CHAPTER 9: LIQUIDS AND SOLIDS

Section 9.1 Liquid/Vapor Equilibrium Vaporization – process in which a

liquid vapor open container

- evaporation continues until all liquid evaporates

closed container

1) Liquid evaporate. 2) Vapor particles collect and

Condense. 3) Eventually,

Rate of Evaporation

Rate of Condensation

DYNAMIC EQUILIBRIUM

Vapor Pressure ● At equilibrium, # molecules/volume is constant. ● Pressure of gas over liquid is constant. ● As long as both liquid and vapor are present,

the pressure exerted bythe vapor is independent of the volume of the container.

Vapor Pressure is dependent on:

a) Characteristics of liquid b) Temperature

Important As long as BOTH liquid and solid are present, the vapor pressure will be constant.

If volume

More liquid will evaporate

Equilibrium Re-establish

Vapor Pressure vs. Temperature

In general, Vapor pressure as Temp. For example, H2O Vapor Pressure of H2O Temp 24 mmHg 25°C 92 mmHg 50°C 760 mmHg 100°C

Higher the Temperature

More Molecules Vaporize

A plot of pressure vs. temperature does not produce a straight line.

This is not a direct relationship!

● Graphs of curves are often difficult to

interpret.

The solution to this problem: Graph manipulated variables of pressure and temperature that will represent a

straight line.

Instead of P vs. T, graph ln P vs. 1/T

● Recall that the general equation of a straight line is y = mx + b (m = slope & b = y-intercept) Here, y = 1n P x = 1/T m = -∆Hvap/R Therefore,

ln P = - TR

H vap 1 + b

If 2 different temps are evaluated:

at T2 : ln P2 = -2

1

TR

Hvap

+ b

at T1 : ln P1 = -1

1

TR

Hvap

+ b

Clausius- Clapeyron Equation

121

2

12

11lnlnln

TTR

H

P

PPP

vap

where R = 8.314 J/K●mole

Boiling Point

Vapor Pressure of Liquid

EQUALS

Pressure Above the Surface of the Liquid

Normal Boiling Point = the temperature a liquid boiling at 1 atm of pressure above the liquid.

● The boiling point of any liquid can be

lowered by reducing pressure above liquid. Varies with altitude.

Critical Temperature – the

temperature above which the liquid state of a pure substance cannot exists regardless of the pressure.

Critical Pressure – the pressure that be applied to cause the condensation of a pure liquid at the critical temperature.

Section 9.2 Phase Diagrams ● Phase Diagram – a graphical way to summarize the conditions under which the different states of a substance are stable.

Water’s Phase Diagram

Liquid

Vapor

Solid

Pre

ssu

re (

atm

)

Temperature (oC)

Critical Point

Normal Boiling Pt.

1.00

.0060

217.75

NormalFreezing Point

Triple Point

0.00 0.01 100.00 373.99

*Not to Scale

D C

B

A

● Melting-Point Curve - Observe the solid/liquid states at different

pressures. - Along the curve both phases are in

equilibrium.

- Special Note: When conditions indicate that a substance is in the liquid or solid state, the vapor of that substance is also present (in equilibrium).

- The question one should ask is “How much

vapor is present?”

- Answer: It depends!!! Sublimation – Transformation of a solid

directly into a vapor. Melting Point – The opposite process

freezing.

Triple Point – the point on a phase diagram representing the temperature and pressure at which three phases of a substance coexist in equilibrium.

For example, water’s triple point is

0.1°C, 0.00603 atm and all phases coexist.

Section 9.3 Molecular Substances: Intermolecular Interactions

1. Nonconductors (when pure) examples: I2, C3H8, C2H5OH

● Most water solutions are also nonconductors ● Some polar molecules form ions when they react with H2O

conduct electricity

For example: HF(g) H+1 (aq) + F-1(aq)

2. Generally, molecular compounds are insoluble in water. 3. They have low melting & boiling points. ● Many are gases (N2, O2, …) ● Some are liquids with melting points <25oC (like H2O, mp = 0oC).

● Some are solids with melting points <300oC (like I2, mp = 114oC).

The boiling point and melting point of molecular substances is

directly related to the strength of their INTERMOLECULAR ATTRACTIVE FORCES

among molecules.

Intermolecular Forces 1. (London) Dispersion Forces ● Found in all molecular substances.

involves a temporary or induced dipole.

Consider the H2 molecule

● Nonpolar bond equal sharing of electrons

For an instant, the electrons within the molecule can concentrate closer to one atom in the molecule.

● Produces a +/ - (dipole) within the

molecule. ● This temporary dipole induces a similar dipole

within another molecule. ● These temporary dipoles result in the two molecules attracting each other.

This attraction is the Dispersion Force! The attraction is dependant on:

1. The # of electrons in the molecules involved.

2. The ease of the electrons in the molecules to be dispersed within the individual molecules.

Larger atoms/molecules

Produce Greater Temporary Dipoles.

In general,

As Molar Mass ,

The dispersion forces ,

The bp & mp of nonpolar

molecules .

2) Dipole Forces ● These interactions occur in polar

molecules. ● The +/ - (dipole) of one polar molecule

lines up with +/ - (dipole) of another polar

molecule (opposites attract). The greater the dipole moments (the

measure of the polarity of a molecule) of the molecules, the stronger the attractive force. ● This interaction (attraction) really only works when the molecules are close together. ● When the molecules are in the gas phase,

the dipole forces of attraction are negligible (as is the case for dispersion forces).

3) Hydrogen Bonding

This attraction occurs in polar molecules, HOWEVER, only in molecules that have X – H bonds where X = N, O, or F.

This attractive force is an unusually strong dipole force of attraction.

Why is the hydrogen bonding such a powerful attractive force?

2 – Reasons: 1. There is a large difference in the

electronegativity of the X and H H(2.2) F(4.0)

H(2.2) O(3.5) H(2.2) N(3.0)

The H – atom almost behave as a “naked” proton.

2. The H – atom is very small.

The lone pair of electrons on F, O, and N can get really close to H.

Let’s look at the pattern of boiling points.

bp(oC) bp(oC) bp(oC)

NH3 -33 H2O 100 HF 19

PH3 -88 H2S -60 HCl -85

AsH3 -63 H2Se -42 HBr -67

SbH3 -18 H2Te -2 HI -35

Note the effect of hydrogen bonding in the first row of boiling points.

IMPORTANT Although these intermolecular force are very important, they are very weak compared to a covalent bond.

Section 9.4: Network Covalent, Ionic, and Metallic Solids

● Most Molecular substances are gases or

liquids at room temperature. ● Most NON-molecular substances (network

covalent, ionic, and metallic) are solids at room temperature.

1. Network Covalent Solids

These solids are made of atoms joined by a continuous network of covalent bonds.

In general, these solids are: a. High melting (over 1000oC)

- In order for this type of solid to melt, bonds need to be broken

- Remember: When molecular solids to

melt, only interactions need to be broken!

b. This type of solid is typically insoluble in all common solvents.

- Why? Bonds need to be broken!

NOT EASY TO DO!!!

c. These solids are poor conductors of electricity.

- Why? No mobile electrons are available

Example: Carbon 2 types of solids exist: Graphite & Diamond Both have very high melting points >3500oC.

Ionic Solids These solids are held together by very strong electrostatic attractive forces (ionic bonds).

1. They are composed of cations/anions. 2. They are non-volatile (do not become gases

very easily). 3. They are high melting (600 – 2000oC). 4. They do not conduct electricity (they only do

when they form aqueous solutions or they are molten).

5. Many (but not all) are soluble in water.

Metallic Solids A structural unit of electrons and metal cations.

Positive metal ions anchored in position with electrons moving around from one metal ion to another.

Metals are highly conductive - Very mobile electrons – metals have very low

electronegativities.

● Metals have a high thermal conductivity - Very mobile electrons vibrate

● Metals are ductile and mobile.

● Metals have high luster.

- Electrons within a metal can absorb and emit light energy very easily.

● Metals are insoluble in common solvents.