Solutions Notes

Words to Know

Solution – homogenous mixture

______– substance present in the largest amount
Solutes – substance present in the smallest amount

______– solutions with water as the solvent

Concentration – the amount of solute in a given volume of solution

______– large amount of solute dissolved in solvent
______– small amount of solute dissolved in solvent
Saturated – a solution that contains as much solute as will dissolve at that temperature
Unsaturated – a solution that hasn’t reached that limit of solute that will dissolve
Supersaturated - a solution that contains as more solute than should dissolve at that temperature

Effect of Temperature on Solubility

Increasing the temperature of a solution, ______the amount of solute that can be dissolved
Decreasing the temperature of a solution, causes the solute to recrystallize

Learning Check

How many grams of NaCl will dissolve in 100 g of H2O at 90°C?

50 g of KCl is dissolved in 100 g of water at 50°C. Is the solution saturated, unsaturated or supersaturated?

Effect of Pressure on Solubility

Pressure has a major effect on the solubility of gas-liquid systems
An increase in pressure ______the solubility of a gas in the liquid

“Like dissolves like” – a solvent usually dissolves solutes that have polarities similar to itself

•______molecules dissolve other polar molecules and ionic compounds.

•______molecules dissolve other nonpolar molecules.

•Alcohols, which have characteristics of both, tend to dissolve in both types of solvents, but will not dissolve ionic solids.

SOLUTES / SOLVENTS
Water / CCl4 / Alcohol
NaCl
I2
C3H7OH
benzene (nonpolar)
Br2
KNO3
toluene (polar)
Ca(OH)2
methanol
NH3
CO2

•Colligative properties - the physical changes that result from adding solute to a solvent. Colligative Properties depend on how many solute particles are present as well as the solvent amount, but they do NOT depend on the type of solute particles.

•Boiling Point ______

•Freezing Point ______

•Osmotic Pressure

•Vapor Pressure ______

Learning Check

1. Which substance will provide the greatest change in freezing point?

A. NaClB. CaCl2C. C6H12O6 D. H2O

2. Which of the following reflect colligative properties?

I. 0.5 m NaBr solution has a higher vapor pressure than a 0.5 m BaCl2 solution.

II. A 0.5 m NaOH solution freezes at a lower temperature than pure water.

III. Pure water freezes at a higher temperature than pure methanol.

A. only I B. only II C. only III D. I and II E. I and III

3. A student measured the conductivity, of unlabeled liquids, after each added drop. The following graph was produced...

a. Identify the line that represents:

–aluminum chloride

–water

–magnesium chloride

–sugar

–sodium chloride

Which line could also represent potassium iodide?

Solution Composition - Mass Percent

Mass percent – describes a solution’s composition

- expresses the mass of solute present in a given mass of solution

Example – A solution is prepared by mixing 1.00g of C2H5OH, with 100.0g of

H2O. Calculate the mass percent of ethanol.

Solution Composition – Molarity

Molarity – measure of concentration

- number of moles of solute per volume of solution in liters

Example – Calculate the molarity of a solution prepared by dissolving 11.5 gNaOH

in enough water to make 1.50L solution.

Example - Calculate the mass of solid AgCl formed when 1.50L of a 0.100M

AgNO3 solution is reacted with excess NaCl.

Example – How many moles of Ag+ ions are present in 25mL of a 0.75M Ag2SO4

solution?

Learning check

Calculate the molarity of a solution prepared by dissolving 25.6 g NaC2H3O2 in enough water to make 200.0 mL solution.

Standard Solution – a solution whose concentration is accurately known

Example – A chemist needs 1.0L of a .200M K2Cr2O7 solution. How much solid

K2Cr2O7 must be weighed out to make this solution?

Dilution – process of adding more solvent to a solution

Moles of solute before dilution = Moles of solute after dilution

Example: What volume of 16M H2SO4 must be used to prepare 1.5L of a 0.10M

H2SO4 solution?

Learning Check: Prepare 500.0mL of 1.00 M HC2H3O2 from a 17.5 M stock solution.

What volume of the stock solution is required?

Notes - Acids and Bases

Acids and Bases

Arrhenius ACIDS – produces hydrogen ions in aqueous solutions, sour taste, low pH, and the fact that they turn litmus paper red

HCl (aq)  H+(aq) + Cl-(aq)

Arrhenius BASES – produces hydroxide ions in aqueous solutions, bitter taste, slippery feel, high pH, and the fact that they turn litmus paper blue

NaOH (aq)  Na+(aq) + OH-(aq)

Arrhenius definition – limits the concept of a base

Brønsted – Lowry definition – gives a broader definition of a base

Brønsted – Lowry ACID – a proton (H+) donor

Brønsted – Lowry BASE – a proton (H+) acceptor

General Reaction –

HA (aq) + H2O (l)  H3O+(aq) + A-(aq)

Conjugate Base – everything that remains of the acid molecule after a proton is lost

Conjugate Acid – the base with the transferred proton (H+)

Conjugate Acid – Base Pair – two substances related to each other by the donating and accepting of a single proton

Examples: Finish each equation and identify each member of the conjugate acid –base

pair.

H2SO4 (aq) + H2O (l) 

CO32- (aq) + H2O (l) 

Learning Check

Write the conjugate ACID

NH3
HCO3-1

Write the conjugate BASE

H3PO4
HBr

Finish each equation and identify each member of the conjugate acid –base pair.

H2SO3 (aq) + H2O (l) 
SO4-2 (aq) + H2O (l) 

The hydronium ion, H3O+, forms when water behaves as a base. This happens when the two unshared pairs of electrons on O bond covalently with the H+.

Water as an Acid and a Base

Amphoteric – a substance that can behave as either an acid or a base

- water is the most common amphoteric substance

Ionization of Water - H2O (l) + H2O (l)  H3O+ (aq) + OH- (aq)

In the shorthand form: H2O (l)  H+ (aq) + OH- (aq)

Ion-product constant – Kw refers to the ionization of water

Kw = [H+][OH-]

At 25C, Kw = [H+][OH-] = [1.0 x 10-7] [1.0 x 10-7] = 1.0 x 10-14

If [H+] increases, the [OH-] decreases, so the product of the two is still 1.0 x 10-14.

There are three possible situations –

A neutral solution, where [H+] = [OH-]
An acidic solution, where [H+]  [OH-]
A basic solution, where [H+]  [OH-]

Example: Calculate [H+] or [OH-] as required for each of the following solutions at 25C,

for each solution state whether it is neutral, acidic, or basic.

a. 1.0 x 10-5 M OH-b. 10.0 M H+

pH scale

pH scale – because the [H+] in an aqueous solution is typically small, logarithms are

used to express solution acidity

pH = -log [H+]pOH = -log [OH-]

Graphing calculatorNon graphing calculator

1. Press the +/- key1. Enter the [H+]

2. Press the log key2. Press the log key

3. Enter the [H+]3. Press the +/- key

Significant Figure Rule – The number of places to the right of the decimal for a log must

be equal to the number of significant figures in the original

number.

Example – Calculate the pH or pOH

a. [H+] = 5.9 x 10-9 Mb. [OH-] = 2.4 x 10-6 M

Since Kw = [H+][OH-] = 1.0 x 10-14, pH + pOH = 14.00

Example - The pH of blood is about 7.4. What is the pOH of blood?

In order to calculate the concentration from the pH or pOH,

[H+] = 10-pH [OH-] = 10-pOH

Graphing calculatorNon graphing calculator

1. Press the 2nd function, then log1. Enter the pH

2. Press the +/- key2. Press the +/- key

3. Enter the pH3. Press the inverse log key

Example - The pH of a human blood sample was measured to be 7.41. What is the

[H+] in blood?

Example – The pOH of the water in a fish tank is found to be 6.59. What is the [OH-]

for this water?

Learning Check

Determine the pH of a solution with a hydrogen ion concentration of 3.2 x10-12 M.

What is the [OH-] concentration of a solution with a hydrogen ion concentration of 8.9x10-4M?

What is the pH of a solution with a hydroxide ion concentration of 5.7x10-10 M?

How Do We Measure pH?

For less accurate measurements, one can use

Litmus paper

•“Red” paper turns blue above ~pH = 8

•“Blue” paper turns red below ~pH = 5

An indicator

•For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

Strong Acids

•The seven strong acids are ______

•These are, by definition, strongelectrolytes and exist totally as ions in aqueous solution.

Strong Bases

•Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).

•Again, these substances dissociate completely in aqueous solution, ______

•Electrolytesare substances which, when dissolved in water, break up into cations (plus-charged ions) and anions (minus-charged ions). We say they ionize. Strong electrolytes ionize completely (100%), while weak electrolytes ionize only partially (usually on the order of 1–10%). The ions in an electrolyte can be used to complete an electric circuit and power a bulb.

•Strong electrolytes fall into three categories: strong acids, strong bases, and soluble salts.

•The weak electrolytes include weak acids, weak bases and insoluble salts.

•Molecules are nonelectrolytes.

Substance / Classification - Strong acid, weak acid, strong base, weak base, soluble salt, insoluble salt, molecule / Strong electrolyte, weak electrolyte, nonelectrolyte
sodium hydroxide
acetic acid
potassium nitrate
hydrobromic acid
silver chloride
Carbon dioxide
chloric acid
barium carbonate
nitric acid
sulfurous acid
strontium sulfate
ethanol
octane (gasoline)

Titration

A known concentration of base (or acid) is slowly added to a solution of acid (or base).

A pH meter or indicators are used to determine when the solution has reached the ______, at which the stoichiometric amount of acid equals that of base.

From the start of the titration to near the equivalence point, the pH goes up slowly.

Just before and after the ______, the pH increases rapidly.

At the equivalence point, moles acid = moles base, and the solution contains only water and the salt from the cation of the base and the anion of the acid.

As more base is added, the increase in pH again levels off.

Neutralization Reaction

Acid + Base  Salt + Water

Salt – ionic compound containing a positive ion other than H+ and a negative ion other

than OH-

Buffered solutions – resists a change in its pH even when a strong acid or base is added to it

- A solution is buffered in the presence of a weak acid and its conjugate base

Summary