Chapter 5 Reactions in Aqueous Solution

Classes of some chemical reactions in solution

Combustion reactions

Gas forming reactions

Dissolution reactions

Precipitation reactions

Acid-base reactions

Oxidation-reduction reactions

Combustion Reactions

3 8 2 2 2C H + 5O 3CO + 4H O

Burning in the presence of oxygen is combustion

In complete combustion all C in the reagents becomes CO2, all H becomes H2O, all N becomes N2

Remaining elements usually combine with oxygen to give the element oxide

In incomplete combustion, C compounds can produce CO or even elemental C

Ex) Combustion of propane

3 8 2 2C H + 7/2 O 3CO + 4H O

The formation of soot is the result of incomplete combustion

Gas Forming Reactions

+2 2 3(g) (l) (aq) (aq)Cl + H O Cl +H O

s (aq) ( ) (g) (aq) (l)3 2 2 3 2 3 2 2

2CH CO H + Na CO CO + 2NaCH CO + H O

(s) (g) (s)2 3

CaO + CO CaCO

Gas can be produced

Ex) The production of carbon dioxide when washing soda and vinegar are mixed

Gases also be consumed.

2 2(l) (g) (l)2HOOH O +2 H O

Ex) Decomposition of Hydrogen peroxide

Ex) Formation of limestone from lime

Ex) Formation of hydrochloric acid

Note : The inorganic gas produced forms acids when combined with water

CO2 (g) + H2O (l) → H2CO3 (aq)

Carbonic acid

SO2(g) + H2O (l)→ H2SO3(aq)

Sulfurous acid

SolubilityThe amount of a substance that can combine with another to give a single phase.

It is governed by the polarity of the two substances involved.

“like dissolves like”

A polar solvent will dissolve polar molecules

Ex) Salt or sugar in Water

A non-polar solvent will dissolve non-polar molecules

Ex) Spices in cooking oil

When the interactions between the solute and solvent molecules are similar to those between solvent molecules and between solute molecules they will mix.

The major component of a single-phase mixture is called the solvent, while the minor component is called the solute

In water these interactions cause water molecules to solvate, i.e. surround, the solute molecule or ion.

Solubility

The partial negative charge on oxygen in water will interact with cations

The partial positive charge on hydrogen in water will interact with anions

When two substances mix, and dissolve into each other they are referred to as miscible

When two substance don’t have compatible polarities they will not mix.

They are referred to as immiscible

Interactions between the different molecules will be weaker than between the same molecules for at least one component.

Ex) Oil and water

SolubilityThe solubility of a substance is usually expressed in the number of grams that dissolves in 100 ml of the solvent at a specified temperature

2H O2 2 7 2 2 7(s) (aq)K Cr O K Cr O

A dissolution reaction can be represented by a chemical equation

Dissolution can be accompanied by dissociation, either partial or complete:

2H O + 2-2 2 7 2 7(s) (aq) (aq)K Cr O 2K + Cr O

Solubility is expressed using its solubility product (Ksp):

AnBm (s) + H2O (l) → n Am+ (aq) + m Bn-

(aq) Ksp = [Am+]n[Bn-]

m

Generally Ksp below 10-6 indicates that a salt is insoluble in water

Large Ksp → Soluble Small Ksp → insoluble

strong electrolyte is one which dissociates completely into cation(s) and anion(s) when dissolved in water.

ElectrolyteAny compound that generates ions when dissolved in water.

weak electrolyte is one which dissociates only partially into cation(s) and anion(s) when dissolved in water.

The undissociated molecules of compound remain neutral.

nonelectrolyte is a compound that dissolves in water but does not dissociate into ions.

Ex) Ionic Salts AnBm (s) + H2O (l) → n Am+ (aq) + m B

n- (aq)

Ex) Ammonia NH3(g) + H2O (l) → NH4+ (aq) + OH

- (aq)

Ex) Vinegar CH3COOH (l) + H2O(l) → CH3COO-(aq) + H3O+(aq)

Table salt NaCl (s) + H2O (l) → Na+(aq) + Cl-(aq)

Large Ksp

Small “Ksp”

Ksp ≈ 0

Which of the following ionic compounds are likely to be soluble in water?(a) LiNO3 (b) Cu3PO4 (c) MgCO3

(d) NiSO4 (e) (NH4)2CO3 (f) Fe(ClO3)3

Solubility of Ionic Compounds

Most Halides are soluble

Alkaline (gr. 1) and ammonium salts are soluble

Hydroxides, phosphates,Sulfites and carbonates tend to be insoluble except for their alkali salts

Note that these anions occur in minerals (rocks).

Nitrates, acetates, chlorates and perchlorates are soluble

YY

NY

NY

The higher the charge of ion the less likely its soluble

Precipitation ReactionsCompounds that exceed their solubility commonly precipitate from solution

The product of a reaction between two soluble salts can be insoluble, resulting in the product to precipitate from solution.

(aq) (aq) (s) (aq)2 2 4 4BaCl +K SO BaSO + 2KCl

2-2+ - +(aq) (aq) (aq) (aq)4

+ -(s) (aq) (aq)4

Ba + 2Cl + 2K + SO

BaSO + 2K + 2Cl

Ionic equation

Spectator ions, K+ and Cl-, are ignored giving the Net Ionic Equation

2-2+(aq) (aq) (s)4 4Ba + SO BaSO

Exercise

Write a net ionic equation for each of the following reactions, and identify the spectator ions:

(a) Mixing potassium hydroxide & copper(II) sulfate solutions

(b) Mixing ammonium phosphate & barium hydroxide solutions

(c) Mixing silver fluoride & magnesium iodide solutions

KOH (aq) + CuSO4 (aq) → Cu(OH)2(s) + K2SO4 (aq)2

2 K+ (aq) + 2 OH_ (aq) + Cu2+(aq) + SO4

2-(aq) → Cu(OH)2(s) +

2 K+ (aq) + SO4

2- (aq)X X X X

2 OH_ (aq) + Cu2+(aq) → Cu(OH)2(s)

(NH4)3PO4 (aq) + Ba(OH)2 (aq) → NH4OH (aq) + Ba3(PO4)2 (s)

6 NH4+

(aq) + 2 PO43- (aq)

2 3 6

+ 3 Ba2+ (aq) + 6 OH_ (aq) → Ba3(PO4)2 (s) + 6 NH4

+ (aq)

+ 6 OH-(aq) X X X

X 2 PO4

3- (aq) + 3 Ba2+ (aq) → Ba3(PO4)2 (s)

AgF (s) + MgI2 (aq) → AgI (s) + MgF2 (aq)

2 Ag+(aq) + 2 F_(aq) + Mg2+(aq) + 2 I

_

(aq) → 2 AgI (s) + Mg2+(aq) + 2 F_

(aq) X X X X

2 2

Ag+(aq) + I_

(aq) → AgI (s)

Acids and BasesThree major ways to define acids and bases introduced by Lewis, Brønsted and Arrhenius.

They differ in the role of water

Arrhenius and Brønsted require water, Lewis does not

BrønstedAcid Donates an H+

Base Accepts an H+

HCl + H2O → H3O+ + Cl- Acid

NaOH + H+ → Na+ + H2O Base

Arrhenius Acid Produces H3O+ when added to water

Base Produces OH- when added to water

Acids and Bases

NH3+ H2O → NH4+ + OH-

HCl + H2O → H3O+ + Cl-

Acids and Bases

LewisAcid Accepts electrons

Donates electronsBase Note: Electrons are not transferred between acids and bases, they are shared.

BH3 + NH3 → BH3NH3

Acid Base

BH H

HN H

H

H

: B N H

H

H

H

H

H

Acids and BasesThe Lewis definition is the most general

Consider a Brønsted Acid It donates a H+ H+ leaves electrons behind

A-H A : _

H+ i.e. A accepts the electrons

A is a Lewis Acid

All Brønsted acids are Lewis Acids

An Arrhenius acid, is a Bronsted Acid, since it produces H3O+ when dissolved in water as it “donates” H+ to H2O.

A strong acid, just like a strong electrolyte, is an acid which dissociates completely when dissolved in water.

Acids and Bases

The concentration of H3O+ is thereby the highest possible, determined exactly by how much acid was added to water

Ex) HCl + H2O → H3O+ + Cl-

Ex) H2SO4 + H2O → H3O+ + HSO4

_

Inorganic acids tend to be strong acids (except HF)

A strong base, just like a strong electrolyte, is dissociates completely when dissolved in water.

The concentration of OH_ is thereby the highest possible, determined exactly by

how much acid was added to water.

Ex) NaOH → Na+ + OH_

The hydroxides of alkali metals are strong bases.

Acids and BasesA weak acid, just like a weak electrolyte, does not dissociate completely when dissolved in water

The concentration of H3O+ is not the highest possible, since much remains in the undissociated form.

Ex) CH3COOH (l) + H2O(l) → CH3COO-(aq) + H3O+(aq)

The concentration of H3O+ is determined from the dissociation constant, similar to Ksp, and the amount of acid added.

Organic acid tend to be weak acids

A weak base, just like a weak electrolyte, does not dissociate completely when dissolved in water.

The concentration of OH- is not the highest possible, since much remains in the undissociated form.

The concentration of OH- is determined from the dissociation constant, similar to Ksp, and the amount of acid added.

Ex) NH3+ H2O → NH4+ + OH-

Metal Oxides and nitrogen containing organic compounds tend to be weak bases

(l) (s) (aq)2 4 10 3 46H O +P O 4H PO

(s) (l) (aq)2 2Na O +H O 2NaOH

(s) (aq) (aq) (l)2 3 3 2Al O + 6HCl 2 AlCl + 3H O

(s) (aq) (l) (aq)2 3 2 4Al O + 2NaOH + 3H O 2Na[Al(OH) ]

Non-metal oxides react with water to give oxoacids

Amphoteric oxides usually do not dissolve with water by themselves, but react with both strong acids and strong bases to give soluble products

The oxides are anhydrides

Metal oxides react with water to give hydroxide bases

Therefore metal oxides are anhydrides of bases

Therefore non-metal oxides are anhydrides of acids

Metalloid oxides are amphoteric: they react with either strong acids or strong bases

The oxides are anhydrides

1 2 13 14 15 16 17

Strength of acids and bases is correlated with the positions of the oxides on the PT

Metal Metalliod Non-metal

pHThe acidity (or basicity) of a solution is reported as pH:

pH = -log [H3O+] or [H3O+] = 10-pH

p = power of

= concentration of H3O+

in mol./l = molar (M)

For pH < 7 solution is acidic

For pH > 7 solution is basic

Ex) 0.10 M solution of HCl [H3O+] = 0.10 M

pH = - log [0.10] = -(-1.00) = 1.00 # of sig. figs. Increased from 2 to 3?

For logarithmic quantities only the decimal numbers are significant.

Therefore a pH = 1.00 has only 2 sig. figs,

Note: pH does not have units

pOHBasicity of a solution can be reported as pOH:

pOH = -log [OH-] Where [OH-] = conc. of OH in mol/l

pH and pOH are related by: pH + pOH = 14 at 25 oC

Therefore an acidic solution as pOH > 7, and a basic solution has pOH < 7

Exercise Determine the pH and pOH of:

a) 0.275 M HNO3 solution

[H3O+] = 0.275 M pH = - log (0.275) = -(0.561) = 0.561

pOH = 14.000 – pH = 14.000 -0.561 = 13.439

b) 0.0051 M NaOH solution

[OH-] = 0.0051 M pOH = - log (0.0051) = -(-2.29) = 2.29

pH = 14.000 – pOH = 14.000 -2.29 = 11.71

Reactivity

A reaction between an acid and a base produces water and a salt

HCl(aq) + NaOH(aq) → H2O (l) + NaCl (aq)

H3O+(aq) + Cl-(aq) + Na+(aq) + OH_ (aq) → H2O(l) + Na+(aq) + Cl-(aq)

H3O+(aq) + OH_(aq) → H2O(l)

Acid Base Water Salt

Ionic Equation

Net equation

A strong acid will react completely with any base.

A strong base will react completely with any acid.

A reaction between a weak acid and a weak base will not go to completion unless there is a driving force (e.g. making a gas):

Reactivity

(a) KOH(aq) + H2SO4(aq) →

(b) H3PO4(aq) + LiOH(aq) →

Write an ionic equation for each of the following acid-base reactions:

Ex) CH3COOH (aq) + NH3 (aq) → CH3COONH4(aq) + H2O(l)

Reduction-Oxidation (REDOX) ReactionsIn chemical reaction bonds, both covalent and ionic, are made and broken by moving electrons.

For all reaction dealt with to date the number of electrons on each atom has been preserved

In REDOX reaction electrons are transferred between atoms as well has bonds being broken and formed.

We therefore need to keep track of the number of electrons of each atom.

The oxidation state of an atom is used for this purpose, which is related to the idea of the formal charge.

2 Na + O2 → 2 NaO 2 Na+ + O2-

The balance between ionic/covalency of the bonds in the reactants is different from that in the product.

0 0 1+ 2-

Na → Na+ + 1 e ½ O2 + 2 e → O2-

Oxidized Reduced

Loss e’s Oxidation = LEO

Gain e’s Reduction = GER

Rules for Assigning Oxidation States

Oxidation StateThe oxidation state of an element is its charge assuming ionic bonding

Electrons are not shared they are placed on the more electronegative element

Pure elements have oxidation states of 0

Ions have oxidation states that add up to the charge of the ion

Hydrogen has an oxidation state of +1 unless bonded to a less electronegative atom. When bound to metals or boron it has an oxidation state of -1.

Fluorine has an oxidation state of -1

Oxygen has an oxidation state of -2 unless bonded to fluorine or another oxygen.

Halogens other than fluorine have oxidation states of -1 unless bonded to oxygen or a more electronegative halogen

The rest are determined by the process of elimination, where the oxidation states must add up to the total charge of the molecule or ion.

Oxidation State

Determine the oxidation state of all the element in the following molecules:

a) H2 H = 0

b) CO2O = -2 0 = C + 2(-2)

c) BF3 F= -1 0 = B + 3(-1)

d) H2SO4H = 1 O = -2 0 = S +(2(+1) + 4(-2)) = S - 6

e) SO2ClF O = -2 F = -1 Cl = -1 0 = S +(2(-2) -1 -1) = S - 6

f) IO2F2

_ O = -2 F = -1 -1 = I + (2(-2) + 2(-1)) = I - 6

C = 4

B = 3

S = 6

S = 6

g) HPO42- O = -2 H = 1

I = 5

-2 = P + (4(-2) + 1) = P - 7 P = 5

Oxidation States of Poly Atomic Ions

PO43- SO4

2-

NO3-

ClO3-

CO32-

NO2-

PO33- SO3

2- ClO2-

4 5

3

5

3

6

4

5

3

Recognizing Redox Reactions

3

3 2 2 4 2

(s) (aq)

(aq) (g) (l)

Cu + 4HNO

Cu(NO ) +N O + 2H O

Consider the following reaction:

Cu = 0

HNO3 N = 5

Reactants Products

Cu2+ = 2

N2O4 N = 4

Cu → Cu2+ + 2 e Oxidation

4 HNO3 + 2 e → 2NO3

- + N2O4 + 2 H2O Reduction

# of e’s gained = # of e’s lost

REDOX equation seem difficult to balance

Recognizing Redox ReactionsThe species that is oxidized in a redox reaction is called the reducing agent

The species that is reduced in a redox reaction is called the oxidizing agent

In the previous example Cu is the reducing agent

In the previous example HNO3 is the oxidizing agent

Concepts from Chapter 5Electrolytes

Solubility Role of PolarityKsp

Miscible vs. Immiscible

Reaction Types

Precipitation Reactions

Gas Forming Reactions

Combustion Reactions

Acid-Base Reactions

Redox Reactions

Lewis vs. Brönsted vs. Arrhenius Definitions of Acids/Bases

Strong Acids/Bases vs. Weak Acids/Bases

Acidity/Basicity of Oxides

pH and pOH

Assigning Oxidation States

Net Ionic Equations