Transcript Slide 1

Ch. 4
Types of Chemical Reactions
and Solution Stoichiometry
Solute
A solute is the dissolved substance in a solution.
Salt in salt water
Sugar in soda drinks
Carbon dioxide in soda drinks
Solvent
A solvent is the dissolving medium in a solution.
Water in salt water
Water in soda
Saturation of Solutions
 A solution that contains the maximum amount
of solute that may be dissolved under existing
conditions is
 saturated.
 A solution that contains less solute than a
saturated solution under existing conditions is
unsaturated.
 A solution that contains more dissolved solute
than a saturated solution under the same
conditions is
supersaturated.
Definition of Electrolytes and
Nonelectrolytes
An electrolyte is:
 A substance whose aqueous solution conducts
an electric current.
A nonelectrolyte is:
 A substance whose aqueous solution does not
conduct an electric current.
Try to classify the following substances as
electrolytes or nonelectrolytes…
Electrolytes?
1.
2.
3.
4.
5.
6.
7.
Pure water
Tap water
Sugar solution
Sodium chloride solution
Hydrochloric acid solution
Ethyl alcohol solution
Pure, solid sodium chloride
Answers…
ELECTROLYTES:
NONELECTROLYTES:
Tap water (weak)
Pure water
NaCl solution
Sugar solution
HCl solution
Ethanol solution
Pure, solid NaCl
But why do some compounds conduct electricity in
solution while others do not…?
Ionic CompoundsDissociate
NaCl(s)  Na+(aq) + Cl-(aq)
AgNO3(s)  Ag+(aq) + NO3-(aq)
MgCl2(s)  Mg2+(aq) + 2 Cl-(aq)
Na2SO4(s)  2 Na+(aq) + SO42-(aq)
AlCl3(s)  Al3+(aq) + 3 Cl-(aq)
Ions tend to stay in solution where they can
conduct a current rather than re-forming a
solid.
The reason for this is the
polar nature of
the water molecule…
Positive ions associate with the negative
end of the water dipole (oxygen).
Negative ions associate with the positive
end of the water dipole (hydrogen).
Some covalent compounds IONIZE in solution
Covalent acids form ions in solution, with the
help of the water molecules.
For instance, hydrogen chloride molecules,
which are polar, give up their hydrogens to
water, forming chloride ions (Cl-) and
hydronium ions (H3O+).
Strong acids such as HCl are completely
100% ionized in solution.
Other examples of strong acids include:
 Sulfuric acid, H2SO4
 Nitric acid, HNO3
 Hydriodic acid, HI
 Perchloric acid, HClO4
Weak acids such as lactic
acid usually ionize less than
5% of the time.
Many of these weaker acids
are “organic” acids
that contain a “carboxyl”
group.
The carboxyl group does not easily give up its
hydrogen.
Because of the carboxyl group, organic acids are
sometimes called “carboxylic acids”.
Other organic acids and their sources include:
o Citric acid – citrus fruit
o Malic acid – apples
o Butyric acid – rancid butter
o Amino acids – protein
o Nucleic acids – DNA and RNA
o Ascorbic acid – Vitamin C
This is an enormous group of compounds; these
are only a few examples.
However, most covalent compounds do not ionize
at all in solution.
Sugar (sucrose – C12H22O11),
and ethanol (ethyl alcohol – C2H5OH) do not
ionize - That is why they are nonelectrolytes!
Molarity
The concentration of a solution measured in moles
of solute per liter of solution.
M = mol
L
Preparation of Molar Solutions
Problem: How many grams of sodium chloride are needed to
prepare 1.50 liters of 0.500 M NaCl solution?
 Step #1: Ask “How Much?” (What volume to prepare?)
 Step #2: Ask “How Strong?” (What molarity?)
 Step #3: Ask “What does it weigh?” (Molar mass is?)
1.500 L
0.500 mol
1L
58.44 g
1 mol
= 43.8 g
Serial Dilution
It is not practical to keep solutions of many different
concentrations on hand, so chemists prepare more dilute
solutions from a more concentrated “stock” solution.
Problem: What volume of stock (11.6 M) hydrochloric
acid is needed to prepare 250. mL of 3.0 M HCl
solution? MstockVstock = MdiluteVdilute
(11.6 M)(x Liters) = (3.0 M)(0.250 Liters)
x Liters = (3.0 M)(0.250 Liters)
11.6 M
= 0.065 L
A. Single Replacement Reactions
A + BX  AX + B
BX + Y  BY + X
Replacement of:
 Metals by another metal
 Hydrogen in an acid by a metal
 Hydrogen in water by a metal
 Halogens by more active halogens
 Ex. Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)
 Ex. 2 Li(s) + 2 H2O(l) 2 LiOH(aq) + H2(g)
The Activity Series of the Metals
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Metals can replace other metals
provided that they are above the
metal that they are trying to
replace.
Metals above hydrogen can
replace hydrogen in acids.
Metals from sodium upward can
replace hydrogen in water
The Activity Series of the Halogens
Fluorine
Chlorine
Bromine
Iodine
Halogens can replace other
halogens in compounds, provided
that they are above the halogen
that they are trying to replace.
2NaCl(s) + F2(g) 
2NaF(s)
??? + Cl2(g)
MgCl2(s) + Br2(g) 
No
Reaction
???
Practice problems - Answers are unbalanced! 
1. Mg + FeCl3  Fe + MgCl2
2. Sodium is added to water.
Na + H2O  H2 + NaOH
3. Lithium is added to hydrochloric acid
Li + HCl  H2 + LiCl
4. Zinc is added to a solution of sodium chloride
Zn + NaCl  N.R.
5. Chlorine gas is bubbled into a solution of potassium iodide
Cl2 + KI  I2 + KCl
6. Chlorine gas is bubbled into a solution of potassium fluoride
Cl2 + KF  N. R.
Double Replacement Reactions
The ions of two compounds exchange places in an
aqueous solution to form two new compounds.
AX + BY  AY + BX
One of the compounds formed is usually a
precipitate (an insoluble solid), an insoluble gas
that bubbles out of solution, or a molecular
compound, usually water.
Double replacement forming a precipitate…
Double replacement (ionic) equation
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
Complete ionic equation shows compounds as aqueous ions
Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) +2 I-(aq)  PbI2(s) + 2K+(aq) + 2 NO3-(aq)
Net ionic equation eliminates the spectator ions
Pb2+(aq) + 2 I-(aq)  PbI2(s)
Solubility Rules – Mostly Soluble
Ion
NO3-
Solubility
Soluble
Exceptions
None
ClO4-
Soluble
None
Alkali
metals
NH4+
Soluble
None
Soluble
None
Cl-,Br-, I- Soluble
Pb2+, Ag+, Hg22+
SO42-
Soluble
Ca2+, Ba2+, Sr2+, Pb2+, Ag+, Hg2+
C2H3O2-
Soluble
Ag+
Solubility Rules – Mostly Insoluble
Ion
Exceptions
CO32-
Solubility
Insoluble
PO43-
Insoluble
Group IA and NH4+
OH-
Insoluble
Group IA and Ca2+, Ba2+, Sr2+
S2-
Insoluble
Groups IA, IIA, and NH4+
CrO4-2
Insoluble
Group IA & NH4+, Ca2+, Sr2+
SO3-2
Insoluble
Group IA and NH4+
Group IA and NH4+
D.R. Practice problems
1. KBr(aq) + AgNO3(aq) 
AgBr(s) + KNO3(aq)
2. Silver nitrate + potassium chromate 
2AgNO3(aq) + K2CrO4(aq)  AgCrO4(s) + 2KNO3(aq)
3. Ammonium chloride + cobalt (II) sulfate 
2NH4Cl(aq) + CoSO4(aq)  (NH4)2SO4(aq) + CoCl2(aq) N.R.
4. Lithium hydroxide + sodium chromate
2LiOH(aq) + Na2CrO4(aq)  2NaOH(aq) + Li2CrO4(s)
5. Zinc acetate + cesium hydroxide 
Zn(C2H3O2)2(aq) + 2CsOH(aq)  Zn(OH)2(s) + 2CsC2H3O2(aq)
6. What is the net ionic equation for the rxn above?
Zn+2(aq) + OH-(aq)  Zn(OH)2(s)
Unstable Compounds!!! (own note paper)
• Ammonium hydroxide
– NH4OH
• Carbonic Acid
– H2CO3
• Sulfurous acid
– H2SO3
• Sulfide salts (ex. Na2S) from acid (H+)
• All break down to form other products!
– NH4OH  NH3(g) + H2O(l)
– H2CO3  CO2(g) + H2O(l)
– H2SO3  SO2(g) + H2O(l)
– S-2  H2S(g)
Unstable Examples:
1. Sodium sulfite + hydrochloric acid
Na2SO3(aq) + 2HCl(aq)  H2SO3(aq) + 2NaCl(aq)
H2SO3(aq)  H2O(l) + SO2(g)
Na2SO3(aq) + 2HCl(aq)  H2O(l) + SO2(g) + 2NaCl(aq)
2. Ammonium sulfate + sodium hydroxide
(NH4)2SO4(aq) + NaOH(aq)  2 NH4OH(aq) + Na2SO4(aq)
2 NH4OH(aq)  2 NH3(g) + H2O(l)
(NH4)2SO4(aq) + 2NaOH(aq)  2NH3(g) + 2H2O(l) + Na2SO4(aq)
What is the net ionic equation for the reaction above?
NH4+(aq) + OH-(aq)  NH3(g) + H2O(l)