Chapter 4: Chemical Reactions in Aqueous Solutions.
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Transcript of 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).