Solutions and Acid/Base
Solutions
Solutions
• Solute – substance in lower concentration, dissolved
• Solvent – substance in higher concentration
• Water is a universal solvent
Electrolyte vs. Non-electrolyte
• Electrolyte conducts electricity in solution due to the formation of
ions
• Acids, bases, most salts (ionic compounds)
• Non-electrolyte does not conduct electricity
• Sugar, ethanol (many covalent compounds)
Rate of solution formation
• Stirring
• Temperature
• Surface area
• How would each of these affect solution formation?
Solubility
• Amount of solute that will dissolve in a given quantity of a solvent at a
specific temperature and pressure to produce a saturated solution
• Expressed in grams of solute per 100 g of solvent
Solubility
• Unsaturated – able to dissolve more solute if added
• Saturated – maximum amount of solute dissolved for the
temperature and pressure
• Supersaturated – contains more solute than possible due to a slow
temperature drop
Factors affecting solubility
• Temperature – most increase solubility with temperature, but a few
drop
• Solubility of gases in liquid solvent is higher at lower temperatures
• Best seen on solubility graph
• Pressure – does not affect solids or liquid solutes as much as gas
solutes
• Solubility increases with pressure
Solubility Graphs
• Think: How does
solubility change as
temperature
increases?
• Is there any
compound that
does not follow the
trend? Why might
that be?
1. A solution of sodium nitrate at
20 ⁰ C has 100 g dissolved. Is it
saturated, unsaturated, or
supersaturated?
2. How much KNO3 would need to
be added to a solution
containing 25 grams at 37 ⁰C in
order to make the solution
saturated?
1. A solution of sodium nitrate at
20 ⁰ C has 100 g dissolved. Is it
saturated, unsaturated, or
supersaturated? above the line =
super saturated
2. How much KNO3 would need to
be added to a solution
containing 25 grams at 37 ⁰C in
order to make the solution
saturated? 25 grams
Concentration calculations
• Molarity = Moles of solute / liters of solution
• Dilutions = M1 x V1 = M2 x V2
• Percent by volume = (V solute/V solution) x 100%
• Percent by mass = (mass solute / mass solution) x 100%
Molarity
• Intravenous (IV) saline solutions are often administered to patients in
the hospital. One saline solution contains 0.90 g NaCl in exactly 100
mL of solution. What is the molarity of the solution?
• Molarity = moles per liter
• 0.90 g NaCl / 58.44 g/mol = 0.015 mol
• 100 mL = 0.1 L
• 0.015 mol/0.1 L = 0.15 M (0.15 mol/L)
Molarity
• Household laundry bleach is a dilute aqueous solution of sodium
hypochlorite (NaClO). How many moles of solute are present in 1.5 L
of 0.70 M NaClO?
Dilutions
• How many milliliters of aqueous 2.00 M MgSO4 solution must be
diluted with water to prepare 100.0 mL of aqueous 0.400 M MgSO4?
Percent by volume
• What is the percent by volume of ethanol (C2H6O) in the final solution
when 85 mL of ethanol is diluted to a volume of 250 mL with water?
Percent by mass
• If you were making a 2000 g solution of glucose that has a 2.8%
(m/m) concentration, how much glucose should you use?
Colligative property summary
• http://www.chemistryexplained.com/Ce-Co/ColligativeProperties.html
• Freezing point depression (solute lowers point of freezing)
• Boiling point elevation (solute raises point of boiling)
Acids, Bases, Salts
Chapter 19
Acids
• Hydrogen-containing compounds that ionize to yield hydrogen ions in
aqueous solution
• Acids are hydrogen-ion donors
Acids
• Properties
•
•
•
•
•
•
Sour taste
Electrolytes
React with metals
Change color of indicator
(litmus paper red)
(phenolphthalein clear)
• Examples
• Acetic acid (vinegar)
• Citric acid (in candy, citrus fruit,
soda)
• HCl, H2SO4, HNO3
Bases
• Compounds that ionize to yield hydroxide ions in aqueous solution
• Bases are hydrogen-ion acceptors
Bases
• Properties
•
•
•
•
•
•
Bitter taste
Slippery feel
Electrolyte
Change color of indicator
(litmus paper blue)
(phenolphthalein pink)
• Examples
• Soap
• Milk of magnesia
Ionization (dissociation)
• Ionization – becoming an ion
• Dissociation – breaking into smaller particles
• All dissociate, only some ionize
Strong and weak
• Strong – dissociate almost completely in water
• Weak – only dissociate slightly in water
Concentration vs. Strength
• Concentrated/dilute – how much (moles) dissolved in water
• Strong/weak – how ionized/dissociated in solution
• HCl is a strong acid, but can be made into a dilute solution.
• Vinegar is a dilute solution of a weak acid.
Acid Base Neutralization Reaction
• Acids and bases react to neutralize and form salt and water
• HCl + NaOH  NaCl + H2O
• Salt = anion from an acid and cation from a base
Ion product constant for water
• For aqueous solutions, the product of the hydrogen-ion concentration
and the hydroxide-ion concentration equals 1.0 x 10-14
• Kw = [H+] x [OH-] = 1.0 x 10-14
• Acidic [H+] greater than [OH-]
• (greater than 1 x 10-7)
• Candies are slightly acidic. If the [H+] in a solution is 4.5 x 10-6 M, is
the solution acidic, basic, or neutral? What is the [OH-] of this
solution?
pH scale and indicators
• 0-14
• Negative logarithm of the hydrogen-ion concentration
• pH = -log [H+]
pH scale and indicators
• pH = -log [H+]
Neutral solution is [H+] = 1 x 10-7 M
pH = -log (1 x 10-7)
pH = 7.0
• What is the pH of a solution with a hydrogen-ion concentration of
5.67 x 10-11 M?
• The pH of an unknown solution is 11.34. What is its hydrogen-ion
concentration?
pOH
• pOH = -log[OH-]
• pH + pOH = 14
• What is the pH of a solution if
[OH-] = 5.6 x 10-3 M?
Summary Questions
1. Give the pH range for acids and bases.
2. What are the units for molarity?
3. You are making hot chocolate. Describe how you will ensure the
best saturation of the solution.
4. What is the difference between a weak acid and a strong acid?
5. If the pH of the solution is 6.5, what is the [H+] concentration?
What is the pOH?
6. When you open a pop bottle what happens to the pressure? How
does this change the solution? What role would temperature play?