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Chapter 13: Solutions Check MasteringChemistry and other deadlines *diagrams from the class textbook 2018 Pearson

Solutions:

Solutions are homogeneous mixtures of two or more substances.

Solutions come in all phases: gas (air), liquid (seawater, vodka,

carbonated water), and solid (brass, an alloy of Zn/Cu)

Solute: the part in the solution that is less

Solvent: the part in the solution that is more

Aqueous solutions have water as the solvent

Polarity:

Substances may be considered polar, nonpolar or ionic

Polar: Polar molecules have polar bonds that do not cancel each other

out. Determining polarity requires knowledge of chemical bonding

and molecular shapes that will be studied later.

For now, be aware of common examples: water and alcohols

(CxHyOH)

Nonpolar: Nonpolar molecules have nonpolar bonding or possibly polar

bonds that cancel each other out.

For now, be aware of examples: hydrocarbons (CxHy), like alkanes

or oils, diatomic elements

Ionic: examples include salts, acids and bases that are made of ion.

When dissolved ionic compounds break up into charged cations and

anions.

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Like dissolves like:

Polar substances will dissolve in polar solvents

Nonpolar substances dissolve in nonpolar solvents

Soluble ionic compounds will dissolve in a strong polar substance like

water

Miscible: two or more substances blend to form a solution (water and

alcohol). Like polarities will blend to make a solution.

Immiscible: two or more substances create layers when added together

(oil and vinegar). Nonpolar and polar substances generally do not

blend into solutions.

Attractions:

In creating solutions the attractions within

the pure substance must be broken

(solute from solute / solvent from

solvent) and allow the formation of

new attractions between unlike

particles (solute and solvent).

Electrolyte solutions : Ionic solids

dissolve individual ions in water

because the polar water molecules

can form solvent cages around the

charged cation or anion in order to

disperse the ions in the solution.

Soluble ionic compounds form electrolyte solutions while soluble

polar or nonpolar molecules are nonelectrolyte solutions.

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Solubility and Saturation:

• The solubility of a compound is the amount of the compound, usually in

grams, that dissolves in a certain amount of liquid, often 100 g of water.

The solubility of sodium chloride at 25 °C is 36 g NaCl per 100 g water.

• A saturated solution holds the maximum amount of solute under the

solution conditions. If additional solute is added, it will not dissolve.

• An unsaturated solution is holding less than the maximum amount of

solute. If additional solute is added, it will dissolve.

• A supersaturated solution is one holding more than the normal

maximum amount of solute. The solution is unstable, the solute will

normally precipitate from (or come out of) a supersaturated solution.

Any disturbance will cause the excess solute to come out.

Solubility

• Solubility rules give a qualitative description of the solubility of ionic

solids. For calcium carbonate, the attraction between Ca2+ ions and

CO32− ions is greater than solvent-solute attractions, and CaCO3 does not

dissolve in water (insoluble). The solubility of CaCO3is close to zero

grams per 100 g water. Ionic solids which dissolve nearly completely in

water (soluble) are strong electrolytes which break into ions: NaCl (aq)

Na+(aq) + Cl- (aq)

• Small polar molecular solids are

generally soluble in water. Table

sugar (C12H22O11) is polar and

soluble in water. Sugar molecules

stay as whole molecules and are

nonelectrolytes.

• Nonpolar molecular solids, such

as lard and vegetable shortening,

are insoluble in water.

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Temperature:

• Generally the solubility of solid solutes in water

increase at higher temperatures. Look at KNO3

on the solubility graph.

• Generally the solubility of gas solutes in water

decrease at higher T. CO2 gas in soda is more

soluble at cold temps than room temp.

Temperature and Solubility to Purify Compounds

One way to purify a solid is a technique called

recrystallization.

Solid is added to a solvent at higher temperature to create a saturated

solution. As the solution cools, the solubility decreases, causing some of

the solid to precipitate, forming crystals as it comes out. The crystalline

structure reject impurities, resulting in a purer solid.

Recrystallization is a way to make rock candy (sugar).

Pressure:

The solubility of gas in water increases as the pressure of that gas above

the liquid increases.

Henry’s Law… Solubility = k Pgas

Liquids exposed to air contain some dissolved gases. Lakes and seawater

contain dissolved oxygen necessary for the survival of fish. Our blood

contains dissolved nitrogen, oxygen, and carbon dioxide. Even tap water

contains dissolved atmospheric gases.

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Example 1:

a) Will C2H6 or CH3OH be more soluble in water? Explain.

b) Which is expected to be more soluble at higher temperatures

Choices: CO2 gas in water or KCl solid in water

c) Will CO2 gas be more soluble in water when the container is

pressurized with air or carbon dioxide?

Quantitative Concentrations:

molality (m) = moles of solute

kg of only the solvent

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Dilution of Solutions:

M1V1 = M2V2

Laboratory Safety Note:

When diluting acids, always add the

concentrated acid to the water.

Never add water to concentrated acid

solutions.

Stoichiometry of Solutions:

Conversions:

Convert volume (or other units given) of A to moles A

Convert moles A to moles B

Convert moles B to desired units (grams, volume, molecules...) B

Example 2:

Solve for the number of grams of sodium hydroxide (NaOH) required

to make 250 ml of 3.00 M NaOH solution.

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Example 3:

For 100 grams of aqueous solution that is 28.0% C2H5OH by mass and

has the density is 0.945 g/ml, solve for…

a) Grams of ethanol in 100 g of solution (mass %)

b) Moles of ethanol in 100 g of solution (MW)

c) Volume of 100 g of solution (density)

d) Molarity (mole/L)

Example 4:

How do you prepare 250 ml of 0.500 M H2SO4 from a stock solution

of 6.00 M H2SO4?

Example 5: (titration)

30.0 ml of 0.240 M NaOH is required to stoichiometrically react with

20.0 ml of an HCl solution. Write the balanced acid-base equation and

calculate and Molarity of the HCl solution?

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Colligative properties depend on amount of particles, not the type of solute

• Adding solute (impurities) to a liquid extends the temperature range in

which the liquid remains a liquid. The mixed solution has a lower

melting point and a higher boiling point than pure liquid; these effects are

called freezing point depression and boiling point elevation.

• For freezing point depression and boiling point elevation, the

concentration of the solution is expressed in molality (m), the number of

moles of solute per kilogram of solvent.

Freezing point depression:

The freezing point of a solution is lower than the pure solvent freezing

point.

Tf =iKf m Kf for water = 1.86 ˚C.kg/mol

i is the van’t hoff factor; i = 1 for nonelectrolytes, and a factor of the

number of particles that the compound breaks into for electrolytes:

NaCl (aq) i ≈ 2

ΔTf is the change in temperature of the freezing point in °C (compared

to the freezing point of the pure solvent).

m is the molality of the solution in (mol solute/kg solvent).

Kf is the freezing point depression constant for the solvent.

Different solvents have different values of Kf.

Boiling point elevation:

The boiling point of a solution is higher than the boiling point of the pure

solvent.

In automobiles, antifreeze not only prevents the freezing of coolant within

engine blocks in cold climates, but also prevents the boiling of engine

coolant in hot climates.

Tb =iKb m Kb for water = 0.512 ˚C.kg/mol

This equation is similar to the freezing point depression equation, ΔTb

is the change in temperature of the boiling point in °C, Kb is the

boiling point elevation constant for the solvent.

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Osmosis:

Osmosis is the flow of a solvent from a less concentrated solution to a

more concentrated solution in the attempt to equal the two concentrations.

Osmosis occurs when solutions containing a high concentration of solute

draw solvent from the lower concentration solution. As a result, fluid

rises on the higher concentration side until the weight of the excess fluid

creates enough pressure to stop the flow.

This pressure is the osmotic pressure of the solution.

Osmotic pressure is the pressure required to stop an osmotic flow.

Osmotic pressure is a colligative property depending only on the

concentration of the solute particles, not on the type of solute.

Why shouldn’t you drink seawater when lost at sea?

Membranes of living cells are semipermeable membranes. Seawater is

approximately 3.5% NaCl and cell tissues have the equivalent of 0.9%

NaCl solution. As seawater

flows through the stomach

and intestine, it draws water

out of cells, causing

dehydration.

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Red blood cells in solutions of different concentration

(a) In blood, the solute concentration of the surrounding fluid should be

equal to that within a blood cell, so there is no net osmotic flow, and the

red blood cell exhibits its typical shape.

(b) If a blood cell is placed in pure water and osmotic flow of water into

the cell causes it to swell up. Eventually, it may burst.

(c) If a blood cell is placed in a concentrated solution like seawater, then

osmosis draws water out of the cell, distorting its normal shape.

A) Intravenous solutions should have osmotic pressure equal to that of

bodily fluids. These solutions are isoosmotic.

B) Solutions having osmotic pressures less than bodily fluids are

hypoosmotic. These pump water into cells.

C) Solutions having osmotic pressures greater than bodily fluids are

hyperosmotic. These solutions take water out of cells and tissues.

When patients are given an IV in a hospital, the majority of the fluid is

usually an isoosmotic saline solution containing 0.9 g NaCl per 100 mL

of solution.