POLAR BONDS AND

MOLECULES

NOTES 16.3

Covalent Bonds

bond in which two atoms

share a pair of electrons.

1. Single bond = 1 shared pair of

electron

2. Double bonds = 2 shared

pairs

3. Triple bonds = 3 shared pairs

Bond Polarity:

1. bonding pairs of electrons in

covalent bonds are pulled

between the nuclei of the

atoms sharing the electrons.

2. When bonds are pulled

equally the bond is a nonpolar

covalent (occurs between like

atoms.)

3. When a covalent bond occurs

between different atoms then

electrons are shared unequally

which is a polar bond.

a. This leads to partial charges

on the atoms (- or + )

b. The polarity of a bond can be

shown with an arrow pointing

to the negative side.

c. Table 16.4 p. 462

Polar Molecules

- when one end of a molecule is

slightly negative and other is

slightly positive.

- Water is polar because the

way the bonds cause the

molecule to bend. You have

partial charges on two sides of

molecule.

- CO2 is not polar because the

double bonds keep the

molecule linear and the

charges cancel.

Attractions

molecules are attracted to

each other through a variety of

forces.

1. These forces are responsible

for determining whether a

compound is a gas, liquid or

solid.

2. Van der Waals forces consist

of two possible types.

a. Dispersion caused by the

movement of electrons.

Increases as the # of

electrons increase.

b. Dipole Interactions

occurs when polar

molecules are attracted to

one another.

1. Hydrogen bonds occur when H

already in a polar compound

bond with a partial negative of

another molecule.

- extremely important in

determining the properties of

water and biological

molecules such as proteins.

Intermolecular Attractions

The physical properties of a

compound depend on the type

of bonding it displays.

WATER AND

AQUEOUS

SYSTEMS

NOTES 17.1

Water Molecule:

1. It is a triatomic molecule with

two polar covalent bonds

(H – O).

2. It has a bent shape leading to

a partially + (δ+) and – (δ)

ends to the molecule.

3. Because of the polarity the

molecule will form Hydrogen

bonds with other water

molecules.

Water Molecule -- Polarity

- Water is polar because the

way the bonds cause the

molecule to bend. You have

partial charges on two sides of

molecule.

H – bonding gives H2O many

of its properties.

1) high surface tension

2) low vapor pressure

3) high specific heat

4) high boiling point

Surface Properties:

1. Water molecules experience

an uneven attraction the

molecules are hydrogen-

bonded on only one side of

the drop. The molecules pull

toward the body of the liquid.

2. This pull is called surface

tension.

3. A liquid that has strong

intermolecular forces has high

surface tension.

4. You can decrease surface

tension by adding a

surfactant, an agent such as

soap that interferes with the

hydrogen bonds.

Specific Heat Capacity

water has a high specific heat

capacity. Ability to absorb heat

without changing temperatures.

WATER VAPOR AND

ICE

NOTES 17.2

Evaporation and Condensation:

1. The hydrogen bonds of water

helps hold the molecules

together, and therefore

requiring a high amount of

energy to break the bonds to

turn to a vapor.

2. The less hydrogen bonding the easier to vaporize.

3. The reverse of evaporation is condensation, when water condenses it releases energy (heat).

4. Temperatures in the tropics would be much higher if water didn’t absorb heat.

5. Temperatures in the polar

regions would be much lower

if water vapor did not release

heat when condensing out of

the air.

Ice:

• A typical liquid cools, it

contracts slightly. Its density

increases because its volume

decreases. The solid would

sink because its density is

higher than the liquid.

As water cools it acts like a

typical liquid, until it reaches

4 oC then the density begins to

decrease. As Ice forms at 0 oC the volume expands and it

has lower density than the

surrounding water.

AQUEOUS

SOLUTIONS

NOTES 17.3

Solvents and Solutes

1. Chemically pure water never

exists in nature, because

water dissolves so many

substances

(Universal solvent).

2. Water samples containing

dissolved substances are

called aqueous solutions.

• In a solution, the dissolving

medium is the solvent.

• In a solution, the particles

dissolved are the solutes.

3. Solutions are homogeneous

mixtures.

4. Solvents and solutes can be

solids, liquids, and gases.

5. Ionic compounds and Polar

covalent molecules dissolve

most readily in water, but

nonpolar covalent do not.

The Solution Process:

1. As you place a solute in a

solvent the particles begin to

collide with one another. The

solvent attracts the solute

particles until substance is

dissolved.

2. In some ionic compounds, the

solvent can’t break the ionic

bonds and the salt doesn’t

dissolve.

3. Polar solvents dissolve ionic

and polar molecules, nonpolar

solvents dissolve nonpolar

compounds.

Raises boiling point –

Salt in water

Lowers freezing point –

Salt on road

Solute added to solution:

Electrolytes & Nonelectrolytes

1. Compounds that conduct an

electric current in aqueous

solution or molten state are

electrolytes.

- All ionic compounds are

electrolytes.

Compounds that do not

conduct electric current are

nonelectrolytes.

- Most molecular compounds

and compounds of carbon

are nonelectrolytes.

3. Some very polar molecular

compounds are nonelectrolytes

in pure state, but electrolytes in

an aqueous state.

4. You can have strong or weak

electrolytes. Depends on how

well the solute dissolves into

ions.

Water of Hydration

1. Water molecules are an

integral part of crystal

structure; this is called water

of hydration. Also, called a

hydrate.

2. Effloresce the ability to

lose the water hydration.

3. Hygroscropic the ability to

remove water from the air.

a. These are used as drying

agents (desiccants).

Ex. Silica gel

b. Agents that became wet

from solutions from H2O in air

when exposed to air are

deliquescent.

HETEROGENEOUS

AQUEOUS

SYSTEMS

NOTES 17.4

Suspensions

mixtures from which particles

settle out of solution upon

standing.

Differs from solution because

component parts are much

larger.

Colloids

contain particles that are

intermediate in size between

suspensions and solutions.

1. The properties of colloids

differ from both suspensions

and solutions.

2. Colloids are cloudy in

appearance when

concentrated but clear to

almost clear when diluted.

3. Particles do not settle of a

mixture.

4. Colloids exhibit the Tyndall

Effect, the scattering of visible

light.

Emulsions dispersions of

liquids in liquids. An

emulsifying agent is essential

for the formation of an

emulsion.

(Ex. Mayo → vinegar, oil, and

egg)

PROPERTIES OF

SOLUTIONS

NOTES 18.1

Solution Formation

1. Solutions are homogeneous mixtures and can be solids, liquids or gases.

2. Factors that affect how fast a

substance dissolves.

a. agitation

b. temperature

c. surface area the smaller

the particle, the faster it

dissolves.

Solubility

the amount that dissolves in a

given quantity of a solvent at a

given temperature to produce

a saturated solution.

Particles can move from solid

to a solvated state and back to

a solid again.

This is a saturated solution

(contains the maximum

amount of solvent.)

A solution that contains less

solute than a saturated

solution is unsaturated.

2. Two liquids are said to be miscible

if they dissolve in each other.

(ex. Water & ethanol)

3. Liquids that are insoluble in each

other are immiscible

(ex. Oil & Vinegar)

Factors Affecting Solubility

1. Temperature

a. for most solids as

temperature increases

solubility increases.

b. For most gases as

temperature decreases

solubility increases.

2. Pressure

a. Gas solubility increases as

the partial pressure of gas

above the solution increases

(Ex. Carbonated drinks

contain dissolved CO2 in

H2O) and decreases as

pressure decreases.

3. Supersaturated solution

contains more solute than it

should theoretically continue

to hold.

a. Crystallization of the

solution can occur by adding a

small crystal (seed crystal).

CONCENTRATIONS

OF SOLUTIONS

NOTES 18.2

Molarity

1. Concentration is a measure of

the amount of solute that is

dissolved in a given quantity

of solvent.

a. dilute solution low

concentration

b. concentrated high

concentration

2. Molarity (M) the number of

moles of a solute dissolved per

liter of solution.

a. also known as molar

concentration and read as

“molar”.

b. To calculate the molarity of

any solution - calculate the

number of moles in 1 L of the

solution.

Molarity = moles of solute

Liters of solution

3. Making Dilution:

a. by adding solvent to a

solution you can lower its

molarity.

1) moles of solute do not

change.

4. Percent Solutions

a. If both solute and solvent are

liquids, a convenient way to

make a solution is to measure

volumes.

1)If 20 mL of pure alcohol is

diluted with water to a total

volume and 100 mL the final

solution is 20% alcohol by

volume.

Percent volume =

Volume of Solute x 100%

Solution Volume

COLLIGATIVE

PROPERTIES OF

SOLUTIONS

NOTES 18.3

Decrease in Vapor Pressure

1. Properties of solutions differ

from those of the pure solvent.

2. Properties that depend on the

number of particles dissolved

in a given mass of solvent are

called colligative properties.

3. 3 Important colligative

properties of solutions:

a. vapor pressure lowering

b. boiling point elevation

c. freezing point depression

4. A solution that contains a

nonvolatile solute always has

a lower vapor pressure than

the solvent.

(Volatile easily vaporized.)

Boiling Point Elevation:

1. By adding a nonvolatile solute

would increase the boiling

point.

2. Attractive forces occur

between the solvent and

solute therefore you need

more energy to overcome

these forces.

Freezing Point Depression:

1. When a substance freezes the

particles of the solid take on an

orderly pattern. The presence of

a solute disrupts this. Therefore

more energy must be withdrawn.

2. The more solute you add the

lower the freezing point.

(Ex. Salt and water)

Percent by Mass

A solute in solution is the number of

grams of solute dissolved in 100 g of

solution.

Percent by Mass =

__Mass of solute__ x 100%Mass of solute + mass of solvent

Example:

% by mass =

10 g NaOH___ x 100%

10 g NaOH + 90 g H2O

= 10 % NaOH

Molarity (M)

Number of moles of solute in

one liter of solution.

Molarity =

# of moles of solute

# of liters of solution

Example:

If 0.500 moles of NaOH is dissolved

in 1.00 L of solution, what is the

molarity that is produced?

M = 0.500 mole NaOH

1.00 L

= 0.500 M NaOH

Moles of solute =

M1 x V1 = M2 x V2

Molality (m)

Concentration of a solution

expressed in moles of solute per

kilogram of solvent.

Molality = # of moles solute__

Mass of solvent (kg)

Example:

If 8.50 g of ammonia, which is one –

half mole of ammonia, is dissolved in

exactly 1 kg of water, what molality is

produced?

m = 0.500 mole NH3

1 kg H2O

= 0.500 m NH3