Solutions

SolutionsI. Classification of liquid mixtures.II. Solution Concentration. Molarity.III. Solubility.

-Dissolving process.-Ionic Equations.

IV. Colligative properties.-Types of solutions.-Solubility curve.-Molality.-Freezing/Boiling Points.

Classification of Matter

Solutions are homogeneous mixtures

Properties of Solutions, Suspensions, and

Colloids

Solutions Solutions: a

homogeneous mixture of two or more substances in a single phase of matter.

In a simple solution where, for example, salt is dissolved in water, the particles of one substance are randomly mixed with the particles of another substance.

SOLUTE – A solute is the dissolved substance in a solution.

EX: CO2, KCl, Na2CO3

SOLVENT – A solvent is the dissolving medium in a solution.

EX: H2O, CCl4

The solute is generally designated as that component of a solution that is of lesser quantity.

If we had a mixture of 25 mL of ethanol and 75 mL of water, the ethanol would be the solute and water would be the solvent.

If we had a 50% to 50% ratio, it would be unnecessary to designate solvent or solution.

There are two types:1. Suspensions2. Colloids

*Suspensions and colloids are not solutions.

Heterogeneous Liquid Mixtures

Suspensions If the particles in a solvent are so

large that they settle out due to gravity unless the mixture is constantly stirred or agitated, the mixture is called a suspension.

These particles (over 1000nm) can be filtered out of the heterogeneous mixture.

Colloids Particles that are

intermediate in size between those in solutions and suspensions form mixtures known as colloidal dispersions.

Particles between 1nm and 1000 nm in diameter may form colloids. After the larger particles settle out (suspensions), the water may still be cloudy because colloidal particles remain dispersed in the water.

Milk is an example of a colloid.

Suspensions Colloids Solutions

The Tyndall Effect

Many colloids appear homogeneous because the individual particles cannot be seen. The particles are, however, large enough to scatter light.

Tyndall effect is a property that can be used to distinguish between a solution and a colloid.

When a laser is passed through a solution and a colloid at the same

time, it is evident which glass contains the colloid. (you can’t see the light

in a colloid)

Colloid Solution

Solutions Colloids SuspensionsHomogeneous Heterogeneous HeterogeneousParticle size: 0.01-1 nm; can be atoms, ions, molecules

Particle size:1-1000 nm, dispersed; can be large molecules

Particle size:Over 1000 nm, suspended; can be large particles

Do not separate on standing

Do not separate on standing

Particles settle out

Cannot be separated by filtration

Most cannot be separated by filtration

All can be separated by filtration

Do not scatter light

Scatter light(Tyndall effect)

Not transparentMay scatter lite

Solution Concentration

Solution Concentration Molarity is simply a measure of the

"strength" of a solution.  A solution that we would call "strong" would have a higher molarity than one that we would call "weak".

If you ever made or drank a liquid made from a powdered mix, such as Kool-Aid or hot cocoa, you probably are familiar with the difference between what is called a "weak" solution or a "strong" solution. 

To make Kool-Aid of "normal“ strength =

4 scoops of powder-----------------------2 quarts of water

                                                                       

To make Kool-Aid twice the "normal" strength… What could you do?

                                                               8 scoops of powder  4 scoops of

powder -------------------- or --------------------  

2 quarts of water   1 quart water

Solution Concentration

Molarity:*Molarity (M) = moles solute

Liter(s) solution

One mole of NaCl (molar mass of NaCl =

22.99 + 35.45 = 58.44 grams) is dissolved in enough water (1 Liter) to make a 1M NaCl solution.

Molarity Calculations Calculate the molarity of:

*35.2 grams of CO2 in 500. mL.M = moles solute/L solution

Step 1 (moles solute): • convert 35.2 g of CO2 into moles

35.2g x (1 mole CO2) = 0.800 mol (44.01 g)Step 2 (Liters of solution): convert mL to L 500. mL x 1 L = 0.500 L 1000 mLStep 3: divide moles by volume in liters

0.800 mol CO2 = 1.60 M CO2

0.500 L

AnswersWS: Molarity Problems

1) NaCl

2) Al2(SO4)3

3) HClO3

4) HCl

5) Ba(OH)2

6) Fe(NO3)2

1) 1.223 M KClO3

2) 0.774 M Na2SO4

3) 1.00 M NaOH

4) 0.655 M AlCl3

5) 0.71 M HCl

6) 0.82 M LiF

7) 2.8 M KOH

8) 0.100 M ZnCl2

Solutions and Solubility

When we talk about the mixing of two or more substances together in solution was much consider solubility.

Solubility is defined as the amount of a substance that can be dissolved in a given quantity of solvent.

Solubility

When deciding what type of solvent to use with a given solute it is important to identify what types of substances you have.1. Polar substances (partial + or – charges) tend to dissolve in polar solvents2. Nonpolar molecules (equal sharing of e-) tend to dissolve in nonpolar solvents

Remember the rule:

LIKE DISSOLVES LIKE

Solvent-Solute Combinations

Remember:

LIKE DISSOLVES LIKE

Solvent Type Solute Type Is solution likely?

Polar Polar Yes

Polar Nonpolar No

Nonpolar Polar No

Nonpolar Nonpolar Yes

The Dissolving Process

Water is a polar solvent and is attracted to polar solutes.

Salt is polar (ionic). Water molecules surround and isolate

the surface ions. The ions become hydrated and move away from each other in a process called dissociation.

Insolubility

Any substance whose solubility is less than 0.01 mol/L will be referred to as insoluble.

We can predict whether a precipitate (insoluble substance) will form when solutions are mixed if we know the solubilities of different substances.

Experimental observations have led to the development of a set of empirical solubility rules for ionic compounds, gases, and molecules.

EX: Experiments demonstrate that all ionic compounds that contain the nitrate anion, NO3

-, are soluble in water.

*Refer to solubility table*

Ionic Equations

Referring to Solubility Tables & Writing Ionic Equations are very useful tools when trying to determine if a

reaction will occur in an aqueous solution.

Ionic Equations: An ionic equation is a chemical

equation in which electrolytes (soluble ions that conduct electricity) are written as dissociated ions.

Ionic equations are used for single and double replacement reactions which occur in aqueous solutions.

In an aqueous reaction ions that are found as both reactants and products are not part of a reaction. They are termed spectator ions and essentially cancel out of the ionic equation.

To write net ionic equations follow these simple rules:

1. Write a balanced equation.2. Determine which substances

are soluble (refer to the solubility rules table)

3. Rewrite the equation in ionic form by dissociating the soluble reactants & products

4. Cancel the spectator ions5. Write the net ionic equation

Hints… What will dissociate?Refer to your Solubility Rules Table

Salts:• Write in ionic form if soluble.

EX: KCl K+ + Cl-

• Write in undissociated form if insoluble. EX: AgCl AgC

Acids:• STRONG ACIDS - Write in ionic form.

They are soluble. (listed on table)

EX: H2SO4 2H+ + SO4

-2

• WEAK ACIDS - Write in undissociated form. They are insoluble. (not listed on table)

EX: H3PO4 H3PO4

Bases (OH-):• STRONG BASES - Write in ionic form.

They are soluble.EX: Ca(OH)2 Ca+2 +

2OH-

• WEAK BASES - Write in undissociated form. They are insoluble.

EX: Mg(OH)2 Mg(OH)2

Oxides:• Always write in undissociated form.

EX: MgO, H2O

Gases:• Always write in undissociated form.

EX: SO2, NH3, H2, O2

Step 1 & 2:AgNO3(aq) + NaCl(aq) AgCl(s) +

NaNO3(aq)

(s) (s) (s)

Step 3 & 4:Ag+ + NO3

- + Na++ Cl- AgCl(s) + Na+ + NO3

-

Step 5:Ag+ + Cl- AgCl(s)

Sample Problem

Types of Solutions

Saturated Solutions A solution at equilibrium with

undissolved solute is said to be saturated.

Additional solute will not dissolve if added to this solution.

It is possible to dissolve less solute than needed to form a saturated solution. These solutions are unsaturated.

A supersaturated solution can be made by dissolving the solute under high temps and then carefully cooling them. These are unstable solutions.

Degrees of Saturation1. Saturated solution Solvent holds as much

solute as is possible at that temperature.

Undissolved solid remains in flask.

Dissolved solute is in equilibrium with solid solute particles.

Less than the maximum amount of solute for that temperature is dissolved in the solvent.

No solid remains in flask.

2. Unsaturated Solution

3. Supersaturated

Solvent holds more solute than is normally possible at that temperature.

These solutions are unstable; crystallization can often be stimulated by adding a “seed crystal” or scratching the side of the flask.

Solids form as solution cools.

Saturated Solution

Supersaturated Solution

Factors Affecting Solubility

Solubility depends on the nature of both the solvents and solutes, temperature, and for gases, on pressure.

The solubility of most solid solutes in water increases as the temp of the solution increases.

This means that more sucrose C12H22O11 can be dissolved in hot water than cold, the basis for making “rock candy”.

The graph represents the solubility of substances, including NaCl, NaNO3, and KNO3 at different temps.

Notice that whentemp increases, solubility increases

for most substances.

Solubility Curves

Solubility of Gases- Temp Solubility of gases

is dependent on both temperature and pressure.

Based on the solubility curve, the solubility of NH3 and SO2 (both gases) decreases as temperature increases.

Gas Solubility increases when pressure increases.

Carbonated beverages are bottled with CO2 under pressure to increase the solubility of CO2 gas.

As bottle is opened, pressure of CO2 decreases and solubility of CO2 decreases.

Therefore, bubbles of CO2 escape from solution.

Temp also effects solubility, colder sodas lose CO2 more slowly than warm sodas.

Solubility of Gases- Pressure

Dissolving ChartSolid Gas

Increasing Solution Temperature

Increases solubility & solubility rate

Decreases solubility

Crushing Solute Increases solubility rate

No Effect

Stirring Solution Increases solubility rate

Increases solubility rate

Increasing Atmospheric (air) Pressure

No EffectIncreases solubility & solubility rate

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Colligative Properties

Colligative comes from the Greek word kolligativ meaning glue together.

We use this term for the properties of substances (solutes and solvents) together.

Colligative properties of solutions is used to describe the effects of antifreeze/summer coolant.

Molality Recall the units for Molarity (M):

moles solute L solutionMolality (m) is the measure of the number

of moles of a solute per 1000g of solvent.moles solute

1kg solvent Molality is best used to describe colligitive

properties and is represented by m.

Boiling Point and Freezing Point Review the

phase diagram of a pure substance.

How will the phase diagram of a solution (freezing and boiling points) differ from those of a pure solvent?

The addition of a nonvolatile solute will require a higher temperature in which to reach boiling point, thus:

Boiling point elevation The addition of a nonvolatile solute

will require a lower temperature in which to reach freezing point, thus

Freezing point depression

Pure water Water with NaCl

The water with the solute of NaCl has fewer liquid molecules becoming gases.

This will increase the temp needed to change the state from (l) (g)

Calculating Freezing and Boiling Points

The following table contains the molal (K) Boiling Point Elevations, Kb, and Freezing Point Depressions, Kf.

Solvent Normal boiling pt

(°C)

Kb

(°C/m)

Normal freezing pt

(°C)

Kf

(°C/m)

Water 100.0 0.52 0.0 1.86Benzene 80.1 2.53 5.5 5.12Ethanol 78.4 1.22 -114.6 1.99CCl4 76.8 5.02 -22.3 29.8

Chloroform 61.2 3.63 -63.5 4.68

The data for the table was found by doing experiments.

It has been found that 1 mole of a nonvolatile solute particles will raise the boiling temperatures of 1 kg of water by 0.52 C°.

The same concentration of solute will lower the freezing point of 1 kg of water by 1.86 C°.

These two figures are the molal boiling point constant (Kb) and the molal freezing point constant (Kf).

A 1m solution of sugar in water contains 1 mole of solute particles per 1 kg of solvent.

A 1m solution of NaCl in water contains 2 mole of solute (because NaCl is an ion, it will dissociate in water into Na+ and Cl- ions) per 1 kg of solvent.

How many mole solute would 1m calcium nitrate, Ca(NO3)2, have per 1kg solvent?

Calculating Changes in Kb and Kf Boiling point elevation is:

ΔTb = Kbmi (molality)

(change in boiling point) (boiling point constant) Freezing point depression:

ΔTf = Kfmi (molality)

(change in freezing point) (freezing point constant)

(moles)

(moles)

If 55.0 grams of glucose (C6H12O6) are dissolved in 525 g of water, what will be the change in boiling and freezing points of the resulting solution?

Step 1: Calculate molality: 55.0 g ( 1 mol ) 0.305 mol =

1 (180.18 g) = 0.525kg0.581 m

Step 2: Obtain molal Kb from table. Step 3: Place values into equation ΔTb = Kbmi

ΔTb = (0.52°C/m)(0.581m)(1) = 0.302 °C

This means that the boiling point will be elevated by 0.302 °C.

Normal Boiling Point + ΔTb = New Boiling Point

100 °C + 0.302 °C = 100.302 °C

This solution will reach boiling point at 100.302 °C.

Now let’s calculate the change in freezing.

Calculate the change in freezing point of 24.5g potassium bromide dissolved in 445 g of water. (assume 100% dissociation)

Step 1: Convert g of KBr into moles24.5g ( 1 mol) 0.206mol =

1 (119.00 g) = 0.445kg0.463m

• Step 2: Obtain molal Kf from table.• Step 3: Place values into equation

• KBr is ionic so the dissociation of KBr makes 2 moles of ions (solute) per kg of solvent:

KBr K+ + Br –

ΔTf = Kfmi

ΔTf = (1.86°C/m)(0.926m)(2) = 1.72 °C

Normal Freezing Point - ΔTf = New Freezing Point

0 °C - 1.72 °C = -1.72 °C *Freezing point has been depressed to -1.72 °C.

Coolant is used because it takes higher temperatures to reach boiling point.

Antifreeze needs lower temperatures in order to freeze.

This also why salt is used on frozen roads and walkways. The salt dissolves in the water and lowers the freezing point of water. It now takes colder temps to turn the water into ice.

A 10-percent salt solution freezes at 20 F (-6 C), and a 20-percent solution freezes at 2 F (-16 C).

Boiling point elevation is:ΔTb = Kbmi

New Boiling Point = normal bp + ΔTb

Freezing point depression:ΔTf = Kfmi

New Freezing Point = normal fp - ΔTf

Solvent Normal boiling pt (°C)

Kb

(°C/m)

Normal freezing pt (°C)

Kf

(°C/m)

Water 100.0 0.52 0.0 1.86

Benzene 80.1 2.53 5.5 5.12

Ethanol 78.4 1.22 -114.6 1.99

CCl4 76.8 5.02 -22.3 29.8

Chloroform 61.2 3.63 -63.5 4.68

Practice problems: Compute both boiling and freezing points of these solutions: (assume 100% dissociation of all ionic compounds)

1. 27.6 g NaBr in 100.g of water.2. 100.0 g of C10H8 (naphthalene) in

250. g of C6H6 (benzene).

3. 25.9 g of C7H14BrNO4 (3-bromo-2-nitrobenzoic acid) in 150. g of benzene.

4. 55.6 g of C12H22O11 in 500. g of water.

5. 1500.g of NaCl in 4500. g of water.

Brief Summary Heterogeneous liquid mixtures are

classified as suspensions (large particles that settle out), or colloids (small particles that stay dispersed).

Homogeneous mixtures are solutions made of a solute dissolved in a solvent.

Solutes and solvents must be alike in polarity in order to produce a solution.

The concentration of a solution is molarity (molar) and has the unit M, which includes moles of solute per unit volume of solvent.

When preparing a dilute solution from a concentrated solution, use the formula:

M1V1 = M2V2

Where initial volume and molarity of concentrated solution (EX: 12M HCl) is compared to final volume and molarity of diluted solution (EX: 6M HCl).

Solubility of solutes can be reflected in a solubility graph.

Solubility of solid substances generally increases as temperature increases.

Solubility of gases decreases with increased temperature.

Ionic equations can be written to express the net reaction occurring in a system after the spectator ions have been removed.

Solubility rules for substances have been experimentally determined. They indicate what substances are or are not water soluble.

Colligative properties demonstrate the properties of the solution rather than solute and solvent independently.

A solution with undissolved solute is termed unsaturated.

A solution with undissolved solute is termed saturated.

A solution that has more dissolved solute at a particular temp due to being dissolved at a higher temp is termed supersaturated.

Boiling points and Freezing points of solutions can be calculated using molality.

Molality (molal) is described by unit m and expresses moles of solute per kg of solvent.

When calculating BP and FP differences use equation: Kfp or bp = Kb or f x m x moles Kb or f is a standard and must be given

m must be calculated and adjusted to express moles contributed.-molecules contribute only 1 mol.-ionic compounds contribute the number of moles they dissociate into.EX: KI 1 mole K+ + 1 mole I-

After calculating difference, refer to normal BP and FP of solvents and:Boiling Point Elevation add difference to normal BP.Freezing Point Depression subtract difference from normal FP.