Metallic Solids

• Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces

• In metals valence electrons are delocalized throughout the solid

• “Electron-sea model”

Types of Bonding in Crystalline Solids

Vapor Pressure and Boiling PointVapor Pressure and Boiling Point

Boiling point ≡ temperature at which vapor pressure equals atmospheric pressure.

Normal boiling point ≡ temperature at which vapor pressure is 760 torr.

Fig 11.24

Physical Properties Physical Properties of Solutionsof Solutions

Chapter 13Chapter 13

Solution - a homogenous mixture of 2 or more substances

Solute - the substance(s) present in the smaller amount(s)

Solvent - the substance present in the larger amount

Table 13.1

Solutions

The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

How Does a Solution Form?

• As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

The Effect of Intermolecular forces

Fig 13.1 Dissolution of an ionic solid

Fig 13.2 Hydrated Na+ and Cl− ions

• The negative end of the waterdipoles point toward the positive ion

• Positive ends point toward the negative ion

• Result is hydrated ions

Three types of interactions in the solution process:

• solute-solute interaction• solvent-solvent interaction• solvent-solute interaction

Hsoln = H1 + H2 + H3

Fig 13.2

Energy Changes in Solution Formation• The enthalpy change of the overall process depends on

H for each of these steps.

Fig 13.4 Enthalpy changes accompanying solution processes

Solutions

Solution Formation, Spontaneity, and Entropy

Enthalpy is only part of the picture

Increasing the disorder or randomness of a system tends to lower the energy of the system

Entropy ≡ degree of randomness or

disorder in a system

Solutions favored by increase in entropy that accompanies mixing

Fig 13.6

Caveat Emptor!

• Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.

• Dissolution is a physical change — you can get back the original solute by evaporating the solvent

• If you can’t, the substance didn’t dissolve, it reacted.

Fig 13.7

Unsaturated solution - contains less solute than the solvent has the capacity to dissolve at a specific temperature

Supersaturated solution - contains more solute than is present in a saturated solution at a specific temperature

Sodium acetate crystals rapidly form when a seed crystal isadded to a supersaturated solution of sodium acetate.

Saturated solution - contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature

Fig 13.10

“like dissolves like”

Two substances with similar intermolecular forces are likely to be soluble in each other:

• non-polar molecules are soluble in non-polar solvents

CCl4 in C6H6

• polar molecules are soluble in polar solvents

C2H5OH in H2O

• ionic compounds are more soluble in polar solvents

NaCl in H2O or NH3 (l)

Factors Affecting Solubility

Factors Affecting Solubility

Acetone is miscible in water

H2O

C6H14

Factors Affecting Solubility

• Glucose (which has hydrogen bonding) is very soluble in water

• Cyclohexane (which only has dispersion forces) is not

Fig 13.12 Structure and solubility

Pressure Effect on Gases in Solution

• Solubility of liquids and solids does not change appreciably with pressure

• Solubility of a gas in a liquid is directly proportional to its pressure

Fig 13.14 Effect of pressure on gas solubility

Henry’s Law

Sg = kPg

where

• Sg ≡ solubility of the gas

• k ≡ the Henry’s Law constant for that gas in that solvent

• Pg ≡ partial pressure of the gas above the liquid

Fig 13.15 Solubility decreasesas pressure decreases

Temperature Effect on Solids and Liquids

• Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature

Fig 13.17 Solubilities of several ion compoundsas a function of temperature

• The opposite is true of gases:

• Carbonated soft drinks are more “bubbly” if stored in the refrigerator

• Warm lakes have less O2 dissolved in them than cool lakes

Temperature Effect on Gases

Fig 13.18 Variation of gas solubility with temperature

Concentration UnitsConcentration UnitsConcentration - the amount of solute present in a

given quantity of solvent or solution.

Percent by Mass

% by mass = x 100%mass of solutemass of solute + mass of solvent

= x 100%mass of solutemass of solution

13.3

Mole Fraction (X)

XA = moles of A

sum of moles of all components

Concentration Units ContinuedConcentration Units Continued

M =moles of solute

liters of solution

Molarity (M)

Molality (m)

m =moles of solute

mass of solvent (kg)

13.3

What is the molality of a 5.86 M ethanol (C2H5OH) solution whose density is 0.927 g/mL?

m =moles of solute

mass of solvent (kg)M =

moles of solute

liters of solution

Assume 1 L of solution:5.86 moles ethanol = 270 g ethanol927 g of solution (1000 mL x 0.927 g/mL)

mass of solvent = mass of solution – mass of solute

= 927 g – 270 g = 657 g = 0.657 kg

m =moles of solute

mass of solvent (kg)=

5.86 moles C2H5OH

0.657 kg solvent= 8.92 m

13.3