Name: Chapter 15 Solutions

1

Chapter 15 Solutions I. What are Solutions A. Solution - homogeneous mixture made up of individual molecules, atoms or ions.

B. Solute - the substance being ________________________

C. Solvent - the substance ________________________

D. Soluble - substance that _________________________________ in a solvent

E. Insoluble - substance that _____________________________________ in a solvent

F. Immiscible – liquids that are __________________________________ in each other

G. Miscible – two liquids that ___________________________________ in each other

H. Types of solutions 1. Aqueous solution -

____________________ is the solvent

2. Tincture - ____________ is the solvent

3. Alloys - solid solution of two or more

metals

- _____________ - an alloy in which one

metal is Hg

II. Electrolytes vs. Nonelectroyles A. Electrolytes: any substance that, when it is dissolved in a solution, will conduct an electric current by means of movement of ions

NaCl(aq) _____ 6M HCl(aq)______ 1M HCl(aq)___________

Glacial Acetic Acid(aq) _____6M HC2H3O2(aq) ______1M HC2H3O2(aq) _____

NaOH(aq) ____ tap H2O _____ distilled H2O _____

Strong Electrolytes - substance that dissolves almost completely in water to produce many ions to conduct electricity

1. NaCl(s)

2. HCl(g)

3. NaOH(s)

Name: Chapter 15 Solutions

2

Weak (or non) electrolyte - poor conductor of electricity

1. HC2H3O2 (aq)

2. H2 O (l)

III. Heterogeneous Mixtures 1. Suspension (NOT A SOLUTION) – a mixture containing particles that remain while thoroughly mixed, but then settles out when left undisturbed. Can you think of examples of a suspension? 2. Colloids (NOT A SOLUTION) –Colloids or colloidal dispersions are

suspensions in which the suspended particles are larger than molecules but too small to separate out of the suspension due to gravity.

Particle size: 10 to 2000 Å. Can you think of examples colloids? 3. Tyndall Effect Dispersed colloid particles that are large enough to scatter light. The path of a beam of light projected through a colloidal suspension can be seen through the suspension. Watch and observe the demo?

Hydrophilic and Hydrophobic Colloids

Water-loving colloids: hydrophilic. Water-hating colloids: hydrophobic. • In the human body, large biological

molecules such as proteins are kept in suspension by association with surrounding water molecules.

• These macromolecules fold up so that hydrophobic groups are away from the water (inside the folded molecule).

• Hydrophilic groups are on the surface of these molecules and interact with solvent (water) molecules.

•Typical hydrophilic groups are polar (containing C–O, O–H, N–H bonds) or charged.

Name: Chapter 15 Solutions

3

• Consider a small drop of oil in water. They will not mix unless soaps act on the oil and water.

• Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water.

• Most dirt stains on people and clothing are oil-based. • Biological application of this principle: • The gallbladder excretes a fluid called bile. • Bile contains substances (bile salts) that form an emulsion with fats in our small

intestine. • Emulsifying agents help form an emulsion. • Emulsification of dietary fats and fat-soluble vitamins is important in their

absorption and digestion by the body.

IV. The Solutions Process

Can you draw how you think a salt crystal dissolves? What do the ions of NaCl and molecules of water look like? How do they interact?

Consider NaCl (solute) dissolving in water (solvent). Water molecules

orient themselves on the NaCl crystals.

H-bonds between the water molecules have to be broken. NaCl dissociates into Na+ and Cl–. Ion-dipole forces form between the Na+ and the negative end of the water

dipole. Similar ion-dipole interactions form between the Cl– and the positive end of

the water dipole. . The process by which the charged particles in an ionic solid

separate from one another is called _____________________________

Solvation – the process of surrounding solute particles with solvent particles to form a solution. If water is the solvent, the interaction is called hydration. Factors that affect rate of solvation – must increase collisions between solute and solvent

1. ____________________________________ 2. ____________________________________ 3. ____________________________________

Dissociation Equation: NaCl (s)

Name: Chapter 15 Solutions

4

Solubility measure the maximum amount of solute that can be dissolved in a given amount of solvent at a specific temperature and pressure. usually expressed a quantity of solute per definite amount of solvent at a particular temperature (ex. 36g NaCl/100g H2O at 20 C) Unsaturated Saturated Supersaturated

Graphs and solubility tables are useful planning tools for making solutions. Can you draw the curve of a graph as a salt increases in solubility as temperature increases? Factors affecting solubility The tendency of a substance to dissolve in another depends on: 1. Nature of the solute and solvent example: at room temp, 1g of PbCl2/100g H2O and 200g ZnCl2/100g H2O Which solute is more soluble and why, can you explain?

Intermolecular forces are an important factor in determining

solubility of a solute in a solvent.

The stronger the attraction between solute and solvent molecules, the

greater the solubility.

For example, polar liquids tend to dissolve in polar solvents.

Favorable dipole-dipole interactions exist (solute-solute, solvent-solvent, and solute-

solvent).

the amount of solute

dissolved is less than the

maximum that could be

dissolved

solution holds the maximum

amount of solute per

amount of the solution

under the given conditions

solutions contain more solute

than the usual maximum amount

and are unstable

They cannot permanently hold

the excess solute in solution and

may release it suddenly

Solution Equilibrium is the physical state in which the opposing processes of dissolving and crystallization of a solute occur at equal rates.

Name: Chapter 15 Solutions

5

Water and ethanol are miscible because the broken hydrogen bonds in both pure liquids

are re-established in the mixture.

However, not all alcohols are miscible with water.

Why? The number of carbon atoms in a chain affects solubility.

The greater the number of carbons in the chain, the more the molecule behaves like a

hydrocarbon.

Thus, the more C atoms in the alcohol, the lower its solubility in water.

Increasing the number of –OH groups

within a molecule increases its

solubility in water.

The greater the number of –OH

groups along the chain, the more

solute-water H-bonding is possible.

Generalization: “Like dissolves like”. Substances with similar intermolecular

attractive forces tend to be soluble in one another.

• The more polar bonds in the molecule, the better it dissolves in a polar solvent.

•The less polar the molecule the less likely it is to dissolve in a polar solvent and the

more likely it is to dissolve in a nonpolar solvent.

• Network solids do not dissolve because the strong intermolecular forces in the

solid are not reestablished in any solution. 2. Pressure The solubility of a gas in a liquid is a function of the pressure of the

gas over the solution. a. solid in liquid?

Name: Chapter 15 Solutions

6

b. gas in liquid? What happen to when the pressure goes up?

c. Henry’s Law

3. Temperature ___ temp, ____ kinetic energy, ____ randomness a. solid in liquid? what if cools back down? b. gas in liquid? What happen to a gas as the liquid warms up? Heat of Solution: overall energy change that occurs during the solution process

Energy Changes and Solution Formation

There are three steps involving energy in the formation

of a solution:

•Separation of solute molecules (ΔH1),

•Separation of solvent molecules (ΔH2), and

•Formation of solute-solvent interactions (ΔH3).

•We define the enthalpy change in the solution process

as:

ΔHsoln = ΔH1 + ΔH2 +ΔH3

• To determine whether ΔHsoln is positive or negative, we

consider the strengths of all solute-solute, solvent-

solvent and solute-solvent interactions:

• Breaking attractive intermolecular forces is always

endothermic.

• ΔH1 and ΔH2 are both positive.

• Forming attractive intermolecular forces is always exothermic.

• ΔH3 is always negative.

At a given temperature the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid.

S =

P =

Name: Chapter 15 Solutions

7

• It is possible to have either ΔH3 > (ΔH1

+ ΔH2) or ΔH3 < (ΔH1 + ΔH2).

Examples:

MgSO4 added to water has ΔHsoln = –

91.2 kJ/mol.

NH4NO3 added to water has ΔHsoln = +

26.4 kJ/mol.

MgSO4 is often used in instant heat

packs and NH4NO3 is often used in

instant cold packs.

How can we predict if a solution will form?

•In general, solutions form if the ΔHsoln is negative.

•If ΔHsoln is too endothermic a solution will not form.

•“Rule of thumb”: Polar solvents dissolve polar solutes.

•Nonpolar solvents dissolve nonpolar solutes.

•Consider the process of mixing NaCl in gasoline.

•Only weak interactions are possible because gasoline is nonpolar.

•These interactions do not compensate for the separation of ions from one another.

•Result: NaCl doesn't dissolve to any great extent in gasoline.

•Consider the process of mixing water in octane (C8H18).

•Water has strong H-bonds.

•The energy required to break these H-bonds is not compensated for by

interactions between water and octane.

•Result: Water and octane do not mix. Factors affecting rate of solution 1. size of particle a. dissolving only occurs at surface b. smaller pieces, ___ surface area, ___ rate of solution 2. stirring a. brings fresh parts of solvent in contact with solute, so ___ rate of sol 3. amount of solute already dissolved a. the more solute dissolved, ______ rate of solution 4. temperature a. solids: _____ temp, _____ rate of solution b. gases: _____ temp, _____ rate of solution

Name: Chapter 15 Solutions

8

Ways of Expressing Concentration

• All methods involve quantifying the amount of solute per amount of solvent (or solution). • Concentration may be expressed qualitatively or quantitatively.

• The terms dilute and concentrated are qualitative ways to describe concentration.

• A dilute solution has a relatively small concentration of solute.

• A concentrated solution has a relatively high concentration of solute.

• Quantitative expressions of concentration require specific information regarding such

quantities as masses, moles, or liters of the solute, solvent, or solution.

Mass percentage, ppm, and ppb

• Mass percentage is one of the simplest ways to express concentration.

•By definition:

• Similarly, parts per million (ppm) can be expressed as the number of mg of solute per

kilogram of solution.

• If the density of the solution is 1g/ml, then 1 ppm = 1 mg solute per liter of solution.

• We can extend this again!

Solution Concentrations what does concentration mean?

100solution of mass total

solnin component of masscomponent of % Mass

610solution of mass total

solnin component of masscomponent of (ppm)million per Parts

Name: Chapter 15 Solutions

9

• Parts per billion (ppb) can be expressed as the number of µg of solute per kilogram of solution.

• If the density of the solution is 1g/ml, then 1 ppb = 1 µg solute per liter of solution.

Mole Fraction, Molarity, and Molality

• Common expressions of concentration are based on the number of moles of one or more

components.

• Recall that mass can be converted to moles using the molar mass. • Recall:

• Note: Mole fraction has no units. • Note: Mole fractions range from 0 to 1.

• Recall:

example: What is the molarity of a solution which contains 4.9 g of H2SO4 in 250 ml of solution? example #2: How many moles of KNO3 are contained in 300 ml of a 0.50 M solution?

• Note: Molarity will change with a change in temperature (as the solution volume increases or

decreases). • We can define molality (m), yet another concentration unit:

• Molality does not vary with temperature. • Note: converting between molarity (M) and molality

(m) requires density.

Colligative Properties:

Those people who live in the colder climates of the world add antifreeze to the water in their automobiles radiators. From the name of the product, it is obvious that its purpose is to prevent the water from freezing. How does the antifreeze prevent freezing? It is not that the antifreeze itself has such a low freezing point. The best protection is afforded by a mixture that is roughly 50% antifreeze and 50% water. Next we will investigate the behavior of solutions in such situations as freezing

M (molarity) # of moles of solute 1 Liter of solution

=

componentsallofmolestotal

componentofmoles,componentoffractionMole X

solvent of kilograms

solute of moles Molality, m

910solution of mass total

solnin component of masscomponent of (ppb)billion per Parts

Name: Chapter 15 Solutions

10

and boiling. We will look at those properties of solvents that are changed by solutes. Properties which are determined by the number of particles in solution rather than by the type of particle in solution, are colligative properties.

Freezing point depression and boiling point elevation A. Vapor Pressure Lowering: 1. VP: the pressure exerted in a closed container by liquid particles that have escaped the liquid’s

surface and have entered the gaseous state. 2. Vapor pressure for water shown in first beaker, a solute lowers the vapor pressure in the

second beaker. 3. There are fewer solvent particles on the surface of the second beaker to enter the gaseous

state and the vapor pressure is lowered.

4. Remember Boiling Point is defined as:

B. Boiling point elevation and freezing point depressions depend on: 1. molal concentration - number of moles of solute in a kilogram of solvent 2. nature of the solvent

Can you think of examples of a BP elevation or FP depression?

Example: #2 What would lower the freezing point better on the roads, sugar (C12H22O11), table salt (NaCl), another salt CaCl2, and why?

C. Molal freezing and boiling point constants

1. A 1 molal solution will lower the freezing point of water 1.86 C. Water

could freeze at -1.86 C instead of 0 C. This value is called the molal freezing point constant for water (1.86 C/molal).

Tf = change in freezing point

kf = constant for water

m = molal concentration of

solution 2. A 1 molal solution will raise the

point of water 0.52 C. boiling

Examples: If 85.0 grams of

sugar are dissolve in 392 grams of water, what will be the boiling point and freezing points of the resulting solution? The molecular formula of sugar is

C12H22O11. (answer = 100.33

Tf = kfm

Tb = Kbm

Name: Chapter 15 Solutions

11

C, -1.18 C)

Lowering the Vapor Pressure

Nonvolatile solutes (with no measurable vapor pressure) reduce the ability of the surface solvent

molecules to escape the liquid.

• Therefore, vapor pressure is lowered. • The amount of vapor pressure lowering depends on the amount of solute.

•

Osmosis

•Semipermeable membranes permit passage of some components of a solution.

•Often they permit passage of water but not larger molecules or ions.

•Examples of semipermeable membranes are cell membranes and cellophane.

• Osmosis is the net movement of a solvent from an area of low solute

concentration to an area of high solute concentration.

•Consider a U-shaped tube with a two liquids separated by a semipermeable membrane.

One arm of the tube contains pure solvent.

The other arm contains a solution.

•There is movement of solvent in both directions across a

semipermeable membrane. •The rate of movement of solvent from the pure solvent to the solution

is faster than the rate of movement in the opposite direction.

•As solvent moves across the membrane, the fluid levels in the arms

become uneven.

•The vapor pressure of solvent is higher in the arm with pure solvent. •Eventually the pressure difference due to the difference in height of liquid

in the arms stops osmosis.

• Osmotic pressure, , is the pressure required to prevent

osmosis.

• Everyday examples of osmosis include:

•If a cucumber is placed in NaCl solution, it will lose water to shrivel up

and become a pickle.

•A limp carrot placed in water becomes firm because water enters via

osmosis. •Eating large quantities of salty food causes retention of water and

swelling of tissues (edema).

•Water moves into plants, to a great extent, through osmosis. •Salt may be added to meat (or sugar added to fruit) as a

preservative.

Name: Chapter 15 Solutions

12

•Salt prevents bacterial infection: A bacterium placed on the salt will lose water through osmosis and

die.