chapter 11 the solid and liquid states. malone and dolter- basic concepts of chemistry 9e2 setting...

Chapter 11Chapter 11

The Solid and Liquid StatesThe Solid and Liquid States

Malone and Dolter- Basic Concepts of Chemistry 9e

2

Setting the Stage – Ice floating on Water

We all know that ice floats in water. This turns out to be very unusual

behavior, since most solids sink in their liquids (the solids are more dense than the liquids).

Life on this planet depends on this odd behavior.


3

Unusual Behavior of Water


4

Setting a Goal – Part AThe Properties of Condensed States and the Forces Involved

You will become familiar with the various forces between molecules and ions and how these forces affect condensed states of matter.


5

Objective for Section 11-1

List the physical properties of solids and liquids that distinguish them from gases.


6

11-1 Properties of the Solid and Liquid State

They have high densities (g/mL). They are essentially incompressible. They undergo little thermal expansion. They have a fixed volume. Solids are rigid and have a definite shape. Liquids flow and do not have a definite

shape.


7

Two Assumptions of the Kinetic Molecular Theory

1. Solid and liquids are composed of particles that have kinetic energy.

2. The average kinetic energy of the particles is related to the temperature.


8

Kinetic Molecular Theory for Liquids and Solids that do not Apply to Gases

3. Particles have a significant attraction for each other and so are held close together.

4. Particles occupy a significant fraction of the volume of the substance.

5. The motion of the particles is not random, but restricted by interaction with other, neighboring particles.


9

Condensed States of Matter

Because of the lack of empty space in solids and liquids, they are referred to as condensed phases.


10

Particle Motion

In solids, particles can move about in place (rotate and vibrate), but they stay in a confined space.

In liquids, particles can move about in three dimensions (translate), as well as rotate and vibrate, but they stay in contact.


11


Describe the types of intermolecular forces that can occur between two molecules and their relative strengths.


12

11-2 Intermolecular Forces and Physical State

Solids have the greater intermolecular forces, followed by liquids. Gases ideally have no intermolecular forces.

Three types of intermolecular forces: London or dispersive forces Dipole-dipole forces Hydrogen bonding forces


13

London or Dispersive Forces

Since molecules are surrounded by negatively charged electrons, we would expect that they would repel each other.

In reality, they are actually attracted to each other.

The most basic force is instantaneous dipole-induced forces.


14

London Forces

Occur in any molecule (including nonpolar)

At any instant, there can be an imbalance in the distribution of positive and negative charge.

For that instant, the molecule is somewhat polar (i.e. it has an instantaneous dipole).

If the molecule in that instant is near another molecule, it will induce a dipole in the other molecule.


15

London Forces

Instantaneous dipoles become more likely for molecules with a greater number of electrons, so London forces are more significant.

Larger molecules are more polarizable because they are surrounded by larger, more diffuse electron clouds.

Hence, London forces increase with molar mass.


16

London Forces


17

Examples of London Forces and Boiling Point Variation

Hydrocarbon

Molar mass/

g/mol

State at 25 oC/1 atm

Boiling point/oC

CH4

Methane

16 Gas -162

C8H18

n-Octane

114 Liquid 126

C18H38

n-octadecane

338 Solid 308

Strengthof

Londonforces


18

Examples of London Forces and Melting Point VariationCompound Molar

mass/g/molMelting point/oC

CH4 16.0 -183

CF4 88.0 -150

CCl4 154 -23.0

CBr4 332 90.0

CI4 520 171

StrengthofLondonforces


19

Dipole-Dipole Attraction Polar molecules

(those with permanent dipoles) can align themselves so that the negative end of one molecule is oriented toward the positive end of another.

Dipole-dipole attractions add to the effect of London forces.


20

Comparison of Forces

For molecules of similar mass (i.e. similar London forces) the polar molecules will have the greater intermolecular forces.

Consider

Substance Molar mass State of matter* CO2 44 g/mol gas

CH3CN 41 g/mol liquid

* At STP


21

Hydrogen Bonding

F, O, and N are small, highly electronegative atoms.

They tend to attract a significant amount of negative charge to themselves and they have lone pairs.

The differences in electronegativity between H and F, O and N are large, but not large enough to yield an ionic bond.


22

Hydrogen Bonding

The bond between H and F, O or N is a highly polar, covalent bond.

For H-F, H-O and H-N The H has lower electron density (+). The F, O, or N has higher electron

density (-). H has no core electrons, so the tiny

bare nucleus is exposed.


23

The Hydrogen Bond

The partially positive H in one molecule is attracted to the lone pairs on F, O or N (which are particularly electron rich), in a different molecule.

The hydrogen bond formed is much stronger than a dipole-dipole interaction, but weaker than a covalent bond.

H F H F _

An intermolecularhydrogen bond

_


24

Effect of Hydrogen Bonding

Hydrogen bonding has a more significant effect on bulk physical properties than dipole-dipole forces.


25

Hydrogen Bonding in Biomolecules

Hydrogen bonding is critical to the properties of biomolecules such as DNA and proteins.

It also plays a part in enzyme reactions (next slide)


26

Hydrogen Bonding in Biomolecules


27


Classify a solid as ionic, molecular, network, or metallic.


28

11-3 The Solid State: Melting Point

Melting point – the temperature at which a crystalline solid melts (turns from a solid to a liquid).

Formally, it is the temperature at which the solid and liquid are in equilibrium (think ice in a glass of water).

The melting point is a definite and constant physical property.


29

Types of Solids

Amorphous solids Have no definite shape Examples are glass, rubber, many

plastics Crystalline solids

Molecules or ions that are arranged in a regular, symmetric structure called a crystal lattice.


30

Ionic Solids

Crystalline solids in which ions are the basic particles making up the crystal lattice.

Forces between the ions are strong. These strong interparticle forces result in

a very high melting point.


31

Examples of Ionic Solids

Lithium iodide LiI m.pt. 447 oC Lithium fluoride LiF m.pt. 867 oC Sodium chloride NaCl m.pt. 801 oC Barium chloride BaCl2 m.pt. 975 oC

Strengthof ionic bonding


32

Molecular Solids

Basic particles of the crystal lattice are individual molecules.

These particles are held together by London forces and in some cases dipole-dipole forces or hydrogen bonding.

Have a wide range of melting points due to the wide range of intermolecular forces possible.


33

Examples of Molecular Solids

Organic solids, such as Benzene m.pt. 5.5 oC

Naphthalene m.pt. 80 oC

Benzoic acid m.pt. 122 oC

HC

HC

CH

CH

CH

HC

HC

HC

CH

C

C

HC

CH

CH

CH

HC

CHHC

HC

HC CH

C C

OH

O


34

Network Solids

Atoms are covalently bonded throughout the entire sample of the solids.

These materials have very high melting points since covalent bonds must be broken (melting points are so high they are hard to establish).

Sometimes called giant molecules.


35

Network Solids: Allotropes of Carbon

Diamond, graphite and Buckminsterfullerene are shown here


36

Metallic Solids

Crystalline solids Positive metal ions are regular positions in the

crystal lattice. Electrons are moving freely among the positive ions.


37

Setting a Goal – Part BThe Liquid State and Changes in State

You will learn about properties of the liquid state and examine how energy is associated with phase changes.


38

Objectives for Section 11-4

List and define physical properties of the liquid state.

Discuss the relationship between intermolecular forces and liquid properties.


39

11-4 The Liquid State: Surface Tension, Viscosity, and Boiling Point

The liquid state has two unique physical properties: surface tension and viscosity.

Surface tension The force that causes the surface of a

liquid to contract. Makes drops of water spherical.

Viscosity is a measure of resistance to flow.


40

Vapor Pressure and Boiling Point Vaporization – liquid changes to the gaseous

state When vaporization occurs at a temperature lower

than the boiling point, it is called evaporation.

Distributionof kinetic energies at two temper-atures


41

Vapor Pressure

Condensation – change from vapor to liquid

At equilibrium, the rate of condensation balances the rate of vaporization.

Equilibrium vapor pressure is the pressure exerted by the vapor above its liquid at a given temperature.

Sublimation is the direct conversion of a solid to a gas (dry ice is a good example).


42

Humidity

The equilibrium vapor pressure of a substance depends on the temperature and the pressure.

When relative humidity is reported as part of the weather, what is being reported is the ratio of the actual vapor pressure to the equilibrium vapor pressure.

At 100% relative humidity, the actual vapor pressure is the same as the equilibrium vapor pressure.


43

The Boiling Point

When the vapor pressure of a liquid equals the restraining pressure, bubbles of vapor form in the liquid and rise to the surface. The liquid is said to boil and the steady temperature at which this occurs is called the boiling point.


44

Normal Boiling and Melting Points

The normal boiling point and normal melting point are measured at 1 atm pressure.

The normal boiling point is the temperature at which the vapor and liquid are in equilibrium.

The normal melting point is the temperature at which the liquid and solid are in equilibrium.


45

Vapor Pressure, Boiling Point and Temperature

The b.pt.is the temperaturewhere thev.p. = 760torr.

B.pt. dependson inter-molecularforces betweenmolecules: weakdipole-dipole forces for diethylEther; hydrogenbonding for the other two (but stronger forwater)


46


Given the appropriate heats of fusion and vaporization, calculate the energy required to melt and vaporize a given compound.


47

11-5 Energy and Changes in State

Heat is involved in the phase changes.1. Heat of Fusion

amount of heat in calories or joules required to melt one gram of a substance (also expressed as J/mol).

Related to the strength of the intermolecular forces.

Melting is endothermic, while freezing is exothermic.


48

Energy and Changes in State

2. Heat of Vaporization Amount of heat in calories or joules

required to vaporize one gram of a substance (also expressed as J/mol).

Specific heat - amount of heat required to raise the temperature of one gram of substance one degree Celsius.


49

Objectives for Section 11-6

Describe the changes that occur on the molecular level in a heating curve.

Calculate the energy required for a change in temperature or physical state.


50

11-6 Heating Curve of Water

The graphical representation of the temperature as a solid is heated through the two phase changes plotted as a function of the time of heating is called the heating curve.


51

Heating Curve for Water

1. Heating ice (specific heat of ice)

2. Melting ice (heat of fusion of ice)

3. Heating water (specific heat of water)

4. Vaporizing water (heat of vaporization of water)

5. Heating steam (heating a vapor)

6. See Tables 11-2 and 11-3 for heats of fusion and vaporization


52

Heating Curve of Water

During heating, average kinetic energy of molecules increases. During phase transitions (melting and boiling), heat input is used to overcome intermolecular forces.


53

Calculating the Heat Involved in the Heating Curve

How much heat (in kJ) is needed to convert 250 g of ice at –15 oC to steam at 100 oC?

Solution

-15 oC 0 oC 0 oC 100 oC 100 oC

Meltings l

Boilingl gHeating

iceHeatingwater

250 g x 2.06 J

g oCx 15.0 oC

= 7.7 kJ

1 kJ

103 Jx

250 g x334 Jgx 1 kJ

103 J= 83.5 kJ

250 g x 4.18 J

g oCx 100 oC

= 105 kJ

1 kJ

103 Jx

250 g x2260 Jgx 1 kJ

103 J= 565 kJ

Total heat required =7.7 + 83.5 + 105 + 565= 761 kJ

chapter 11 the solid and liquid states. malone and dolter- basic concepts of chemistry 9e2 setting...

Documents

basic force

various forces

types of intermolecular

greater intermolecular

liquids flow

neighboring particles

physical properties

physical state solids