Transcript Solute - Bryant School District

PHall AP info. NOT integrated into this presentation
Chapter 4.1-4.5
Types of Chemical
Reactions and Solution
Stoichiometry
Sugar & Potassium Chlorate video not sure where it goes...
Chapter Four:
TYPES OF CHEMICAL
REACTIONS AND
SOLUTION
STOICHIOMETRY
Chapter 4
Table of Contents
4.1
4.2
4.3
4.4
4.5
Water, the Common Solvent
The Nature of Aqueous Solutions: Strong and Weak
Electrolytes
The Composition of Solutions
Types of Chemical Reactions
Precipitation Reactions
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Section 4.1
Water, the Common Solvent
•
•
•
One of the most
important substances
on Earth.
Can dissolve many
different substances.
A polar molecule
because of its unequal
charge distribution.
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Dissolution of a Solid in a Liquid
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4–
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
“Like Dissolves Like”
Nonpolar solutes dissolve best in nonpolar
solvents
Fats
Benzene
Steroids
Hexane
Waxes
Toluene
Polar and ionic solutes dissolve best in polar
solvents
Inorganic Salts
Water
Sugars
Small alcohols
Acetic acid
Section 4.2
The Nature of Aqueous Solutions: Strong and Weak Electrolytes
Nature of Aqueous Solutions
•
•
•
Solute – substance being dissolved.
Solvent – liquid water.
Electrolyte – substance that when dissolved in
water produces a solution that can conduct
electricity.
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Section 4.2
The Nature of Aqueous Solutions: Strong and Weak Electrolytes
Electrolytes
•
•
•
Strong Electrolytes – conduct current very
efficiently (bulb shines brightly).
Weak Electrolytes – conduct only a small
current (bulb glows dimly).
Nonelectrolytes – no current flows (bulb
remains unlit).
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Electrolytes
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4–
Electrolyte Behavior
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4–
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.
Electrolytes vs. Nonelectrolytes
The ammeter measures the flow of electrons (current)
through the circuit.
If the ammeter measures a current, and the bulb
glows, then the solution conducts.
If the ammeter fails to measure a current, and the
bulb does not glow, the solution is non-conducting.
Try to classify the following substances as
electrolytes or nonelectrolytes…
1.Pure water
2.Tap water
• Sugar solution
• Sodium chloride solution
• Hydrochloric acid solution
• Lactic acid solution
• Ethyl alcohol solution
• Pure sodium chloride
Answers to Electrolytes
ELECTROLYTES:
NONELECTROLYTES:
Tap water (weak)
Pure water
NaCl solution
Sugar solution
• HCl solution
• Lactate solution (weak)
• Ethanol solution
• Pure NaCl
Ionic Compounds “Dissociate”
Na+(aq) + Cl-(aq)
NaCl(s) 
+(aq) + NO -(aq)
Ag
AgNO3(s) 
3
Mg2+(aq) + 2 Cl-(aq)
MgCl2(s) 
2 Na+(aq) + SO42-(aq)
Na2SO4(s) 
AlCl3(s) 
Al3+(aq) + 3 Cl-(aq)
solution where they can
conduct a current rather
than re-forming a
The reason for this is the
solid.
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
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
o
o
o
o
o
Citric acid – citrus fruit
Malic acid – apples
Butyric acid – rancid butter
Amino acids – protein
Nucleic acids – DNA and RNA
Ascorbic acid – Vitamin C
This is an enormous group of compounds;
these are only a few examples.
S
o
l
u
t
i
o
n
s
solutionsxt
Solution Concentration
Calculations of Solution Concentration:
Molarity
Molarity is the ratio of moles of
solute to liters of solution
Section 4.3
The Composition of Solutions
Chemical Reactions of Solutions
•
We must know:
 The nature of the reaction.
 The amounts of chemicals present in
the solutions.
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Section 4.3
The Composition of Solutions
Molarity
•
Molarity (M) = moles of solute per
volume of solution in liters:
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Section 4.3
The Composition of Solutions
Exercise
A 500.0-g sample of potassium phosphate
is dissolved in enough water to make 1.50 L
of solution. What is the molarity of the
solution? (calculate #moles per L to get molarity)
1.57 M
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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
58.44 g
1 L
1 mol
= 43.8 g
Section 4.3
The Composition of Solutions
Concentration of Ions
•
For a 0.25 M CaCl2 solution:
CaCl2 → Ca2+ + 2Cl–
 Ca2+: 1 × 0.25 M = 0.25 M Ca2+
 Cl–: 2 × 0.25 M = 0.50 M Cl–.
Note that there were 2 times as many ions of Cl- formed, so the molarity was
doubled due to having twice as many moles like in the titration lab with citric
acid which had 3 moles hydrogen ions produced for each mole of the base.
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Section 4.3
The Composition of Solutions
Concept Check
Which of the following solutions contains
the greatest number of ions?
a)
b)
c)
d)
400.0 mL of 0.10 M NaCl.
300.0 mL of 0.10 M CaCl2.
200.0 mL of 0.10 M FeCl3.
800.0 mL of 0.10 M sucrose.
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Section 4.3
The Composition of Solutions
Let’s Think About It
•
Where are we going?

•
To find the solution that contains the greatest
number of moles of ions.
How do we get there?


Draw molecular level pictures showing each
solution. Think about relative numbers of ions.
How many moles of each ion are in each
solution?
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Section 4.3
The Composition of Solutions
Notice
•
The solution with the greatest number of
ions is not necessarily the one in which:
 the volume of the solution is the
largest.
 the formula unit has the greatest
number of ions.
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Section 4.3
The Composition of Solutions
Dilution
•
•
•
The process of adding water to a
concentrated or stock solution to achieve
the molarity desired for a particular
solution.
Dilution with water does not alter the
numbers of moles of solute present.
Moles of solute before dilution = moles
of solute after dilution
M1V1 = M2V2
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Section 4.3
The Composition of Solutions
Concept Check
A 0.50 M solution of sodium chloride in an open
beaker sits on a lab bench. Which of the
following would decrease the concentration of
the salt solution?
Add water to the solution.
Pour some of the solution down the sink drain.
Add more sodium chloride to the solution.
Let the solution sit out in the open air for a
couple of days.
e) At least two of the above would decrease the
concentration of the salt solution.
a)
b)
c)
d)
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Serial Dilution
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
Dilution
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4–
Section 4.3
The Composition of Solutions
Exercise
What is the minimum volume of a 2.00 M
NaOH solution needed to make 150.0 mL of
a 0.800 M NaOH solution?
60.0 mL
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Section 4.4
Types of Chemical Reactions
•
•
•
Precipitation Reactions
Acid–Base Reactions
Oxidation–Reduction Reactions
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Precipitation Reactions
Graphic: Wikimedia Commons User Tubifex
Section 4.5
Precipitation Reactions
Precipitation Reaction
•
A double displacement reaction in which
a solid forms and separates from the
solution.
 When ionic compounds dissolve in
water, the resulting solution contains
the separated ions.
 Precipitate – the solid that forms.
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Section 4.5
Precipitation Reactions
The Reaction of K2CrO4(aq) and Ba(NO3)2(aq)
•
Ba2+(aq) + CrO42–(aq) → BaCrO4(s)
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Precipitation of Silver Chloride
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4–
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…
Lead(II) nitrate + potassium iodide lead(II) iodide + potassium nitrate
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
Solubility
Exceptions
NO3-
Soluble
None
ClO4-
Soluble
None
Na+
Soluble
None
K+
Soluble
None
NH4+
Soluble
None
Cl-, I-
Soluble
Pb2+, Ag+, Hg22+
SO42-
Soluble
Ca2+, Ba2+, Sr2+, Pb2+, Ag+, Hg2+
Solubility Rules – Mostly Insoluble
Ion
Solubility
Exceptions
CO32-
Insoluble
Group IA and NH4+
PO43-
Insoluble
Group IA and NH4+
OH-
Insoluble
Group IA and Ca2+, Ba2+, Sr2+
S2-
Insoluble
Groups IA, IIA, and NH4+
Table 4.1 Simple Rules for the
Solubility of Salts in Water
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4–
Solubility Rules
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4–
Solubility
Chart:
Common
salts
at 25C
S = Soluble
I = Insoluble
P = Partially
Soluble
X = Other
Section 4.5
Precipitation Reactions
Precipitates
•
•
•
Soluble – solid dissolves in solution; (aq)
is used in reaction.
Insoluble – solid does not dissolve in
solution; (s) is used in reaction.
Insoluble and slightly soluble are often
used interchangeably.
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Section 4.5
Precipitation Reactions
Simple Rules for Solubility
1. Most nitrate (NO3) salts are soluble.
2. Most alkali metal (group 1A) salts and NH4+ are soluble.
3. Most Cl, Br, and I salts are soluble (except Ag+, Pb2+,
Hg22+).
4. Most sulfate salts are soluble (except BaSO4, PbSO4,
Hg2SO4, CaSO4).
5. Most OH are only slightly soluble (NaOH, KOH are
soluble, Ba(OH)2, Ca(OH)2 are marginally soluble).
6. Most S2, CO32, CrO42, PO43- salts are only slightly
soluble, except for those containing the cations in Rule 2.
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Solubility Trends
 The solubility of MOST solids increases with
temperature.
 The rate at which solids dissolve increases
with increasing surface area of the solid.
 The solubility of gases decreases with
increases in temperature.
 The solubility of gases increases with the
pressure above the solution.
Therefore…
Solids tend to dissolve best when:
o Heated
o Stirred
o Ground into small particles
Gases tend to dissolve best when:
o The solution is cold
o Pressure is high
Solubility Chart
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.
Section 4.5
Precipitation Reactions
Concept Check
Which of the following ions form compounds
with Pb2+ that are generally soluble in water?
a)
b)
c)
d)
e)
S2–
Cl–
NO3–
SO42–
Na+
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Section 4.5
Precipitation Reactions
END OF SLIDES FOR SECTION 4.1-4.5
• END OF SLIDES FOR SECTION 4.1-4.5
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