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Mixtures, Water, & Solutions
Chapter 2,15, & 16
Water
Properties
• Water is a polar molecule.
• Hydrogen bonds form between water molecules.
– High surface tension
– Low vapor pressure
– Maximum density at 4oC (as a liquid!)
• What would happen to aquatic life if ice were
denser than water?
Solutions (Homogeneous Mixtures)
• Solute: becomes dispersed in the solvent
• Solvent: dissolves the solute
• Aqueous Solutions:
substance dissolved in water.
(water = solvent)
Examples
Kool – Aid
Solute – powered substance
Solvent – water
Chocolate Milk
Solute – chocolate powder or syrup
Solvent – milk
Identify the solvent and the solute
• A teaspoon of sugar is dissolved in 200.2g of water.
• Sterling silver is made by adding small amounts to
copper to pure silver.
• Jell-O consists of solid particles that were dissolved and
then left suspended in water.
• NaCl(aq)
“Like Dissolves Like”
– Dissolving is a physical process in which particles of a
solute are held apart by particles of the solvent.
• RULE: Like Dissolves Like
– Nonpolar solvents (ex. oil) dissolve nonpolar
compounds.
– Polar solvents (ex. water) dissolve polar compounds
and most ionic compounds (ex. NaCl).
“Like Dissolves Like” Examples
• Oil (nonpolar) does not
dissolve in water (polar).
• Alcohols (ex. CH3CH2OH –
ethanol) have both a polar
and a nonpolar end.
– Dissolve polar and nonpolar
solutes, but NOT ionic
compounds.
What happens when compounds
dissolve?
Dissolving Ionic Compounds
• NaCl (s)  Na+ (aq) + Cl- (aq)
• Dissociation in Water causes Ions to formed.
What happens when compounds
dissolve?
Dissolving Covalent Compounds
• C12H22O11 (s)  C12H22O11 (aq)
• NO dissociation because NO ions
• Sucrose dissolves in water because sugar is polar
(-OH group), but dissociation does not occur.
Sucrose molecules are simply separated from
each other. No ions are formed
Electrolytes
• Electrolyte: a compound that conducts an
electric current when it is in an aqueous state.
– Mobile ions are required for the conduction of electric
current.
– Ex. Ionic Compounds (salts), acids, and bases.
• Nonelectrolytes: cannot conduct electricity
– Ex. glucose, alcohol
Solubility
• Solubility: The amount of solute that dissolves in
a given quantity of a solvent at a specified
temperature and pressure.
Concentration of Solute in a Solution
• Concentrated: Relatively more solute in the
solution.
• Dilute: Relatively less solute in the solution.
– Example: A 0.02 M solution is more dilute than a
2 M solution.
Saturation
 Unsaturated Solution: the amount of solute dissolved
is less than the maximum that could be dissolved.
 Ex. Earth’s oceans = unsaturated salt solution
 The ocean can hold a LOT more salt than it actually has in it!
 Saturated Solution: the solution holds the maximum
amount of solute.
 Supersaturated Solution: contains more solute than
the usual maximum amount and is unstable (may
release solute suddenly).
 Created by dissolving solute in the solution at a high
temperature then slowly cooling the solution.
Solubility
Curves
1. How many grams
of potassium
chloride would
dissolve in 100 g
of water at 90oC?
55 grams
2. How much would
dissolve in
200 g of water?
55 grams x 2
= 110 grams
Solubility
Curves
1. A solution has
132 g of NaNO3
in 100 g of water
at 75oC.
Is the solution
UNSATURATED
–
saturated,
falls
BELOW the line
unsaturated,
or
of saturation
supersaturated?
2. What about 95 g SUPERSATURATED –
of KNO3 in 100g falls ABOVE the line
water at 50oC?
of saturation
Factors Affecting
Solubility
1) Temperature
A. Solids dissolving in liquids
 solubility increases with
increasing temperature
B. Gases dissolving in liquids
 solubility decreases
with
increasing temperature
–
–
Carbonated soda
Thermal pollution from
industry
affecting dissolved oxygen
in lakes
Factors Affecting Solubility
2) Pressure
– Solids and Liquids  little effect
– Gases  solubility increases under increased
pressure.
- ex. carbonated beverages
3) Stirring / Solute Particle Size /
Solute Surface Area (crushing)
Calculating the Concentration of Solutions
• Molarity (M) = moles of solute
liters of solution
• 3M NaCl = “three molar solution of sodium
chloride”
Molarity Practice
1. What is the molarity of a solution in which 67
g of NaCl are dissolved in 1 L of solution?
2. How many grams of KNO3 should be used to
prepare 2 L of a 1 M solution?
Making Dilutions
• When you dilute a solution, you increase the
amount of solvent (the number of solute particles
stays the same).
• M 1 x V1 = M 2 x V2
• M = molarity
V = volume (in mL or L  must be same for both)
Practice
1. A chemist starts with 50.0 mL of a 0.40 M NaCl
solution and dilutes it to 1000. mL. What is the
concentration of NaCl in the new solution?
2. If you dilute 175 mL of a 1.6 M solution of LiCl to 1.0 L,
determine the new concentration of the solution.
1. A chemist wants to make 500. mL of 0.050 M HCl by
diluting a 6.0 M HCl solution. How much of that
solution should be used?
Solution Compared to Pure Solvent
 Freezing-Point Depression – presence of solute
particles disrupts formation of orderly pattern
found in solid phase.
 Ex. Antifreeze
 Vapor-Pressure Lowering (nonvolatile solute)
 Boiling-Point Elevation – lowers vapor pressure
 For aqueous solutions, BP will be higher than 100oC
Acids & Bases
Acids
• Sour taste
• Examples:
– HCl (stomach acid), H2SO4
– HC2H3O2 (acetic acid) – (vinegar = acetic acid & water)
– Citric acid (lemon juice)
• Releases H+ when dissolved in water,
producing hydronium ions!
H+ + H2O  H3O+
• Electrolyte
Acids (cont’d)
• React with metals to produce H2 gas.
– Ex. HCl + Zn  ZnCl2 + H2
• When diluting acids, always slowly
pour the acid into water while stirring.
• Acid/Base Indicators:
– Turns litmus paper red.
– Phenolphthalein does not change color.
Bases
• Bitter taste, slippery feel
• Examples: NaOH, Mg(OH)2 (milk of magnesia),
NH3 (ammonia), soap, household cleaners
• Releases OH- (hydroxide ions) when dissolved in
water.
• Electrolyte.
• Acid/Base Indicators:
– Turns litmus paper blue.
– Phenolphthalein turns bright pink.
Naming Common Acids
Textbook Table 9.5
Anion
Ending
Example
Acid Name
Example
-ide
Chloride, Cl-
Hydro-(stem)-ic acid
Hydrochloric acid
-ite
Sulfite, SO32-
(stem)-ous acid
Sulfurous acid
-ate
Phospate, PO43-
(stem)-ic acid
Phosphoric acid
3 Definitions of Acids & Bases
1) Arrhenius Theory
• Acids: ionize to produce H+ ions in aqueous solution
– Monoprotic: HNO3, HCl, HC2H3O2
– Diprotic: H2SO4, H2SO3
– Triprotic: H3PO4
• Bases: dissociate to produce OH- ions in aqueous
solution
– NaOH, KOH, Ca(OH)2, Al(OH)3
Definitions (cont’d)
2) Brønsted-Lowry Theory
– Hydrogen ion (H+) = a proton
• Acids: proton donors – ex. HCl
• Bases: proton acceptors – ex. NH3
Conjugate Acids and Bases
• Show the direction of H+ transfer.
• Label: Acid, Base, Conjugate Base, Conjugate Acid
• NH3 (aq) + H2O (l)
NH4+ (aq) + OH- (aq)
• HCl (g) + H2O (l)
H3O+ (aq) + Cl- (aq)
More Examples + WS
• Show the direction of H+ transfer.
• Label: Acid, Base, Conjugate Base, Conjugate Acid
• H2SO4 + OH-
HSO41- + H2O
• HSO41- + H2O
SO42- + H3O+
Definitions (cont’d)
3) Lewis Theory
• Acids: electron-pair acceptor
• Bases: electron-pair donor
• HCl (g) + H2O (l)
H3O+ (aq) + Cl- (aq)
Strong Acids and Bases
• Strong Acids: completely ionize in water
– Ex. HCl, HBr, HI, HNO3, H2SO4, HClO3
• Strong Bases: completely dissociate into ions
in water
– Ex. NaOH, LiOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak Acids and Bases
• Weak Acids: only some molecules ionize in water
– Ex: acetic acid (less than 0.5% of molecules ionize)
• Weak Bases: do not completely dissociate into
ions in water
– Ex: ammonia (only 0.5% of molecules dissociate)
Concentrated vs. Strong
• “Concentrated” – refers to the amount dissolved
in solution.
• “Strong” – refers to the fraction of molecules that
ionize.
• For example, if you put a lot of ammonia into a
little water, you will create a highly concentrated
solution. However, since only 0.5% of ammonia
molecules ionize in water, this basic solution will
not be very strong.
pH Scale
pH Scale
• pH Scale: logarithmic scale in which [H3O+] is
expressed as a number from 0 to 14.
– [OH-] = [H3O+] when pH = 7.
pH Scale
Acidic
Neutral
Basic
0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
100 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14
[H3O+]
10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1 100
[OH-]
pH equations
•
•
•
•
pH= - log [H3O+]
pOH = - log [OH-]
pH + pOH = 14
[H+] x [OH-] = 1 x 10-14
pH – Examples
1) What is the pH if [HCl] = 1 x 10-4 M?
2) What is the [H+] if the pH = 9?
3) What is the pH if [NaOH] = 1 x 10-2 M
4) What is the concentration of [OH-] if the pOH is 3?
5) What is the concentration of [H+] if the pOH is 10?
pH calculations – easy practice
1) What is the pH of the solution?
a) [H3O+] = 1 x 10-4 M
b) [H+] = 1 x 10-10 M
c) [HCl] = 1 x 10-2 M
2) What is the concentration of H3O+ if the pH is 5?
3) What is the concentration of H+ if the pH is 11?
(cont’d)
4) What is the pOH of each solution?
a) [OH-] = 1 x 10-4 M
b) [NaOH] = 1 x 10-10 M
5) What is the pH of a solution if the pOH is 4?
6) What is the pH of each solution?
a) [OH-] = 1 x 10-8 M
b) [KOH] = 1 x 10-3 M
7) What is the [H3O+] in a solution if [OH-] = 1 x 10-3 M?
8) What is the [OH-] in a solution if [H3O+] = 1x 10-5 M?
pH equations
•
•
•
•
pH= - log [H3O+]
pOH = - log [OH-]
pH + pOH = 14
[H+] x [OH-] = 1 x 10-14
Practice – Using the equations
Find the pH of the following solutions.
Is the solution acidic or basic?
1) 0.01 M HCl
2) 0.050 M Ca(OH)2
3) 2.6 x 10-12 M Mg(OH)2
4) 1 x 10-7 M HC2H3O2
5) Find the concentration of hydrogen ions if the pH is
3.
6) Find the concentration of hydroxide ions if the pH
is 5.6.
7) Find the [H3O+] in a solution if [OH-] = 3 x 10-6 M
Neutralization Reaction
• ACID + BASE  SALT + WATER
– Salt: ionic compound formed from the negative
part of the acid and the positive part of the base.
• Example:
2HCl + Mg(OH)2 MgCl2 + 2H2O
• What type of reaction is this? Synthesis,
Decomposition, Single Replacement, Double
Replacement, or Combustion?
Complete and Balance the
Neutralization Reactions
1) HCl + NaOH 
2) HC2H3O2 + Ca(OH)2 
3) HBr + Al(OH)3 
Acid-Base Titration
• Uses a neutralization reaction to
determine the concentration of
an acid or base.
• Standard Solution: the reactant that has a
known molarity
• Endpoint: the point at which the unknown has
been neutralized.
Titration Examples
• Example #1) 8.0 mL of 0.100M NaOH is used
to neutralize 20.0 mL of HCl. What is the
molarity of HCl?
Titration Examples (cont’d)
• Example #2) A 0.1M Mg(OH)2 solution was
used to titrate an HBr solution of unknown
concentration. At the endpoint, 21.0 mL of
Mg(OH)2 solution had neutralized 10.0 mL of
HBr. What is the molarity of the HBr solution?
Titration Practice
Example #3) What is the molarity of an Al(OH)3 solution if
30.0 mL of the solution is neutralized by 26.4 mL of a 0.25
M HBr solution?
Equilibrium
Factors Affecting
Chemical Reaction Rates
1) Temperature
 Higher temperature = faster reaction (ex. baking a
cake)
 Lower temperature = slower reaction (ex. batteries
in refrigerator)
2) Concentration of Reactants
– Increased concentration = higher reaction rate
(higher frequency of collisions between
reactants).
Factors Affecting
Chemical Reaction Rates
3) Smaller Particle Size/Greater Surface Area
– Increases the amount of reactant exposed for
reacting.
4) Catalysts/Inhibitors
– Added to a reaction to speed it up (or slow it down).
– Is not permanently changed or used up during the
reaction.
– A catalyst speeds up reaction by lowering activation
energy.
– An inhibitor interferes with the catalyst.
– Enzymes = biological catalysts
LeChâtelier’s Principle
Regaining Equilibrium
3 Stresses that Upset Equilibrium
1. Concentration: Changing the amount, or
concentration, of reactants or products in a system at
equilibrium disturbs the equilibrium.
Ex. Removing products to increase yield
Chicken Egg Production: Chicken egg
If the farmer removes the product (egg), the chicken produces more.
Ex. H2CO3 (aq) CO2 (aq) + H2O (l)
Add H2CO3 SHIFT _____________
Add CO2 SHIFT ___________
Remove CO2 SHIFT ___________
3 Stresses that Upset Equilibrium
2. Temperature: Increasing the temperature causes the
equilibrium to shift in the direction that absorbs
heat. Think of heat as a product or reactant and treat
like concentration changes.
Ex. 2SO2 (g) + O2 (g) 2SO3 (g) + heat
Increase temperature SHIFT ____________
Decrease temperature SHIFT __________
3 Stresses that Upset Equilibrium
3. Pressure: A change in the pressure on a system
affects only an equilibrium that has an unequal
number of moles of gaseous reactants and products.
Ex. Formation of Ammonia
N2 (g) + 3H2 (g) 2NH3 (g)
Increase pressure, shifts to side with fewer mol of gas
SHIFT _________
Decrease pressure, shifts to side with more mol of gas
SHIFT __________
Equilibrium
• The rate of the forward reaction is equal to
the rate of the reverse reaction.
• The concentration of products and reactants
stays the same, but the reactions are still
running.
• Shown with the double arrow.
Measuring equilibrium
• At equilibrium the concentrations of products
and reactants are constant.
• We can write a constant that will tell us where
the equilibrium position is.
• Keq = [Products]coefficients
[
[Reactants]coefficients
– Keq = equilibrium constant
– Square brackets [ ] means concentration in molarity
(moles/liter)
Writing Equilibrium Expressions
• General equation
nA + mB
xC + yD
Keq = [C]x [D]y
[A]n [B]m
Practice
• Write the equilibrium expressions for the
following reactions.
• 3H2(g) + N2(g)
2NH3(g)
• 2H2O(g)
2H2(g) + O2(g)
Calculating Equilibrium
• Keq is the equilibrium constant, it is only effected by
temperature.
Calculate the equilibrium constant for the
following reaction.
3H2(g) + N2(g) 2NH3(g)
If at 25ºC there 0.15 mol of N2 , 0.25 mol of
NH3, and 0.10 mol of H2 in a 2.0 L container.
What it tells us
• If Keq > 1 Products are favored
– More products than reactants at equilibrium
• If Keq < 1 Reactants are favored
Practice
1. At 1500K, an equilibrium mixture of N2, O2, and NO gases
consists of 6.4 mol/L of N2, 1.7 mol/L of O2, and 0.011 mol/L of
NO. What is the equilibrium constant for the system? What is
favored at equilibrium?
N2 (g) + O2 (g) 2NO (g)
2. At equilibrium at 2500K, [HCl] = 0.0625 mol/L and [H2] = [Cl2] =
0.0045 mol/L for the reaction. Find the value of Keq. What is
favored at equilibrium?
H2 (g) + Cl2 (g)
2HCl (g)