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9/22/2010
A solution is a homogenous mixture of 2 or more
substances
The solute is(are) the substance(s) present in the
smaller amount(s)
The solvent is the substance present in the larger
amount
Reactions in Aqueous Solution
Chapter 4
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g)
N2
O2, Ar, CH4
Soft Solder (s)
Pb
Sn
4.1
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Conduct electricity in solution?
An electrolyte is a substance that, when dissolved in
water, results in a solution that can conduct electricity.
Cations (+) and Anions (-)
A nonelectrolyte is a substance that, when dissolved,
results in a solution that does not conduct electricity.
Strong Electrolyte – 100% dissociation
NaCl (s)
H2O
Na+ (aq) + Cl- (aq)
Weak Electrolyte – not completely dissociated
CH3COO- (aq) + H+ (aq)
CH3COOH
nonelectrolyte
weak electrolyte
strong electrolyte
4.1
4.1
Nonelectrolyte does not conduct electricity?
Hydration is the process in which an ion is surrounded
by water molecules arranged in a specific manner.
No cations (+) and anions (-) in solution
C6H12O6 (s)
H2O
C6H12O6 (aq)
Strong Electrolyte Weak Electrolyte
Nonelectrolyte
HCl
CH3COOH
(NH2)2CO
HNO3
HF
CH3OH
HClO4
HNO2
C2H5OH
NaOH
H2O
C12H22O11
Ionic Compounds
H2O
4.1
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Writing Net Ionic Equations
Precipitation Reactions
Precipitate – insoluble solid that separates from solution
precipitate
1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes
3. Determine precipitate from solubility rules
Pb(NO3)2 (aq) + 2NaI (aq)
PbI2 (s) + 2NaNO3 (aq)
4. Cancel the spectator ions on both sides of the ionic equation
molecular equation
Pb2+ + 2NO3- + 2Na+ + 2I-
PbI2 (s) + 2Na+ + 2NO3-
ionic equation
Pb2+
PbI2
+ 2I-
PbI2 (s)
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-
Na+ and NO3- are spectator ions
AgCl (s)
4.2
4.2
Acids
Solubility Rules for Common Ionic Compounds
In water at 250C
Soluble Compounds
Exceptions
Have a sour taste. Vinegar owes its taste to acetic acid. Citrus
fruits contain citric acid.
Compounds containing alkali
metal ions and NH4+
React with certain metals to produce hydrogen gas.
NO3-, HCO3-, ClO3Cl-,
Br-, I-
Halides of
Ag+,
Hg22+,
Pb2+
SO42-
Sulfates of Ag+, Ca2+, Sr2+, Ba2+,
Hg2+, Pb2+
Insoluble Compounds
Exceptions
CO32-, PO43-, CrO42-, S2-
Compounds containing alkali
metal ions and NH4+
OH-
Compounds containing alkali
metal ions and Ba2+
React with carbonates and bicarbonates to produce carbon
dioxide gas
Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
4.2
4.3
Arrhenius acid is a substance that produces H + (H3O+) in water
A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
Arrhenius base is a substance that produces OH - in water
base
acid
acid
base
A Brønsted acid must contain at least one
ionizable proton!
4.3
4.3
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Neutralization Reaction
Monoprotic acids
H+ + Cl-
HCl
Strong electrolyte, strong acid
H+ + NO3-
HNO3
Strong electrolyte, strong acid
H+ + CH3COO-
CH3COOH
acid + base
Weak electrolyte, weak acid
salt + water
Diprotic acids
H2SO4
HSO4
-
H+ + HSO4-
Strong electrolyte, strong acid
H+ + SO42-
Weak electrolyte, weak acid
H+
Triprotic acids
H3PO4
H2PO4HPO42-
HCl (aq) + NaOH (aq)
+
Cl-
+
Na+
+
OH-
H+ + OH-
H+ + H2PO4H+ + HPO42H+ + PO43-
Weak electrolyte, weak acid
NaCl (aq) + H2O
Na+ + Cl- + H2O
H2O
Weak electrolyte, weak acid
Weak electrolyte, weak acid
4.3
4.3
Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg (s) + O2 (g)
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2Mg
O2 + 4e-
2O2-
Reduction half-reaction (gain e-)
2Mg + O2 + 4e-
2Mg2+ + 2O2- + 4e-
2Mg + O2
2MgO
Zn (s) + CuSO4 (aq)
Cu2+ + 2e-
4.4
Oxidation number
ZnSO4 (aq) + Cu (s)
Zn2+ + 2e- Zn is oxidized
Zn
4.4
Zn is the reducing agent
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
Cu Cu2+ is reduced Cu2+ is the oxidizing agent
1. Free elements (uncombined state) have an oxidation
number of zero.
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq)
Cu2+ + 2e-
Cu
Ag+
Cu(NO3)2 (aq) + 2Ag (s)
+
1e-
Ag
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
Ag+
is reduced
Ag+
is the oxidizing agent
4.4
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
4.4
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4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
IF7
Oxidation numbers of all
the elements in the
following ?
F = -1
7x(-1) + ? = 0
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
I = +7
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.
Oxidation numbers of all
the elements in HCO3- ?
O = -2
K2Cr2O7
NaIO3
HCO3-
O = -2
Na = +1 O = -2
H = +1
3x(-2) + 1 + ? = 0
3x(-2) + 1 + ? = -1
7x(-2) + 2x(+1) + 2x(?) = 0
I = +5
C = +4
K = +1
Cr = +6
4.4
Types of Oxidation-Reduction Reactions
Combination Reaction
C
0
+4 -2
S + O2
A + BC
SO2
0
+4
+1 +5 -2
+1 -1
0
Sr(OH)2 + H2
0
0
-1
Hydrogen Displacement
+2
Ti + 2MgCl2
-1
Cl2 + 2KBr
0
AC + B
+2
TiCl4 + 2Mg
A+B
0
2KClO3
+1
Sr + 2H2O
Decomposition Reaction
C
Types of Oxidation-Reduction Reactions
Displacement Reaction
A+ B
0
4.4
Metal Displacement
0
2KCl + Br2
Halogen Displacement
2KCl + 3O2
4.4
The Activity Series for Metals
4.4
Types of Oxidation-Reduction Reactions
Disproportionation Reaction
Displacement Reaction
M + BC
Element is simultaneously oxidized and reduced.
AC + B
0
Cl2 + 2OH-
M is metal
BC is acid or H2O
B is H2
+1
-1
ClO- + Cl- + H2O
Chlorine Chemistry
Ca + 2H2O
Ca(OH)2 + H2
Pb + 2H2O
Pb(OH)2 + H2
4.4
4.4
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Solution Stoichiometry
Classify the following reactions.
The concentration of a solution is the amount of solute
present in a given quantity of solvent or solution.
Ca2+ + CO32-
CaCO3
NH3 + H+
NH4+
Zn + 2HCl
M = molarity =
Acid-Base
ZnCl2 + H2
Ca + F2
Precipitation
CaF2
Redox (H2 Displacement)
Redox (Combination)
moles of solute
liters of solution
What mass of KI is required to make 500. mL of
a 2.80 M KI solution?
M KI
volume KI
500. mL x
moles KI
1L
1000 mL
x
2.80 mol KI
1 L soln
M KI
x
grams KI
166 g KI
1 mol KI
= 232 g KI
4.4
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
Gravimetric Analysis
How would you prepare 60.0 mL of 0.2 M
HNO3 from a stock solution of 4.00 M HNO3?
1. Dissolve unknown substance in water
2. React unknown with known substance to form a precipitate
MiVi = MfVf
Mi = 4.00
Vi =
Mf = 0.200
MfVf
Mi
=
Vf = 0.06 L
4.5
3. Filter and dry precipitate
4. Weigh precipitate
Vi = ? L
5. Use chemical formula and mass of precipitate to determine
amount of unknown ion
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
4.6
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Titrations
What volume of a 1.420 M NaOH solution is
Required to titrate 25.00 mL of a 4.50 M H2SO4
solution?
In a titration a solution of accurately known concentration is
added gradually added to another solution of unknown
concentration until the chemical reaction between the two
solutions is complete.
Equivalence point – the point at which the reaction is complete
WRITE THE CHEMICAL EQUATION!
Indicator – substance that changes color at (or near) the
equivalence point
H2SO4 + 2NaOH
M
volume acid
Slowly add base
to unknown acid
UNTIL
25.00 mL x
the indicator
changes color
4.7
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
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