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
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