Chapter Four

1
Chapter Four

• Chemical Reactions in Aqueous Solutions

2
Electrostatic Forces

• Unlike charges ( and ) attract one another.
• Like charges ( and , or and ) repel one
  another.

3
Conduction Illustrated
Electron flow

• Flow of charged particles is electric current.
• One type of current electrons flowing through
  a wire, from cathode (negative electrode) to
  anode (positive electrode).
• Another type of current anions and cations
  moving through a solution as shown here. Cations
  move to the cathode, anions move to the anode.
• Of course, an external source of potential
  (voltage) is required in either case!

4
Arrheniuss Theory OfElectrolytic Dissociation

• Why do some solutions conduct electricity?
• Original hypothesis electricity produced ions
  in solution, and those ions allowed the
  electricity to flow.
• Arrhenius proposed that certain substances
  dissociate into cations and anions when dissolved
  in water.
• The ions already present in solution allow
  electricity to flow.

5
Electrolytic Properties of Aqueous Solutions

• Electrolytes dissociate to produce ions.
• All electrolytes dissociate, but not necessarily
  to the same extent!

6
The Litany Of Electrolytes

• Ionic compounds that dissolve in water dissociate
  into their ions.
• Ionic compounds that dissolve in water dissociate
  into their ions.
• Ionic compounds that dissolve in water dissociate
  into their ions.
• Ionic compounds that dissolve in water dissociate
  into their ions

7
Types Of Electrolytes

• Nonelectrolyte does not dissociate.
• A nonelectrolyte is present in solution almost
  exclusively as molecules.
• Nonelectrolyte solutions do not conduct
  electricity.
• Strong electrolyte dissociates completely.
• A strong electrolyte is present in solution
  almost exclusively as ions.
• Strong electrolyte solutions are good conductors.
• Weak electrolyte dissociates partially.
• Weak electrolyte solutions are poor conductors.
• Different weak electrolytes dissociate to
  different extents.

8
Is it a strong electrolyte, a weak electrolyte,
or a nonelectrolyte?

• Strong electrolytes include
• Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4)
• Strong bases (IA hydroxides, Ba(OH)2)
• 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 of them are
  molecular)

How do we tell whether an acid (or base) is weak?
9
Ion Concentrations in Solution

• Trick question
• What is the concentration of Na2SO4 in a solution
  prepared by diluting 0.010 mol Na2SO4 to 1.00 L?
• The answer is
• zero
• WHY??
• Sohow do we describe the concentration of this
  solution?

10
Calculating Ion Concentrations in Solution

• For 0.010 M Na2SO4
• two moles of Na ions are formed for each mole of
  Na2SO4 in solution, so Na 0.020 M.
• one mole of SO42- ion is formed for each mole of
  Na2SO4 in solution, so SO42- 0.010 M.
• Only one concentration is reported for an ion in
  a solution, even if the ion has two or more
  sources.
• THIMK! A solution is prepared by dissolving 1.5
  mol Na2SO4 and 1.0 mol NaCl in 2.0 L solution.
  What is the sodium ion concentration?

11
Strong And Weak Acids

• Strong acids
• Strong electrolytes completely ionized in water.
• Weak acids
• Some of the dissolved molecules ionize the rest
  remain as molecules (double arrow partial
  ionization).
• Some acids have more
• than one ionizable hydrogen
• atom. They ionize in steps
• (more in Chapter 15).

its all hydrogen ion and chloride ion
Dissolved HCl contains no HCl
Only a little H forms
Dissolved acetic acid is mostly CH3COOH
12
Strong And Weak Bases

• Strong bases
• Most are ionic hydroxides (Group IA and IIA,
  though most IIA hydroxides arent very soluble).
• Weak bases
• Like weak acids, they ionize partially, but most
  of the dissolved base is in the form of
  molecules

Only a little OH- forms
Dissolved ammonia is mostly NH3
13
Common Strong AcidsAnd Strong Bases
14
Acid-Base ReactionsNeutralization

• In the reaction of an acid with a base, the
  identifying characteristics of each cancel out.
• Neutralization is the (usually complete) reaction
  of an acid with a base.
• The products of this neutralization are water and
  a salt.
• Indicators are commonly used to tell when a
  neutralization is complete, or if a solution is
  acidic or basic.

15
Acid-Base Reactions Titrations

• In a titration, two reactants in solution are
  combined carefully until they are in
  stoichiometric proportion (what does that term
  mean?).
• One reactant is placed in a buret and the other
  reactant is placed in a flask along with a few
  drops of an indicator.
• The reactant in the buret (called the titrant) is
  then slowly added to the reactant in the flask
  until the indicator changes color.

16
Titration Techniques
A measured portion of HCl and some indicator
added to a flask
and NaOH is added from the buret, just until
the indicator changes color NaOH and HCl are
present in stoichiometric proportions.
17
Titration Calculations Example
Titration calculations are simply stoichiometry
calculations that involve solutions. They are
NOT a new type of calculation!
18
Titration Calculations (contd)
19
Titration Calculations (contd)
20
Titration Calculations (contd)
21
Some Acid-Base Reactions Form Gases
22
Precipitation Reactions

• There are limits to the amount of a solute that
  will dissolve in a given amount of water.
• If the maximum concentration of solute is less
  than about 0.01 M (or roughly 1 g/100 mL) the
  solute is generally considered to be insoluble in
  water.
• A reaction that forms such a solute is called a
  precipitation reaction the solid that comes out
  of solution is called a precipitate (duh!)
• AgNO3(aq) KCl(aq) ? AgCl(s)
  KNO3(aq)

23
Silver Iodide Precipitation
Silver iodide ppt.
What are the solute species in this solution of
silver nitrate?
What solute species are left over in this
solution?
24
Solubility Rules

• Soluble
• Nitrates, acetates, perchlorates
• Group IA salts and ammonium salts
• Mostly soluble
• Chlorides, bromides, and iodides, except for
  those of Ag, Pb2, Hg22 mercury(I) and Ca2
  (borderline)
• Sulfates, except those of Sr2, Ba2, Pb2, and
  Hg22
• Mostly insoluble
• Carbonates, hydroxides, phosphates, sulfides,
  except as noted under Soluble.

25
Precipitation Reactions ofPractical Importance
26
Net Ionic Equations

• In many reactions (acid-base, precipitation, some
  others) the reacting species are not molecules,
  but ions.
• The equation
• AgNO3(aq) KCl(aq) ? AgCl(s)
  KNO3(aq)
• is misleading because the species that actually
  react arent AgNO3 and KCl. They are ____ and
  ____

27
Net Ionic Equations

• In a net ionic equation, only the species that
  actually react and are formed are shown.
• Spectator ions that sit on the sidelines are
  left out of the equation.
• AgNO3(aq) KCl(aq) ? AgCl(s)
  KNO3(aq)
• Ag NO3 K Cl ? AgCl(s)
  K NO3
• Ag Cl ? AgCl(s) (Net ionic)

Spectator ions
28
To Write a Net Ionic Equation

• Write the balanced equation.
• Identify the strong electrolytes and dissociate
  them (strong electrolytes dissociate
  completely)
• Cancel spectator ions (identical on both sides of
  the equation).
• 2 CH3COOH(aq) Ba(OH)2 (aq) ? Ba(CH3COO)2
  2 H2O(aq)
• 2 CH3COOH Ba2 2 OH ? Ba2 2
  CH3COO 2 H2O
• 2 CH3COOH 2 OH ? 2 CH3COO 2 H2O
  (Net ionic) or
• CH3COOH OH ? CH3COO H2O
  (Net ionic)

29
Oxidation and Reduction

• Oxidation Loss of electrons
• Reduction Gain of electrons
• Both 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.

30
Oxidation Numbers

• An arbitrary bookkeeping device that helps when
  working with oxidation-reduction (redox)
  reactions.
• Can be thought of as The charge an atom would
  have, if it really had a charge.
• Example 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 (since oxygen
  would be expected to have a 2 charge).
• The carbon in CO2 has an oxidation number of 4
  (can you figure out why?)

31
Rules For AssigningOxidation Numbers (in order
of precedence)

• For the atoms in a neutral species, the sum of
  the oxidation numbers is zero.
• In their compounds, the Group 1A metals all have
  an oxidation number of 1 Group 2A metals all
  have an oxidation number of 2.
• In its compounds, the oxidation number of
  fluorine is 1.
• In its compounds, the oxidation number of
  hydrogen is 1.
• In most of its compounds, oxygen has an oxidation
  number of 2.
• In binary compounds with metals, halogens have an
  oxidation number of 1, Group 6A elements have an
  oxidation number of 2, Group 5A elements have an
  oxidation number of 3.

32
Examples of Oxidation Numbers
33
Identifying Oxidation-Reduction Reactions

• In a redox reaction, the oxidation number of a
  species changes during the reaction.
• Oxidation occurs when the oxidation number
  increases (loses electrons).
• Reduction occurs when the oxidation number
  decreases (gains electrons).

34
A Redox Reaction
Electrons are transferred from Mg (light blue)
metal to Cu2 (darker blue) ions. Products are
Cu (red-brown) metal and Mg2 (pink) ions.
Mg Cu2 ? Mg2 Cu
35
Writing And BalancingOxidation-Reduction
Equations

• Must balance both mass and electric charge.
• In aqueous solution, water can be either a
  reactant or a product.
• In acidic solution, H can be either a reactant
  or a product.
• In basic solution, OH can be either a reactant
  or a product.
• In disproportionation reactions, a portion of the
  reactant is oxidized and a portion of that same
  reactant is reduced.

36
Oxidizing Agents and Reducing Agents

• 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.
• What are the oxidizing and reducing agents?
• Mg Cu2 ? Mg2 Cu

37
Oxidation Numbers OfNonmetals

• The maximum oxidation number Group number.
• The minimum oxidation number Group number 8.
• Species at the maximum oxidation number are
  oxidizing agents
• Species at the minimum oxidation number are
  reducing agents.
• Species in the middle of their oxidation number
  range can act as either oxidizing or reducing
  agents.

38
Activity Series of Some Metals
In the activity series, any metal above another
can displace (react with) that lower metal from a
solution of the metal ion.
Example Mg Cu2 ? Mg2 Cu because Mg is
above Cu in the series.
Question Why does a gold coin remain unchanged,
even after centuries underwater?
39
Some Practical Applications ofOxidation and
Reduction

• In analytical chemistry to determine the
  concentration of various species.
• In organic chemistry to convert one functional
  group to another.
• In industry to produce consumer goods.
• Everyday to clean our clothes (bleach) and
  swimming pools (chlorine) and to burn the
  food we eat.
• Combustion is a redox reaction.

40
A Redox Titration
MnO4 (Ox. of Mn is 7)
At endpoint, excess MnO4 colors the solution pink
Fe2
Fe2 Mn2 (and Fe3 formed)
41
Oxidation of Ethanol by Dichromate Ion in Acid
Potassium dichromate
Ethanol has been oxidized to acetic acid.
Cr2O72- (orange) has been reduced to Cr3
(gray-green)
Ethanol
42
Summary

• Soluble ionic compounds are completely
  dissociated into ions in aqueous solution and are
  called strong electrolytes.
• Most molecular compounds exist in solution only
  as molecules and are non-electrolytes.
• Neutralization reactions between acids and bases
  are conveniently represented by ionic and net
  ionic equations.
• Another important type of reaction in solution is
  one in which ions combine to form an insoluble
  solid a precipitate.

43
Summary (continued)

• We use the concept of oxidation numbers to deal
  with a third reaction type known as
  oxidation-reduction (redox).
• Oxidation is an increase in oxidation number
  (loss of electrons) and reduction is a decrease
  in oxidation number (gain of electrons).
• Oxidizing agents oxidize another reactant and are
  themselves reduced. Reducing agents reduce
  another reactant and are themselves oxidized.
• The activity series is a ranking of metals by
  their relative strengths as reducing agents.

44
For each strong electrolyte, write an equation
representing the dissolving process. Write n/a
for those which are not strong electrolytes.

• CaCl2
• HNO3
• CH3COOH, acetic acid
• NaClO4
• H3PO4, phosphoric acid
• Ba(ClO3)2
• FeI3
• K2CO3
• Ni(NO3)2
• Cs2HPO4

• Magnesium chloride
• Ammonia
• HN3, hydrazoic acid
• Zinc sulfate
• Chromium(III) sulfate
• Aluminum nitrate
• Sodium hydrogen carbonate
• Carbon dioxide
• Oxygen difluoride

45
Complete and balance the equation for the
reaction.

• NaOH HNO3
• KOH HClO4
• Barium hydroxide HClO3
• Fe(OH)3 HI
• H3PO4 CsOH
• CH3COOH calcium hydroxide