Transcript chapter 4 lecture slides

Reactions in Aqueous Solution
Chapter 4
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Reactions Between Ions in Aqueous Solutions
General properties of aqueous solution
A solution is a homogenous mixture of 2 or more
substances = solute dissovled in solvent
The solute is(are) the substance(s) present in the
smaller amount(s)
The solvent is the substance present in the larger
amount
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g)
N2
O2, Ar, CH4
4.1
Solute
An electrolyte is a substance that, when dissolved in
water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved,
results in a solution that does not conduct electricity.
nonelectrolyte
weak electrolyte
strong electrolyte
4.1
Strong and Weak Electrolytes
Conduct electricity in solution?
Cations (+) and Anions (-)
Strong Electrolyte – 100% dissociation
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
Weak Electrolyte – not completely dissociated
CH3COOH
CH3COO- (aq) + H+ (aq)
Acetic acid is a weak electrolyte because its
ionization in water is incomplete.
A reversible reaction. The reaction can
occur in both directions.
4.1
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s)
H 2O
C6H12O6 (aq)
How to Predict Electrolytes
• Water soluble and ionic = strong electrolyte (probably)
• Water soluble and molecular, and a weak acid or
weak base = weak electrolyte
• Otherwise, the compound is probably a nonelectrolyte.
4.1
Metathesis Reactions
Double Displacement Reactions
• cations(A and C) & anions (B and D) change partners
in the reaction
AB + CD –> AD + CB
Example:
• Pb(NO3)2(aq) + 2KI(aq) –> 2KNO3(aq) + PbI2(s)
Precipitation of Lead Iodide
PbI2
Precipitation Reactions
Precipitate – insoluble solid that separates from solution
4.2
Solubility is the maximum amount of solute that will dissolve
in a given quantity of solvent at a specific temperature.
4.2
Molecular equation, Ionic equation
and net Ionic equation
molecular equation
precipitate
Pb(NO3)2 (aq) + 2NaI (aq)
PbI2 (s) + 2NaNO3 (aq)
ionic equation
Pb2+ + 2NO3- + 2Na+ + 2I-
PbI2 (s) + 2Na+ + 2NO3-
net ionic equation
Pb2+ + 2I-
PbI2 (s)
Na+ and NO3- are spectator ions: ions that are not
involved in the overall reaction
4.2
Writing Net Ionic Equations
1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes
completely dissociated into cations and anions. Weak and
non electrolytes are written as molecules
3. Cancel the spectator ions on both sides of the ionic equation
4. Check that charges and number of atoms are balanced in the
net ionic equation
Write the net ionic equation for the reaction of silver
nitrate with sodium chloride.
AgNO3 (aq) + NaCl (aq)
AgCl (s) + NaNO3 (aq)
Ag+ + NO3- + Na+ + Cl-
AgCl (s) + Na+ + NO3-
Ag+ + Cl-
AgCl (s)
4.2
The net ionic equation tells us:
1. changes in ionic strength
(more ions present before the rxn than after)
2. what actually changed during a reaction
Example: Cd2+ (aq) + S2-(aq) –> CdS (s)
Writing ionic equations, ask:
1. is substance soluble ?
2. is substance a strong electrolyte?
**If yes to both questions, write substance as ions.
3. Weak and non electrolytes are written as molecules.
Conditions that favor product
formation:
• formation of a precipitate
• formation of a soluble weak electrolyte
• formation of a nonelectrolyte
• formation of a gas
• oxidation-reduction reactions
Acids-Base Reactions
Acids
*Have a sour taste. Vinegar owes its taste to acetic acid.
*Cause color changes in plant dyes.
*React with certain metals to produce hydrogen gas.
2HCl (aq) + Mg (s)
MgCl2 (aq) + H2 (g)
*React with carbonates and bicarbonates to produce carbon
dioxide gas
2HCl (aq) + CaCO3 (s)
CaCl2 (aq) + CO2 (g) + H2O (l)
*Aqueous acid solutions conduct electricity.
4.3
Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
Cause color changes in plant dyes.
Aqueous base solutions conduct electricity.
4.3
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
Arrhennius’s definition can be applied only to aqueous solution.
4.3
A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
base
acid
acid
base
A Brønsted acid must contain at least one
ionizable proton!
4.3
Identify each of the following species as a Brønsted acid,
base, or both. (a) HI, (b) CH3COO-, (c) H2PO4HI (aq)
H+ (aq) + I- (aq)
CH3COO- (aq) + H+ (aq)
H2PO4- (aq)
Brønsted acid
CH3COOH (aq)
H+ (aq) + HPO42- (aq)
H2PO4- (aq) + H+ (aq)
H3PO4 (aq)
Brønsted base
Brønsted acid
Brønsted base
Amphoteric species
4.3
Monoprotic acids
HCl
H+ + Cl-
HNO3
H+ + NO3H+ + CH3COO-
CH3COOH
Strong electrolyte, strong acid
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Diprotic acids
H2SO4
H+ + HSO4-
Strong electrolyte, strong acid
HSO4-
H+ + SO42-
Weak electrolyte, weak acid
Triprotic acids
H3PO4
H2PO4HPO42-
H+ + H2PO4H+ + HPO42H+ + PO43-
Weak electrolyte, weak acid
Weak electrolyte, weak acid
Weak electrolyte, weak acid
4.3
Neutralization Reaction
Neutralization occurs when a solution of an acid reacts
with a base. The products are salt and water.
acid + base
salt + water
Salt = ionic compound
cation from a base + anion from an acid
HCl (aq) + NaOH (aq)
H+ + Cl- + Na+ + OH-
H+ + OH-
NaCl (aq) + H2O
Na+ + Cl- + H2O
H2O
However, a weak base and an acid gives only a salt
HCl +NH3 NH4Cl
4.3
Oxidation-Reduction Reactions
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
4.4
4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds (e.g. LiH,
CaH2). In these cases, its oxidation number is –1.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the molecule or
ion. For example,NH4+ , the sum of oxidation numbers is
-3+4(+1)=+1
7. Oxidation numbers do not have to be integers.
Oxidation number of oxygen in the superoxide ion,
O2-, is -½.
Oxidation numbers of all
the elements in HCO3- ?
HCO3-
O = -2
H = +1
3x(-2) + 1 + ? = -1
C = +4
4.4
The oxidation numbers of elements in their compounds
*Metallic elements:
only positive
oxidation number;
nonmetallic
elements: positive
or negative
oxidation number;
*The highest
oxidation number of
an element in
group1A-7A is its
group number.
4.4
IF7
Oxidation numbers of all
the elements in the
following ?
F = -1
7x(-1) + ? = 0
I = +7
NaIO3
Na = +1 O = -2
3x(-2) + 1 + ? = 0
I = +5
K2Cr2O7
O = -2
K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
4.4
Oxidation and Reduction always occur together
Zn (s) + CuSO4 (aq)
Zn
Zn2+ + 2e- Zn is oxidized
Cu2+ + 2e-
ZnSO4 (aq) + Cu (s)
Zn is the reducing agent
Cu Cu2+ is reduced CuSO4is the oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq)
Cu
Ag+ + 1e-
Cu(NO3)2 (aq) + 2Ag (s)
Cu2+ + 2eAg Ag+ is reduced AgNO3 is the oxidizing agent
4.4
Types of Oxidation-Reduction Reactions
Combination Reaction
A+B
0
C
+3 -1
0
2Al + 3Br2
2AlBr3
Decomposition Reaction
C
+1 +5 -2
2KClO3
A+B
+1 -1
0
2KCl + 3O2
4.4
Types of Oxidation-Reduction Reactions
Combustion Reaction
A + O2
B
0
0
S + O2
0
0
2Mg + O2
+4 -2
SO2
+2 -2
2MgO
4.4
Types of Oxidation-Reduction Reactions
Single Displacement Reaction
A + BC
0
+1
+2
Sr + 2H2O
+4
0
TiCl4 + 2Mg
0
AC + B
-1
Cl2 + 2KBr
0
Sr(OH)2 + H2 Hydrogen Displacement
0
+2
Ti + 2MgCl2
-1
Metal Displacement
0
2KCl + Br2
Halogen Displacement
4.4
The Activity Series for Metals
The Activity Series
* Arranges metals according
to their ease of oxidation
*The higher the metal on the
Activity Series, the more
active that metal (the easier it
is oxidized.)Any metal can be
oxidized by the metal ions
below it.
*Any metal above hydrogen
will displace it from water or
from an acid.
The Activity Series
• arranges metals according to their ease of oxidation
• can be used to predict reactions
The higher the metal on the Activity Series, Figure 4.16,
the more active that metal (the easier it is oxidized.)
Any metal can be oxidized by the ions of elements below
it on Figure 4.16.
Cu + AgNO3 --> Cu(NO3)2 + Ag
• Cu is above Ag in the activity series; Cu is more active.
Therefore Cu will displace Ag+ from a solution of AgNO3.
• Silver metal will come out of the solution (reduction.)
• The solution will begin to turn blue from the
presence of Cu2+ as the copper metal reacts
(oxidizes.)
3. Halogen Displacement
F2 > Cl2 > Br2 > I2
F2 is the greatest oxidizing halogen
I2 is the least oxidizing halogen
Example:
-1
0
-1
0
2 Br- + Cl2 --> 2 Cl- + Br2
Br2 + Cl- --> no reaction (NR)
Types of Oxidation-Reduction Reactions
Disproportionation Reaction
Element is simultaneously oxidized and reduced.
0
Cl2 + 2OH-
+1
-1
ClO- + Cl- + H2O
Chlorine Chemistry
4.4
Solution Stoichiometry
The concentration of a solution is the amount of solute
present in a given quantity of solvent or solution.
moles of solute
mole
n
=
=
M = molarity =
liters of solution
liters
V
What mass of KI is required to make 500 mL of
a 2.80 M KI solution?
M KI
volume of KI solution
500 mL x
1L
1000 mL
moles KI
x
2.80 mol KI
1 L soln
x
M KI
166 g KI
1 mol KI
grams KI
= 232 g KI
4.5
Dilution is the procedure for preparing a less concentrated
solution from a more concentrated solution.
Dilution
Add Solvent
Moles of solute
before dilution (i)
=
Moles of solute
after dilution (f)
MiVi
=
MfVf
4.5
How would you prepare 60.0 mL of 0.200 M
HNO3 from a stock solution of 4.00 M HNO3?
MiVi = MfVf
Mi = 4.00
Vi =
Mf = 0.200
MfVf
Mi
Vf = 0.06 L
Vi = ? L
0.200 x 0.06
=
= 0.003 L = 3 mL
4.00
3 mL of acid + 57 mL of water = 60 mL of solution
4.5
Titrations
Slowly add base
to unknown acid
UNTIL
the indicator
changes color
What volume of a 1.420 M NaOH solution is
Required to titrate 25.00 mL of a 4.50 M H2SO4
solution?
WRITE THE CHEMICAL EQUATION!
H2SO4 + 2NaOH
M
volume acid
25.00 mL x
acid
2H2O + Na2SO4
rx
moles acid
4.50 mol H2SO4
1000 mL soln
x
coef.
M
moles base
2 mol NaOH
1 mol H2SO4
x
base
volume base
1000 ml soln
1.420 mol NaOH
= 158 mL
4.7