unit 9: liquids & solids. three states of matter state shapevolumewhy? particles far apart;...

Unit 9: Liquids & SolidsUnit 9: Liquids & Solids

Three States of MatterThree States of Matter

State Shape Volume Why?

Gas None None Particles far Particles far apart; apart;

forcesforcessmallsmallLiquid None Fixed Particles Particles

closercloser forces greaterforces greater

Solid Fixed FixedParticles touchParticles touchforces greatforces great

Why Gas Laws and Not Solid or Why Gas Laws and Not Solid or Liquid Laws?Liquid Laws?

Gases are mainly empty space; have weak attractions between

molecules

Solids/liquids have particles which are closer together and

have more varied forces between particles.

Phase TransitionsPhase TransitionsDefinition:Definition:

PhysicalPhysical changes that changes that result in changes of state.result in changes of state.

ALL phase changes involve ALL phase changes involve

energy (enthalpy). energy (enthalpy).

Review: Review: Enthalpy = Enthalpy = ΔΔH = heat at H = heat at const Pconst P

Phase Changes that Require Energy

If you have to put energy into a reaction to make it happen, it is an endothermic reaction.

Endothermic Phase Changes: Melting (solid liquid) a.k.a

“fusion” Vaporization (liquid gas) Sublimation (solid gas)

Phase Changes that Release Energy

If energy is released or given off by a reaction, it is an exothermic reaction.

Exothermic Phase Changes: Condensation (gas liquid) Deposition (gas solid) Freezing (liquid solid)

NOTE: What is constant at every phase change above????

VaporizationVaporization Definition: conversion of liquid to

gas.(endothermic)

We commonly call this evaporation. Condensation is the reverse

process(exothermic)

LIQUID VAPOREvaporationEvaporation

condensationcondensation

Enthalpy of VaporizationEnthalpy of Vaporization Heat that is absorbed to vaporize a given amount

of liquid at a constant temperature and pressure.

Units: Energy (J, kJ or cal) per g or mol

LIQUID + heat ---> VAPOR

Compd. ∆Hvap (kJ/mol)

H2O 40.7 (at 100 oC)

SO2 26.8 (-47 oC)

Xe 12.6 (-107 oC)

Heat of Condensation is same Heat of Condensation is same quantity but opposite in sign!quantity but opposite in sign!

∆Hvap = - ∆Hcond

VAPOR ---> LIQUID + heatVAPOR ---> LIQUID + heat

∆∆HHcondcond for water = - 40.7 kJ/mol for water = - 40.7 kJ/mol

Example 9. 1: How many kilojoules of heat are required to vaporize 598.5 g of ethanol? The heat of vaporization is 43.3 kJ/mol.

Vapor PressureVapor Pressure

Definition: Pressure exerted by a vapor in a closed flask in equilibrium with its liquid

Equilibrium: two opposing processes occur at same rate.

A: Rate vap> Rate cond

B: Rate vap = Rate cond

Example 9.2 Liquid ethanol has a vapor pressure of 43.9 mm Hg at 20 deg. C.

What is the minimum volume of a flask needed to vaporize 1.00 g of liquid ethanol?

Graphing Vapor PressureGraphing Vapor Pressure

As Temperature increases, vapor pressure increases.

As attractive forces between molecules increase, vapor pressure decreases.

Attractive Forces:

Water > Ethanol > Ether

Liquids with higher vapor pressure at a given T are said to be more volatile.

Questions:

1. What does this graph tell you about the relative attraction between molecules for substances a – e?

2. Which substance is most volatile?

Boiling PointBoiling Point A liquid boils when

vapor pressure = atmospheric pressure.

Normal Boiling Point: Temp. where P = 760 mm Hg (1 atm) on vapor pressure curve.

Dependency on pressureAs pressure increases, boiling point increases; as pressure decreases, boiling point decreases.

Normal Boiling PointNormal Boiling Point

Melting and Freezing PointMelting and Freezing PointConversion of a solid to a liquid is called

melting, or fusion.

Conversion of a liquid to a solid is called freezing.

The freezing point = melting point

Energy needed to melt a given quantity of solid is called the enthalpy of fusion, or Δ H fus.

Example 11.3 How much energy is required to melt 100.0 g of ice? The heat of fusion of water is 6.01 kJ/mol.

Example 11.4 How much energy in kJ is required to heat 100.0 g of liquid water from zero to 100 deg. C, and then vaporize all of it?

Strategy:

1. Calculate energy needed to heat water (Q equation from Chapter 6)

2. Calculate how much energy is needed to vaporize the water.

3. Add the two amounts together.

Example 11.4 How much energy in kJ is required to heat 100.0 g of liquid water from zero to 100 deg. C, and then vaporize all of it?

Example 11.5 How much energy is required to heat 75.0 g of ice from 0.0 deg. C to 185.0 deg. C? The heat capacity of steam is 1.84 J/g deg. C.

Strategy:

1. Calculate energy needed to melt ice.

2. Calculate how much energy is needed to heat the water to boiling.

3. Calculate how much energy is needed to vaporize the water.

4. Calculate how much energy is needed to heat the steam.

5. Add the four amounts together.

Example 11.5 How much energy is required to heat 75.0 g of ice from 0.0 deg. C to 185.0 deg. C? The heat capacity of steam is 1.84 J/g deg. C.

Cooling Curve: H2O (g) H2O (s)

IntraIntramolecular Forcesmolecular Forces

The attractive forces that hold particles together in ionic, covalent and metallic bonds are called intramolecular forces

“Intra-” prefix = within The forces inside a molecule holding the

individual atoms together Ex.) Covalent bonds

Ionic bondsMetallic bonds

InterIntermolecular Forces (IMF’s) :molecular Forces (IMF’s) : Bonds Bonds between between MoleculesMolecules

1.1. Much weaker than Chemical Much weaker than Chemical Bonding Bonding withinwithin moleculesmolecules

2.2. Chemical Bonds (ionic and covalent) Chemical Bonds (ionic and covalent) determine determine chemicalchemical propertiesproperties

3.3. Intermolecular forces determine Intermolecular forces determine physicalphysical properties properties

e.g. density, mp, bp, solubility, vapor e.g. density, mp, bp, solubility, vapor pressure, etc.pressure, etc.

Three General Types of Three General Types of Intermolecular Forces (IMF’s)Intermolecular Forces (IMF’s)

Dispersion (London, van der Waals)

Dipole/dipole

Hydrogen Bonding

London Dispersion ForcesLondon Dispersion ForcesDefinition: IMF between two non-polar

molecules formed by temporary positive and negative attractions due to the shifting of electron cloud.

Found in all substances, but become important when they are the only IMF present.

Strength increases as molar mass increases.

Formation of a dipole in two nonpolar I2 molecules:

London Dispersion ForcesLondon Dispersion Forces

Higher molar mass ---> larger Higher molar mass ---> larger dispersion forcesdispersion forces

MoleculeMolecule Boiling Point Boiling Point ((ooC)C)

CHCH44 (methane) (methane) - 161.5- 161.5

CC22HH66 (ethane) (ethane) - 88.6 - 88.6

CC33HH88 (propane) (propane) - 42.1- 42.1

CC44HH1010 (butane) (butane) - 0.5- 0.5

Higher boiling point means GREATER IMF’s!

London Dispersion ForcesLondon Dispersion Forces

Example 11.6. Account for the fact that chlorine is a gas, bromine is a volatile liquid, and iodine is a volatile solid at room temperature.

Dipole-Dipole ForcesDipole-Dipole Forces

Definition: Attractions between oppositely charged regions of polar molecules.

Caused by attraction of one dipole for another.

Present in all polar substances!

Solubility and dipole-dipole forces:

“like dissolves like”

Dipole-Dipole ForcesDipole-Dipole Forces

Hydrogen BondingHydrogen Bonding A special type of dipole-dipole force occurring

only between molecules with a H atom bonded to either a F, O, or N atom.

How to recognize: F, O, or N directly bonded to H

Two reasons why hydrogen bonds are stronger than dipole-dipole forces:

a. F, O, N very electronegativeb. H is a small atom

Hydrogen bonding is FON!

Hydrogen BondingHydrogen Bonding

Hydrogen Bonding in Hydrogen Bonding in HH22OO

Hydrogen Bonding in Hydrogen Bonding in HH22OO

H-bonding is especially strong in water because

The O—H bond is very polar

There are 2 lone pairs on the O atom

Accounts for many of water’s unique properties.

H bonds ---> abnormally high specific heat capacity of water (4.184 J/g•K)

This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.

Hydrogen Bonding in HHydrogen Bonding in H22OO

Hydrogen BondingHydrogen BondingHydrogen BondingHydrogen Bonding

H bonds H bonds

lead tolead to

abnormallabnormallyy

highhigh

boiling boiling

point of point of

water.water. See Screen 13.7See Screen 13.7

Hydrogen Bonding in DNA Hydrogen bonding plays a key role in

maintaining the double helix structure of DNA

Example:Identify the types of intermolecular forces present in compounds of:

Hydrogen Fluoride

Pentane (C5H12)

Hydrochloric Acid

Ethanol (Ethyl Alcohol)

Example Rank these substances in terms of increasing

boiling point.

N2 CCl4 CH3Cl NH3

Liquids Viscosity – a measure of the resistance of a

liquid to flow The particles in a liquid are close enough together

that their attractive forces slow their movement as they flow past one another

The stronger the attractive forces (intermolecular forces), the more viscous the liquid is.

As temperature increases, viscosity decreases.

Surface tension – an inward force that tends to minimize the surface area of a liquidA measure of the inward pull by particles in the interiorThe stronger the intermolecular forces, the higher the

surface tension

Liquids

In water, this is due mainly to

hydrogen bonding!

Liquids

Surfactant – any substance that interferes with the hydrogen bonding between water molecules & reduces surface tension

Surfactants used to clean up oil spills as wellExxon Valdez oil spill in 1989 spilled over 700,000

barrels of oil into the water near Alaska

Network Covalent SolidsNetwork Covalent SolidsGiant molecules connected by strong covalent bonds

Properties: hard, high mp, nonconductors

Network Covalent SolidsAtoms that can form multiple covalent bonds

(look for C, Si, and other Group 14 elements)

are able to form network covalent solids.

All atoms in the entire structure are bonded together with covalent chemical bonds.

Physical Properties of Graphite vs. Physical Properties of Graphite vs. DiamondDiamond

PropertyProperty GraphiteGraphite DiamondDiamondDensity Density (g/mL)(g/mL)

2.27 3.51

Hardness Hardness Very soft Very hardColorColor Shiny black Colorless/

transparentElectrical Electrical ConductivityConductivity

High None

HHcombcomb (kJ/mol)(kJ/mol)

-393.5 -395.4

Metallic and Ionic Solids: See Metallic and Ionic Solids: See Chem Act. 24 and 25Chem Act. 24 and 25

Metallic and Ionic Solids: See Metallic and Ionic Solids: See Chem Act. 24 and 25Chem Act. 24 and 25

Ionic SolidsIonic Solids A compound where each cation is

simultaneously attracted to an anion.

Review: How can we identify an ionic compound?

Properties of Ionic SolidsProperties of Ionic Solids1. Molecules, atoms or 1. Molecules, atoms or

ions locked into a ions locked into a CRYSTAL CRYSTAL LATTICELATTICE

2. Particles are CLOSE 2. Particles are CLOSE togethertogether

3. STRONG IM forces3. STRONG IM forces

4. Highly ordered, rigid, 4. Highly ordered, rigid, incompressibleincompressible

ZnS, zinc sulfideZnS, zinc sulfide

Metallic SolidsMetallic Solids

A solid consisting of entirely metals.

Characteristics: Electrons are “delocalized” (they can move freely)

Good conductors, malleable, ductile

Metallic Solids Metallic solids – positive metal ions surrounded

by a sea of mobile electrons Mobile electrons make metals malleable and ductile

because electrons can shift while still keeping the metal ions bonded in their new places

Metallic solids are good conductors of heat and electricity

Metallic Bonds

Amorphous Solids Amorphous solid – a solid in which the

particles are not arranged in a regular, repeating pattern “Amorphous” = “without shape” Often form when a molten material cools too

quickly to allow enough time for crystals to form Common examples: glass, rubber, many

plastics

Types of SolidsTypes of Solids

TYPETYPE EXAMPLEEXAMPLE FORCEFORCE

Ionic Ionic NaCl, CaFNaCl, CaF22, ZnS, ZnS Ion-ionIon-ion

MetallicMetallic Na, FeNa, Fe MetallicMetallic

MolecularMolecular Ice, IIce, I22 DipoleDipoleInd. dipoleInd. dipole

NetworkNetwork DiamondDiamond ExtendedExtendedGraphiteGraphite covalentcovalent