Solutions Part I

I. States of Matter

______ – molecules/particles in chaotic motion whose average kinetic energy is greater than the attractive forces between molecules

______ – attractive intermolecular forces between molecules is comparable to kinetic energies of molecules: these molecules are held in close proximity, but still move in a rather chaotic motion.

______ – intermolecular forces are sufficiently strong relative to kinetic energy. Molecules are virtually locked in place. Often this is very orderly (crystalline structures).

II. Phase Changes

1. vaporization (+H); liquid to gas; H2O(l)  H2O(g) absorbs 40.7kJ/mol therefore H=+40.7kJ/mol

2. ______(-H); gas  liquid

3. ______(+H); solid  gas

4. ______(-H); gas  solid

5. melting (+H); solid  liquid

6. freezing (-H); liquid  solid

III. Characteristics/Descriptions of Liquids

We will start this section by discussing the very unique properties of water:

(1) it covers ______of the earth

(2) it is the foundation of all ______things

(3) it contains 2 very polar bonds because

(a) difference in ______is 1.24 therefore the H-O bond is polar covalent

(b) the molecule itself is very polar due to the 2 unshared pairs of electron which cause it to bend

(c) water molecules are held together by ______types of intermolecular forces: dipole-dipole and hydrogen bonding

(d) it is the combination of shape (______) and IM forces (dipole-dipole and H-bonding) which is responsible for all the other properties of water that follow.

(4) ______ Pressure – the pressure exerted by a gas above a liquid in a closed container. Vapor pressure increases with increasing temperature. Because the rate of evaporation increases with increasing temperature, vapor pressures of liquids always increases as temperature increases. Vapor pressure decreases with increased IM forces. Water has low vapor pressure. It is good that water has a low vapor pressure, otherwise what would happen to the rivers & lakes?

An easy way to remember vapor pressure is to ask yourself how fast the liquid evaporates, the faster the rate of evaporation, the higher the vapor pressure. Example: water and rubbing alcohol

(5) ______ Point – when the vapor pressure of the liquid equals the atmospheric pressure the liquid boils. Water boils at l00.C. Water has a relatively high boiling point. Draw water’s dot structure and determine its molar mass:

What accounts for the difference in boiling points between water and H2S, with a boiling point of –60.0 C?

Other facts about boiling points for all liquids:

a) normal boiling point is the boiling point of a substance at l atm of ______

b) boiling point is unique for each liquid; it does not depend on volume or ______area

c) boiling point increases with increased molecular mass if IM forces are the same

d) boiling point increases with increased IM forces

(6) ______ Tension – a measure of the inward forces that must be overcome to expand the surface area of a liquid. The “skin;” surface tension is responsible for bugs walking on water, a pin floating on water, etc. Water has relatively high surface tension. The greater the surface tension, the more spherical the drop of liquid will be.

Ex: surface tension of water vs. hexane

Surfactant – SURFAce ACTive ageNT – a substance that breaks down surface tension: detergents, _____

(7) ______ Capacity (or specific heat) – the amount of heat needed to change the temperature of 1.0 g of a substance by l.0°C. Because water has such a high heat capacity, it is cooler at the beach on a warm day. The water absorbs the heat and lowers the air temperature.

(8) Heat of Vaporization – heat needed to change l.0 g of liquid to a vapor at its boiling point. Water has a high heat of ______.

Other facts about water:

1) solid water (ice) is less dense than liquid water. Why? Its molecules expand as it freezes. At about 4.0°C, water begins to freeze and the molecules get farther apart and thus, less dense. That’s why ice floats on water.

2) water ______most ionic solids and polar liquids

3) water is sometimes present as water of hydration (hydrates)

CuSO4 5H20 – copper(II) sulfate pentahydrate Heating this substance will drive the water out of the ionic crystal.

IV. Types of Liquid Mixtures (Solutions, Suspensions, and Colloids)

1. Solution (homogeneous mixture)- any substance (solid, liquid, gas) that is evenly dispersed throughout another substance. page 398 (Not the same as a chemical reaction!!)

Ex: sugar water, salt water (do not scatter light)

Components of a Solution

1. Solute – substance dissolved

2. Solvent – substance that does the dissolving (water is the universal solvent)

2. Suspensions (heterogeneous mixtures) – particles in a solvent are so large that they settle out unless the mixture is constantly stirred Ex: muddy water, vegetable soup, page 398 (may scatter light, but are transparent)

3. Colloids (heterogeneous mixtures) – particles are intermediate in size between those in solutions and suspensions. Example: After large soil particles settle out of muddy water the water is often still cloudy because colloidal particles remain dispersed in the water. Ex: milk, mayonnaise , page 398 (do scatter light – Tyndall Effect)

V. The Solution Process (Solvation)

Solvation is the process by which a solute dissolves in a solvent.

Miscible: when solutes and solvents are soluble in each other (solvation occurs)

Immiscible: when solutes and solvents are not soluble in each other (solvation does not occur)

Aqueous solutions – solvent is water.

What happens when:

Ionic molecules are the solute / Polar covalent molecules are the solute / Nonpolar covalent molecules are the solute

and water is the solvent?

/ and water is the solvent? / and water is the solvent?
Dissociation of ionic molecules occurs (ions separate). Water is then attracted to the positive and negative ions. When all molecules have been “surrounded” the molecule is called hydrated.
Miscible / Dissociation does NOT occur. Water is polar and its “oppositely charged poles” will be attracted to other polar molecules' “oppositely charged poles.” When a solution is made between two polar molecules it is called molecular solvation.
Miscible / A solution will NOT occur. Water and any nonpolar molecule will not mix! Think of putting water and oil together. Water is polar and oil is nonpolar. The polar water is not attracted to the oil, because the nonpolar oil does not have any oppositely charged poles! Immiscible

VI. Like Dissolves Like

We don’t always use water as the solvent! Solutions can be made from various substances – a rule of thumb to follow when trying to determine if two substances will form a solution is “like dissolves like.”

Polar Molecules + Polar Molecules

Nonpolar Molecules + Nonpolar Molecules

Ionic and Ionic

Ionic + Polar Molecules

Polar Molecules + Nonpolar Molecules

Ionic Molecules + Nonpolar Molecules

VII. Solubility

Solubility: the maximum amount of a substance that will dissolve in a solvent (at a specific temperature)

According to solubility, solutions can be either:

  1. unsaturated – a solution that is able to dissolve more solute (not enough)
  2. saturated – a solution that cannot dissolve any more solute (just enough)
  3. supersaturated – a solution that contains more solute than can be dissolved (too much!!)

The solubility of substances varies widely. For example 0.189 grams of Ca(OH)2 dissolves in 100 grams of water at 0C. 122 grams of AgNO3 dissolves in 100 grams of water at 0C. (page 404 in your book)

VIII. Factors Effecting Rate and Solubility

A. Factors Effecting Rate:

1. Agitation – stirring or mixing the solution will increase the rate or how fast the solute dissolves, but it will not change how much solute can be dissolved. If you add 35.9 grams of salt to water (at 20C) it will all eventually dissolve, but if you stir the solution it will dissolve much quicker. (As you stir the particles are constantly being moved, allowing for interactions between solute and solvent to occur more quickly.)

2. Surface Area – increasing the surface area of the solute will increase the rate or how fast the solute dissolves, but it will not change how much solute can be dissolved.

3. Temperature – increasing temperature will increase the rate or how fast the solute dissolves in the solvent. (As temperature increases the particles begin to move faster and faster and collide with more particles quicker, which means the solute and solvent particles have an increased chance of coming into contact with each other.)

B. Factors Effecting Solubility:

1. Increasing Temperature - solubility of a solid solute in a liquid solvent generally increases with an increase in temperature. At 20C 35.9 grams of salt will dissolve in 100 grams of water, but at 100 C 39.2 grams of salt will dissolve in 100 grams of water!

2. Decreasing Temperature-increases the solubility of a gaseous solute in a liquid solvent. What would you rather drink, a hot coke or a cold coke?

3. Pressure – The solubility of a gas increases as the pressure of the gas above the liquid increases. Carbonated drinks have CO2 dissolved in them. They are also bottled under a high pressure of CO2, which forces the CO2 into solution. When the bottle is opened, the pressure above the solution decreases, and bubbles of CO2 form in the liquid, then escape. Eventually, most of the CO2 escapes and the drink becomes “flat.”

Henry’s Law- “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.”

S1 = S2

P1 P2