Chapter 11 Gases, Liquids and Solids (Sections 1-5)

Comparison of the characteristics of solids, liquids, and gases:

Solids: Maintain shape and have constant volumeDo not expand to fill containerAre virtually incompressibleDo not flowDiffusion within a solid occurs very slowlyIntermolecular attractive forces keep

particles locked together in a dense, orderly array – causing solids to be rigid.

Liquids: Assume the shape of their containerDo not expand to fill containerAre virtually incompressibleFlow readily Diffusion within a liquid occurs slowlyParticles closely packed; randomly orientedIntermolecular attractive forces between

molecules are strong enough to keep them together, but weak enough to allow the particles to flow around each other.

Gases Assume shape and volume of their container

Expand to fill containerAre compressibleFlow readilyDiffusion within a gas is rapidMolecules are widely separated and in

constant, random motion – resulting in us being able to ignore interactions between gas molecules.

The state of a substance depends largely on the balance between

Kinetic energy of the particles (which act to keep particles apart from each other)

Attractive forces between the particles (which act to keep particles close to one another)

Both decreasing temperature and increasing pressure allows particles to form attractions to each other and condense to a liquid or settle into a solid state.

INTERMOLECULAR FORCES – Forces that exist between

molecules (substances composed of nonmetals).

* they act to hold neighboring molecules together* weak compared to ionic, metallic, or covalent

bondswhich are intramolecular forces.

It is the intermolecular forces between molecules that loosen or break to allow a change in state

the strength of intermolecular forces determines

the state of a molecular substance at room temp (RT)- weak = gas- moderate = liquid- strong = solid

* the stronger the intermolecular forces, the higher the melting and boiling points

Three types of intermolecular forces are important in neutral molecules:

These are all electrostatic forces – due to attraction between oppositely charged particles due to shifting of electrons

1. Dispersion forces : forces caused by the attraction of

temporary, induced dipoles.* are present in all molecular substances* molecules that are not permanent dipoles

may have dipole moments when their electron clouds shift causing the molecule

to become a temporary dipole* the momentary dipole causes the neighboring

molecule’s electrons to shift, causing it to become a temporary , induced dipole, and causing a momentary attractive force

between the temporary dipoles.

* these are weaker than dipole-dipole forces, but important

Strength depends on the polarizability of the molecule, which is the ease that the electron cloud can be distorted (its squishiness). - In general polarizability increases as the

number of electrons in the molecule increase, so…dispersion forces increase with increasing atomic mass or molecular size and shape.

- Larger molecules will have larger dispersion forces , leading to higher melting, boiling, and sublimation points.

- Longer, more linear molecules have more intermolecular contact and greater dispersion forces than more spherical molecules.

2. Dipole – dipole forces : attractions between oppositely charged ends of close neighboring permanent dipoles (polar molecules)

* moderately strong forces For molecules of approx. equal mass and size,

the strength of dipole-dipole bonds increases with increasing polarity.

(Dipole-Dipole and Dispersion forces are collectively termed

van der Waals forces)

3. Hydrogen bonding : strong dipole-dipole forces between binary (2 atom) molecules made up of H bonded to strong electronegative atoms (F, N, O)Ex: Water

* stronger than regular dipole-dipole forces - b/c the H does not have any other e- around it,

and becomes a definite positive charge on the molecule when it’s electron is pulled

towards the other atom AND because H is so small, it can get close to the neighboring electronegative atom. (decreased distance between atoms = stronger attraction/bond)

* the Hydrogen bond is the attraction between the H on one molecule and the non-bonding electron pairs on a nearby small electronegative atom ( F, O, or N) on a neighboring molecule

* Hydrogen bonding causes compounds such as NH3, HF,

and H2O to have unusually high boiling points.

Some important consequences of H bonding:a. Stabilize proteins 3D shapeb. Allows the means by which DNA can carry genetic

info (joins the two strands of DNA)c. Cause ice to be less dense than water because as

the KE of the molecules decrease, the H bonds form a rigid more open pattern than in water

Ion-Dipole Forces : Form between an ion and a polar molecule

Strength of these increase with increasing ionic charge or increasing dipole moment.

The strength of these forces is what causes ionic compounds to dissolve in polar solvents.

Relative strengths of intermolecular forces:1. Generally

H bonding > dipole-dipole forces > dispersion forces

2. When the molecules of two substances have similar molecular weights and shapes, dispersion forces will be approx equal.

3. Increasing polarity causes increased strength of dipole-dipole forces.

4. When molecules are very different in molecular weight, dispersion forces will determine which has the strongest intermolecular attractions, (the larger molecule will have the larger dispersion force.

5. Effects of all the intermolecular forces present are additive.

6. Attractive forces between ionic compounds are stronger than between covalent compounds.

Practice problems:

#1: In which of these substances is H bonding likely to play an important role in determining the physical properties?

CH4 – no H bonds (H-F; H-O, H-N)H2NNH2 - this molecule contains H bondsCH3F – no H-F bonds (the F is bonded to C)H2S – no H bonds

#2 List the substances BaCl2, H2, CO, HF, and Ne in order of increasing boiling point. (KEY – 1. Ionic bonds > intermolecular forces

2. the stronger the intermolecular forces, the

higher the bp) BaCl2 is ionic and will have the highest bp. HF will have H bonding , will have highest bp of

the molecular compounds.

CO is polar, so it will have higher bp than hydrogen (non polar) and neon

Between H2 and Ne, both non-polar, we must consider dispersion forces. Ne has a greater molecular weight, so it’s dispersion forces will be greater and it’s bp will be higherLowest Highest bpH2 < Ne < CO < HF < BaCl2

Group problems:#3 In which of the following molecules is significant H bonding possible?CH2Cl2 PH3 HOOH CH3COCH3

#4 Identify the intermolecular forces present in the following molecules and select the substance with the highest boiling point.CH3CH3 CH3OH CH3CH2OH

Properties of Liquids

Viscosity: the resistance of a liquid to flow The greater the viscosity, the more slowly it

flows It is related to the ease of molecules in the

liquid moving around in relation to each other.- Depends on intermolecular forces- Shape of the molecule (long molecules

tangle)- Generally, for a series of related

compounds, viscosity increases with molecular weight.

- Viscosity decreases with increasing temp

Surface Tension: a characteristic of liquids in which their surfaces behave like a thin, elastic film.

Due to an imbalance of molecular forces at the surface of the liquid- Molecules in the interior are attracted in all

directions and experience no net force- Molecules at the surface are only attracted

inward, which tends to pull the surface molecules inward, causing them to pack more closely together along the surface.

* Solutes that disrupt and lessen intermolecular forces will decrease the surface tension.This includes non-polar substances like detergents and surfactants.

Cohesive forces : intermolecular forces that bind similar molecules together- (dispersion, dipole-dipole, H bonding)

Adhesive forces : intermolecular forces that bind a substance to a surface- meniscus (curved surface). A U shaped

meniscus is formed when water is in a glass tube b/c the adhesive forces between the water and glass are greater than the cohesive H bonding between water molecules.

- An arch (or inverted U) is formed when the cohesive forces are greater than the adhesive forces, such as in Hg in a glass cylinder.

Capillary action : the rise of liquids up very narrow tubes is due to a combination of

adhesion of the liquid to the walls and cohesion of the molecules to each other.

Phase Changes: changes of state. See chart below identifying the phase changes

and the energy changes associated with each.

1. Each phases change is accompanied by a change of energy in the system. Heat of fusion (enthalpy of fusion, ∆Hfus) is the

heat absorbed by a solid melting to form a liquid.(Conversely it is the heat given off when the liquid solidifies, aka freezes)

2. Heat of vaporization, ∆Hvap is the heat absorbed in the change of liquid to gas.* Heats of vaporization tend to be greater than heats of fusion b/c the intermolecular forces must be broken instead of just loosened.

3. Heat of sublimation, ∆Hsub, is the energy associated with the change of a solid directly to a gas. It is equal to the sum of ∆Hfus + ∆Hvap

4. Applications to real life….- Use of ice cubes to cool drinks- Evaporation of sweat to cool our bodies- Absorption of heat by the refrigerant in a refrigerator

to keep our food cold.

5. Heat curves : Graphs that show the temperature and phase changes that occur as a substance is

heated.* the temp of a substance will remain

constant

until all of it has changed state* the temp of the substance will only

change once all of it has changed state.

* You can calculate the ∆H for each section of the heat curve by using the

equations ∆H = mC∆T for the parts where the temp changes (J)

ΔH = (ΔHfus * moles) for the solid ↔ liquid change (kJ) ΔH = (ΔHvap * moles) for the liquid ↔ gas change (kJ)

Interpreting a heat curve:

Recall that specific heat is the amount of heat needed to raise the temperature of a substance 1°C. The greater the specific heat, the more heat we must add to increase the temperature.

In the horizontal sections indicated by BC and DE, the added heat is used to overcome intermolecular forces instead of increasing their kinetic energy – the substance is changing from one phase to another, but at a constant temperature.

In the sections AB and CD, the heat is mainly being used to increase the KE of the particles.AB the temperature of the ice is increasingBC the ice is changing to water at 0°CCD the temperature of the water is increasingDE the water is changing to gas at 100°CEF the steam (gas) is increasing temp.

Practice problem: Calculate the enthalpy change upon converting 1.00 mol of ice at -25°C to steam at 125°C under a constant pressure of 1 atm. Given…

Specific heat, C, of ice = 2.03 J/gKSpecific heat of liquid water = 4.18 j/gK Specific heat of steam is 1.84 j/gK∆Hfus

= 6.01 kJ/mol∆Hvap = 40.67 kJ/mol

To solve, calculate the ∆H of each section of the heat curve and add them together:

1. Calculate mass of 1 mol water = 18 g2. Ice from -25 to 0°C ΔT = Tf – Ti = 0 - -25 = 25

∆H = mC∆T = (18)(2.03)(25) = 914J 0.914 kJ

Ice to water at 0°C. There is no temp change here, so

we use the enthalpy of fusion value here and multiply it by the moles substance:

∆Hfus = 6.01 kJ/mol * 1mol = 6.01 kJ

(this value is positive because the heat is being input)

Water from 0° to 100°C∆H = (18)(4.18)(100) = 7520 J 7.52 kJ

Water to steam at 100° using ∆Hvap 40.67 kJ/mol * 1mol = 40.67 kJ

(this value is positive because heat is input)

Steam from 100° to 125°C ∆H = (18)(1.84)(25) = 830 J 0.83 kJ

Sum all heat values = 56.0 kJ

Group practice:What is the enthalpy change during the process in which 100g water at 50°C is cooled to ice at -30°C? Use specific heats and enthalpy values from practice problem.

Supercooling: when a liquid is temporarily cooled to below

it’s freezing point without it forming a solid. This occurs when conditions do not allow the

particles time to arrange themselves in ordered structure of the solid.

A supercooled liquid is very unstable and will solidify when triggered

Vapor Pressure: The pressure exerted by vapor particlesin the space above a liquid in a closed container when the liquid and vapor are in dynamic equilibrium.Since at any time liquid molecules are evaporating and escaping the surface to become vapor at the same rate that vapor molecules are colliding with the surface and becoming liquid.

As temperature increases, vapor pressure increases.

* when non-volatile solutes are added to the liquid, they displace some of the solvent particles on the surface, and thus get in the way of and slow down evaporation. This causes a reduced vapor pressure.

Volatility

1. When vaporization occurs in an open container, the vapor moves off, and eventually the liquid will completely evaporate

2. Substances with high vapor pressure evaporate more quickly than substances with low vapor pressure. (hot water will evaporate faster than cold water)

3. Liquids that evaporate easily are said to be volatile.

4. Volatility is affected by temperature and intermolecular forces. Increasing temp will increase volatility Weaker intermolecular forces allow molecules

to escape the liquid more easily and will cause a substance to have a higher vapor pressure.

Vapor pressure and boiling point:A liquid boils when its vapor pressure = external pressure.

Boiling point will increase with increasing pressure

Normal boiling point – the temp at which a substance will boil at 1 atm.

Practical applications…Pressure cooker – increases pressure, causing water inside to boil at a higher temp – increasing speed of cooking

Cooking at higher elevations where the pressure is less than 1 atm, causes water to boil at temperatures less than 100°C. At a lower temperature food takes longer to cook.

Practice problem: Use the graph above to estimate the boiling point of diethyl ether at 0.80 atm. (Hint- how are boiling point and vapor pressure related?)

Clausius-Chapeyron equation: the natural log (ln P) of the vapor pressure of a liquid is inversely related to its temperature.Meaning a graph of the log of vapor pressure vs 1/T will be linear. (you will use this for the vapor pressure lab).

Phase Diagrams: a graphic that summarizes the conditions under which equilibria exist between the different states of matter at any given T and P

Ex. The temp at which a solid and a liquid exist in equilibrium is the melting point of the solid and the freezing point of the liquid.

The liquid-vapor interface (line) starts at the T (triple point), at which all three states are in equilibrium, and ends at the critical point (C), above this T and P, the liquid and vapor are indistinguishable and exist as a supercritical fluid. Each point along this liquid-vapor line is the boiling point of the substance at that pressure.

The interface between the liquid and the solid marks the melting point at each pressure.

A substance cannot exist in the liquid form below the triple point.

The sublimation point at each pressure is along the solid-gas interface line.

Phase Diagram of Water

The Critical point of water occurs at an unusually high temp and pressure due to the strong H bonding between water molecules.

The melting curve (solid-liquid interface) is slightly sloped backwards, meaning that with increasing pressure, the melting point decreases for water. Why? Because water is very unusual in that its liquid form is more dense that its solid form.

At 1 atm, water can exist at any phase and is determined by the temperature (as we know from our common weather)

Phase diagram of CO2

Note that CO2 cannot exist as a liquid at normal pressure (1atm)

Under normal pressures, solid CO2 sublimates.

Practice Problem: Use the phase diagram for methane, below to answer the following questions:

What are the approximate T and P of the critical point?

What are the approximate T and P of the triple point?

Is methane a solid, liquid, or gas at 1 atm and 0oC?

If solid methane at 1 atm is heated while pressure is held constant, will it melt or sublimate?

If methane at 1 atm and 0oC is compressed until a phase change occurs, in which state is the methane when the compression is complete?