Liquids & Solids

• Kinetic Molecular Theory explains liquids and solids as well as gas.

Figure 10.1: Schematic representations of the three states

of matter.

What compound is in all three phases RIGHT NOW ??

WATER

So why aren’t all the other gases in the air - like CO2 or O2 or N2 in all three phases ??

Liquid N2 is –196 oCDry Ice is –78 oC

Liquid O2 is –183 oC

Condensed Phase

PROPERTY GAS LIQUID SOLID

shape Assumes ____________ ___________

Assumes ____ __________ -flows

____________ shape

Kinetic Energy

Density High

Order Moderate

Compressibility Low

Importance of Intermolecular forces

Thermal expansion

Small

Condensed Phase

PROPERTY GAS LIQUID SOLID

shape Assumes volume of container

Assumes shape of container - flows

Retains own shape

Kinetic Energy High (a lot of movement)

Medium Low (vibration only motion)

Density Low High High

Order None Moderate High

Compressibility High Low Low

Importance of Intermolecular forces

Low Medium High

Thermal expansion

Large Small Small

What holds things together?Intramolecular Forces

• Hold particles together _________ molecule (covalent bonds).

Intermolecular Forces

• Cause attractive interactions _______ __________ (molecules, ions or atoms) of a substance.

WITHIN

BETWEEN

Particles

Intermolecular ForcesIntermolecular forces hold solids

and liquids together! There are four types:

1.Ion-dipole forces

2.Dipole-dipole forces

3.(London) Dispersion forces

4.Hydrogen bonding

ElectronegativityElectronegativity relationships

between atoms ultimately determine the nature of bonding.

Electronegativity

A=B A<B A<<B

Electron “Share”

A : B A+ : B- A : B

Type of Interaction

_____ sharing

______ sharing transfer

Bonds formed

Elements Involved

Non PolarCovalent

PolarCovalent

Equal Unequal

Ionic

Likenonmetals

Unlikenonmetals

Metal &nonmetal

Complete

Bond PolarityBecause of a difference in

electronegativity, polar covalent bonds have shift in electron density There is a permanent unequal distribution of electrons due to different electron pulling power.

H - F

H - F

+

-H is partially (+) because it has lost some e- density

F is partially (-) because it has gained some e- density

Molecular PolarityElectronegativity differences

between atoms combine with molecular shape to determine if the whole molecule will be polar. Molecules will be polar if:

1.They contain polar bonds

2.The polar bonds are asymmetrically arranged around the central atom

Molecular PolarityExamples:

polar bonds

polar bonds

polar bonds

polar bonds

geometry geometry geometry geometry

molecule molecule molecule moleculeC

Cl C Cl

Cl

Cl

Cl C H

Cl

Cl

H

H

O O C O

asym sym asym sym

Polar nonpolar Polar nonpolar

Ion-Dipole ForcesIon-dipole Forces are the electrostatic

attraction between an ion and the oppositely charged end of a polar molecule.

Example:

Ions in water

 

Interaction _______ energy = (enthalpy) of hydration = Hhyd

Cl-

OHH

O H

H

O

H

H

OHH

O H

H

OHH

K+

releases

Dipole-Dipole ForcesDipole-Dipole Forces are the

electrostatic attraction between oppositely charged ends of polar molecules

Example: HCl (l) Water (l)

HCl HCl

ClH ClH OHH

OHH

OHH

+- +-

+-

+ - + -

PolarizabilityPolarizability is the tendency for the

electron density of atom or molecule to be deformed or ”sloshed around”

Polarizability will increase with:

• _________ mass

• _____________ of electrons

• electrons _____________

increased

greater number

less tightly held

PolarizabilityPolarization of molecules/atoms results

in “____________” dipoles.

Instantaneous dipole moments are due to the polarizability of a molecule!

Example: I2 iodine has no permanent dipole (non-polar)

instantaneous

London Dispersion ForcesDispersion Forces are the electrostatic

attraction involving __________ dipole moments in (non-polar) molecules 

The greater the polarizability of the atoms/molecules, the ________ the interaction

• ___ molecules are polarizable

• ___ molecules experience dispersion forces (even polar molecules)

instantaneous

stronger

All

All

Which would be affected more by London Dispersion Forces?

H2 or I2

How many electrons does each have?

Hydrogen BondingHydrogen Bonding is an interaction

that exists _______ the H-atom in a polar covalent bond with ________ and the _______ electrons on a nearby highly electronegative atom (________).

Special case of dipole-dipole interaction

Since H _____ shares all its e-, the positive nucleus is relatively exposed and is attracted to unshared electron pairs on another molecule

betweenF, O, or N

lone-pairF, O, or N

always

Hydrogen BondingHF H2O NH3

H-bonds are much ______than covalent or ionic bonds.

H-bonds are ________ than other dipole-dipole or dispersion (London) forces.

H-bonds are strong enough to determine structure of _____________!

H F

H F H F

OHH

OHH

OHH

OHH

NH H H

NH H H

NH H H

weaker

stronger

ice and proteins

Cool feature of water

• In the solid crystal form, the solid water (ice) is __________ than liquid water. This is NOT true for other solids.

• This allows– -– -

Less dense

Ice to float

Fish to live at bottom of lake

Due to arrangement and hole in middle

http://www.its.caltech.edu/~atomic/snowcrystals/ice/iceIh.gif

IM Force PracticeWhat are the predominant forces in the

following interactions?

• Dissolving candle wax with kerosene

• Ethyl alcohol absorbing into your blood stream

• Iodine gas molecules condensing into solid iodine

-> dispersion forces (both are nonpolar)

-> dipole-dipole

-> dispersion forces (both nonpolar)

IM Force Practice• Detergent removing gravy from a

garment, in water

• Dissolving table salt in water

• Dissolving sugar in water

-> dispersion forces (nonpolar end of detergent and fat)

ion-dipole (detergent and water)

-> ion-dipole

-> dipole-dipole

IM Force PraciticeWhich has stronger intermolecular

forces?

•  Dissolving table salt vs. sugar in water

•  Dissolving I2 vs. O2 in water

-> table salt; ion-dipole stronger than dipole-dipole.

-> I2; both are dispersion forces, but I2 is more polarizable giving it a larger instantaneous dipole.

IMF and Solid/Liquid Physical Properties

Intermolecular Forces (IMF) between particles determine physical properties:

• Melting point (mp)

• Boiling point (bp)

• Viscosity

• Surface tension

• Vapor pressure

Liquid Properties

1. high density compared to gas due to ___________ of molecules

2. Incompressibility

3. Ability to ______ - due to fluidity

compactness

diffuse

Physical PropertiesViscosity is the resistance of a liquid to

flow (viscosity is high if molecules _______move past each other easily).

• The _______ the IMF, ______ the viscosity of the liquid.

• At _______, viscosity _________; particles have ________ and can ________ attractive forces (note: viscosity is also affected by molecular chain lengths, physical entanglement)

Do NOT

stronger higher

Higher T decreasesmore KE

overcome

Liquid Properties

4.Surface tension - a force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the ____________________. Varies from liquid to liquid.

5. Capillary action-attraction of the surface of a _____ to the surface of a _____

smallest possible size

liquidsolid

Physical PropertiesSurface Tension is the energy required

to _________________ of a liquid.

For a given number of liquid molecules:

• Highest surface area is to put them all in a row.

• Lowest surface area is a sphere

 ______ IMF results in liquids that have _______ surface tensions.

increase surface area

Strongerhigher

Physical PropertiesCohesive forces are a measure of how

well molecules _______________.

Adhesive forces are a measure of how well molecules stick to ___________ (e.g., to the molecules of a container). 

If adhesive force > cohesive force, then the molecules will maximize surface area to maximize adhesive interactions!

-this determines _________________!

stick to each other

other molecules

direction of meniscus

Physical Properties ____ ____

 

Capillary action is the result of _______________.

Paper chromatography is a separation technique based on __________ in adhesive forces.

meniscus

Ad>Co Co>Ad

adhesive forces

differences

Nonpolar liquid mercury forms a convex meniscus in a glass tube, whereas polar water

forms a concave meniscus.

http://www.wwnorton.com/chemistry/overview/ch9.htm#phase_diagram

Liquid Properties Evaporation and Boiling

6.Evaporation - process by which particles escape from the surface of a _________ liquid and enter the gas state.

Boiling - is a change of a liquid to ________ of vapor.

non-boiling

bubbles

Evaporation

• Escape of molecules from a liquid to a gas _________________________________

• Warmer molecules have _________________ & they are able to ____________________ This leaves the ______________ molecules behind.

Without boilingHigh Kinetic Energy (KE)

Escape the surface

cooler

Evaporation

• This is _________________ in nature

– _____________________________

– ______________________________

crucial

Sweat COOLS

Water cycle for RAIN

Liquid Properties

7.     freezing or solidification - change of liquid to solid by _______________

removal of heat

Physical PropertiesMelting Point/Boiling Point:

• The _______ the IMF, ______ the melting point of the solid.

• The _______ the IMF, ______ the boiling point of the liquid.

 

Stronger IMF means that _________ is required to disrupt interactions!

stronger higher

stronger higher

more heat

Physical PropertiesVapor pressure is the pressure of the

vapor of a substance when the liquid and vapor phases are in _______________.

If a liquid is in a closed container, some molecules have enough KE to escape the IMF and form vapor phase in container.

 

 

____ container will _____ reach equilibrium

liquid liquid

dynamic equilibrium

open never

Physical PropertiesSubstance reaches dynamic

equilibrium when there is __________ in particles in gas phase (even though gas and liquid particles in constant flux!)

Vapor pressure _______ w/ _________

________ T ________ # of molecules with sufficient energy to escape surface of liquid.

Liquid with high vapor pressure is ______!

no net change

increases Increases T

Increasing increases

volatile

Volatile liquids

• Liquids ____________ easily

• Examples:

• Have _______________ attractive forces _______________ molecules

• Would water be volatile?

• Would CH4 be volatile?

evaporate

weakBETWEEN

No – lots of attraction, H bonding

Yes, non-polar, no attraction

Copyright © Houghton Mifflin Company. All Rights Reserved. 47

Behavior of a Liquid in a Closed Container

Just closed containerInitially only moleculesLeave surface

Container atEquilibrium

BOTH leavingAND returning

Equilibrium Vapor Pressure

Vapor Pressure

Physical PropertiesBoiling is the conversion of liquid to

vapor within a liquid as well as its surface. Occurs when ____________________ = ________________

Boiling point is the temperature where equilibrium vapor pressure of the liquid is _______ the ambient atmospheric pressure. (i.e., equal to ____________ __________________________). 

equilibrium vapor pressureatmospheric pressure

exactly pressure acting on the surface by the atmosphere

Physical Properties

Example: At sea level, temp at which vapor pressure is 1 atm, the bp = 100oC for water

At this point, bubbles of vapor can form ______ the bulk of the liquid

molecules of gas not just coming from the surface but _____________

within

EVERYWHERE

Physical PropertiesNormal boiling point refers to standard

pressure, 1 atm.

If Patm , then vp must be _ for boiling to be reached (requires __ T).

If Patm , then vp must be _ for boiling to be reached (requires __ T).

Since bp occurs at lower T, takes ______ to cook things at high alt (low P)

longer

Boiling Point changes due to pressure

• Lots Little

Kinetic Molecular TheoryGases consist of particles that are very

far apart.

• If KE > IMF, substance ___________

• If KE < IMF, substance is __________

Molecules of condensed phase are close together. Energy (heat) must be added to _________________; energy must be released when ____________ ________ in liquid or solid.

must be a gas

liquid or solid

separate the molecules particles come together

Heating/Cooling Curve

When temperature is changing,Energy is increasing orDecreasing Kinetic Energy (KE)

When temperature is NOT changing,Energy is used to overcome attractive forces like H-bonding or dipole-dipole

Phase ChangeReactions often involve changes in

phases of substances which also involve energy or enthalpy changes.

The enthalpy terms associated with common phase changes include:Type of phase change Symbol

Heat of fusion

Heat of vaporization

Hfus

Hvap

Molar Heat of Fusion• Amount of heat energy _______ to

melt 1 mole of solid at its melting point. Must ________ attractive forces. (solid liquid)

• Or the heat energy ________ when one mole freezes. (liquid solid)

• ____Temperature change during phase change.

• Temperature this occurs varies for every solid

required

overcome

released

NO

Molar Heat of Vaporization• Amount of heat energy _______

to vaporize 1 mole of liquid at its boiling point. (liquid gas)

• Varies for every liquid. The ___________ the intermolecular attraction, the higher the energy needed to separate, the higher the temperature.

• Or, amount of heat ________ to condense 1 mole of gas (gas liquid)

• ____Temperature change

needed

STRONGER

released

NO

Phase ChangeH’s for opposite changes will have

the same mag., but negative sign. The phase transitions of water:endo

phase change

Symbol exo phase change

Melting+6.02 kJ/mol

HfusFreezing

Vaporization+40.8 kJ/mol

HvapCondensation

-6.02 kJ/mol

-40.8 kJ/mol

Heating CurveEach segment of the curve has an

enthalpy change (H).

The total H is sum of each segment.

For T segments, H = Csp x mass x T

For phase change segments (T = 0)

H = Hvap x n

H = Hfus x n n = # moles or mass in grams

Energy change with grams for water

• Solid Liquid (freezing or melting): – Given g H2O x 335 Joules/gram

• Heating or cooling liquid: – Given g H2O x 4.18 J/g oC x T oC

• Liquid Gas (boiling or condensing): – Given g H2O x 2260 J/g

• Heating ice or heating steam– Given g H2O x 2 J/g oC x T oC

T = change in temperature

Energy Change TablesSolid Liquid (freezing or melting)

g H2O 335 J

1 g

Heating or cooling liquid

g H20 4.184 J T oC

1 g oC

Liquid Gas (boiling or condensing):

g H2O 2260 J

1 g

Heating ice or heating steam

g H20 2.09 J T oC

1 g oC

Heating/Cooling CurvePlot change in T of a substance as a

function of heat energy added:

Energy is _________ _______when a substance changes state. T is _______ during phase change!

absorbed or released

constant

Heating Curve PracticeWhat is the H for the process in which

2.00 g water at 27oC is converted to ice at –3oC?

Heating Curve PracticeWhat is the H for the process in which

2.00 g water at 27oC is converted to ice at –3oC?

J 905

J 47905.

J 412.5kJ

J 1000kJ 00.667J 93225.

C032.00gCg

J2.09

18.02g

molg 2.00

mol

kJ6.01C270g 2.00

Cg

J4.184

ΔT)mass(CMW

1massΔHΔTmassC

C)3- toC0 from ice ΔH(cooling water)gΔH(freezinC)0 toC27 from water ΔH(coolingΔH

oo

oo

icespfusH2Osp

ooootot

How a Wall of Water works

• Heat is _____________________ to boil

• Heat is _____________________ to freeze

Added / absorbed

Released

Phase DiagramsPlot of P vs. T, for phases:

Triple point – all phases (s, l, g,) are in equilibrium

Critical point - critical T & P to a substance where it cannot exist as a liquid above the critical temperature

For waterCritical temperature = 373.99 oCCritical pressure = 217.75 atm

http://www.wwnorton.com/chemistry/overview/ch9.htm#phase_diagram

Structure and Properties of Solids

Solids can be divided into two categories:

• Amorphous

• Crystalline

• Ionic crystals

• Metallic crystals

• Molecular crystals

• Network crytals

SolidsAmorphous solid: particles have

___ orderly structure (IMF vary in strength throughout structure)

Example: ____

Crystalline solid: atoms, ions, or molecules ordered in __________ arrangements (causes crystals to have regular shapes)

Example: ______________________

glass

NaCl, ice, diamond, all metals

NO

well-defined

Crystalline SolidsIonic Solids: ions are held together by

ionic (electrostatic) interactions;

 

bond strength ________ with ion size, and _______ with ion charge!

• ionic bonds strong

• high melting point, boiling point

• low vapor pressure

Like a magnet

decreasesincrease

Crystalline SolidsMetallic Solids: all atoms present are

metals (inc. hydrogen at very low T)

Electrons in valence shells shared by many atoms, called __________electrons.

*____ electrical conductivity (mobile e-)

*____ thermal conductivity (tightly packed identical atoms)

*malleable and ductile

delocalized

good

good

Crystalline SolidsMolecular Solids: have ______

intramolecular bonds (covalent) and _______ intermolecular bonds, like H-bonds.

Example: sugar, water

• soft substances

• ___ melting and boiling points

• _____ vapor pressures

low

higher

strong

weaker

Crystalline SolidsNetwork Solids: are single giant

molecules with an endless number of covalent bonds among atoms. All bonds are _____________.

Example: diamond, graphite, sulfur

• can be very hard (diamond)

• tend toward high melting, boiling

• low vapor pressure

equally strong

Crystalline SolidsIn summary:crystalline solid type

Example Structural Unit

Forces Between Units

Typical Properties

Ionic NaCl ions ionic: (+) and (-)

hard, brittle, high mp, poor conductorwater soluble

Metallic Cu(s) atoms Metal atoms with delocalized e-

malleable, ductile, luster, conductive

Molecular sugar molecules dispersion forces + other intermolecular forces (depends on type of molecule)

low to moderate mpsoft, poor conductor

Network graphite covalently bonded atoms

covalentbonds

wide range of mp, poor conductor (some exceptions)