + BASE SALT + WATER

HCl + NaOH NaCl + H2OA double replacement reaction!

In the reaction above, the HCl, NaOH, and NaCl all are strong electrolytes and dissociate completely.

The actual reaction occurs between ions.

Acid–Base Reactions:Example

HCl + NaOH H2O + NaCl

H+ + Cl– + Na+ + OH– H2O + Na+ + Cl–

H+ + OH– H2OA net ionic equation shows the species actually involved in the reaction.

Na+ and Cl– are spectator ions.

Strong acid and base titration

Short-cut equation:

(# H+)MAVA = (# OH-) MBVB

Where #H+ and #OH- are obtained from the formula

Volumes can be in mL or L, but must be same units on both sides

of equation!

Procedure for TitrationProcedure for Titration

Total mols of H+ from the acid

Total mols of H+ from the acid

Total mols of OH- from the base

Total mols of OH- from the base==

*Remember: (conc)(vol in L) = moles

*Remember: (conc)(vol in L) = moles

At the equivalence point (end point):

Sample ProblemExample How many milliliters of 0.0195 M HCl

are required to titrate 25.00 mL of 0.0365 M KOH?

Sample ProblemExample How many milliliters of 0.0108 M

Ba(OH)2 are required to titrate 25.00 mL of 0.0213 M HCl?

Oxidation: Loss of electrons Reduction: Gain of electrons Both oxidation and reduction must occur

simultaneously. A species that loses electrons must lose them to

something else (something that gains them). A species that gains electrons must gain them from

something else (something that loses them). Historical: “oxidation” used to mean “combines with

oxygen”; the modern definition is much more general.

Reactions InvolvingOxidation and Reduction

An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion.

Examples: in NaCl, the oxidation number of Na is +1, that of Cl is –1 (the actual charge).

In CO2 (a molecular compound, no ions) the oxidation number of oxygen is –2, because oxygen as an ion would be expected to have a 2– charge.

The carbon in CO2 has an oxidation number of +4 (Why?)

Oxidation Numbers

In a redox reaction, the oxidation number of a species changes during the reaction.

Oxidation occurs when the oxidation number increases (species loses electrons).

Reduction occurs when the oxidation number decreases (species gains electrons).

If any species is oxidized or reduced in a reaction, that reaction is a redox reaction.

The two processes occur at the same time

Identifying Oxidation–Reduction Reactions

A Redox Reaction: Mg + Cu2+ Mg2+ + Cu

… the products are Cu metal and Mg2+ ions.

Electrons are transferred from Mg metal to Cu2+ ions and …

An oxidizing agent causes another substance to be oxidized.

The oxidizing agent is reduced. A reducing agent causes another substance to be

reduced. The reducing agent is oxidized.

Mg + Cu2+ Mg2+ + Cu

What is the oxidizing agent? What is the reducing agent?

Oxidizing and Reducing Agents

Examples of redox reactions: Displacement of an element by another element Combustion Extraction of metal from ores Metabolic reactions in living organisms Manufacturing chemicals Mg(s) + Ca 2+(aq) Mg 2+ (aq) + Ca(s)

Ox: Mg(s) Mg 2+ (aq) +2e-

Red: Cu 2+ (aq) + 2e- Cu(s)

Redox equations must be balanced according to both mass and electric charge.

For now, our main goals will be to: Identify oxidation–reduction reactions. Balance certain simple redox equations by

inspection. Recognize, in all cases, whether a redox

equation is properly balanced.

Oxidation–Reduction Equations

Neutral Species

For an isolated atom, a molecule, or a formula unit—the oxidation numbers is 0. ex: Cl2, Fe

Monatomic Ions (Groups 1, 2, 17)

Group 1 elements all have an oxidation number of +1

Group 2 elements all have an oxidation number of +2.

Fluorine always has an oxidation number or -1.

Oxygen

In most compounds, oxygen has an oxidation number of –2.

Rules for Assigning Oxidation Numbers

Sum of oxidation numbers for neutral and charged species:Hydrogen +1 when bonded to a nonmetal,

-1 when bonded to a metal

The sum of the oxidation numbers of all atoms (or ions) in a neutral compound = 0.

The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on the polyatomic ion.

ex: CaCl2

Example: Assign oxidation numbers to the elements in the following species:

CaC2O4 Cr2O72-

N2O N2O4

ClO1- ClO41-

6. Identifying redox reactions Look for changes in oxidation numbers

Al (s) + Fe2O3(s) 2 Fe (l) + Al2O3

0 3+ 2- 0 3+ 2-

Example 4.7 Identify the element reduced, the element

oxidized, the reducing agent and the oxidizing agent:

A) Fe2+(aq) + Cr2O72-(aq) + H+(aq)

Fe3+(aq) + Cr3+(aq) + H2O(l)

B) 3 Cl2 (g) + 2 Cr(OH)3 (aq) + 10 OH-

2CrO4-(aq) + 6Cl-(aq) + 8H2O(l)

Applications of Oxidationand Reduction Analytical Chemistry KMnO4 is the most commonly used oxidizing

agent in chemical laboratories.

5Fe2+(aq) + MnO4-(aq) + 8H+(aq) 5Fe3+(aq)+ Mn2+ (aq) + 4H2O

(l)

Oxidation and Reduction in Organic Chem.

Potassium dichromate

Ethanol

Initially the solution turns the orange of Cr2O7

2–

After a while the alcohol is oxidized to a ketone, and the Cr2O7

2– is reduced to Cr3+

In industry: to produce iron, steel, other metals, and consumer goods.

In foods and nutrition: redox reactions “burn” the foods we eat; antioxidants react with undesirable free radicals.

Applications of Oxidationand Reduction

Applications of Oxidationand ReductionEveryday life: to clean (bleach) our clothes, sanitize our swimming pools (“chlorine”), and to whiten teeth (peroxide).