liquids and solids slides

AP Chemistry Rapid Learning Series - 17

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Liquids and Solids

AP Ch i t R id L i S i

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AP Chemistry Rapid Learning Series

Wayne Huang, PhDKelly Deters, PhDRussell Dahl, PhD

Elizabeth James, PhDDebbie Bilyen, M.A.



Learning Objectives

What intermolecular forces are

P ti f li id

By completing this tutorial you will learn…

Properties of liquids

How vapor pressure is affected by intermolecular forces and temperature

How solids are structured

How matter changes states

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The energy changes during a phase change

Concept Map

Chemistry

Studies

Previous content

New content

Matter

States

Has different

Solids Liquids

StructuresBonding

Structures

One is One is

Can change by

Have different possibilities of

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IntermolecularForces

Phase Changes

Vapor Pressure

g y

Involves breakingor forming

Boiling and melting point determined by



Intermolecular Forces

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Definition: Intra & Intermolecular

Intramolecular Forces –Chemical bonds within aChemical bonds within a molecule.

Intermolecular Forces (IMF) –Physical attractions between

t l l

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separate molecules.



London Dispersion ForcesAll molecules have London Dispersion Forces.

Due to a temporary “ganging up” of electrons on one side of the molecule.

δ-

+ +----

-- --

--

--

--Electrons

Nucleus

A th l t d th l l th

δ+

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As the electrons move around the molecule, they can temporarily end up on one side.

This results in a portion of the molecule that has a partially negative charge and a portion with a partial positive charge.

London Dispersion Forces PropertiesProperties of London Dispersion Forces:

TemporaryThe electrons continue moving and “spread out” p y g pagain.

Weakest IMF

All molecules have them

It’s temporary, and therefore weak.

All molecules have electrons which can “gain up”

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The larger the molecule, the larger the force

up”.

The more electrons in a molecule, the greater the effect of “ganging up”.



Dipole-Dipole ForcesPolar molecules exhibit dipole-dipole forces.

Polar molecules have a permanent partial separation of charges.

δ+ δ-

Polar Moleculee.g. sugar and water

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The opposite charges on separate polar molecules are attracted to one another.

They are not chemical bonds, just physical attractions between opposite charges.

Ion-Dipole ForcesAn ion and a polar molecule exhibit ion-dipole forces.

Polar molecules have a permanent partial separation of charges.

δ+ δ-

Polar Molecule

+

Catione.g. water and Na+

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Ions are attracted to the partially charged regions of a polar molecule.

They are not chemical bonds, just physical attractions between opposite charges.



Dipole Forces PropertiesProperties of Dipole Forces:

Permanent dipolesPartial separation of charge within the p gmolecule is permanent.

Stronger IMF

Only polar molecules

Although the IMF is not permanent, the ability to form the IMF is.

A permanent dipole (separation of charges) is

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can exhibit them

The stronger the dipole, the stronger the force

(separation of charges) is needed.

Stronger dipoles have greater attraction to other charges.

Hydrogen BondingHydrogen bonding is an especially strong case of dipole-dipole forces.Hydrogen atoms contain only 1 proton and 1 electron.

HH HH---- -- --OO

-- -- ----

δ+

δ-

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When hydrogen is bonded to a very electronegative atom (N, O or F), the separation of charges is very large as there are no other electrons around the hydrogen proton at all.

The hydrogen in this extreme dipole can be attracted to the lone pairs on an N, O or F atom on another molecule.



Hydrogen Bonding PropertiesProperties of Hydrogen Bonds:

Extreme case of dipole H has no other electrons, it is very positive when

forces sharing electrons with a very electronegative atom.

Strongest IMF

Molecules with an H on

Stronger than typical dipole forces

With lone pairs on N, O or

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Molecules with an H on an N, O or F can hydrogen bond

With lone pairs on N, O or F atoms in another molecule

Intermolecular Forces SummaryA summary of the 4 IMF’s:

Type of force Happens with Relative strength

All molecules

Polar molecules

Ion & polar molecule

Weakest

Medium

Medium

London Dispersion

Dipole-dipole

Ion-dipole

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Ion & polar molecule

H on an N, O or F with another N, O or F

Medium

Strongest

Ion dipole

Hydrogen bonding



Liquids

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Properties of LiquidsSome general properties of liquids:

Property Example

Definite volumePouring a drink into a larger glass doesn’t make more of the drink.

No definite shape Drink pools out when poured on a table.

A drop of food coloring

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Molecules are free to move past each other

Not very compressible

A drop of food coloring moves through the liquid over time.

A full bottle of water (no air) cannot be compressed.



Definition: Vapor Pressure

Vapor Pressure – Pressure t d b l bcreated above a sample by

particles evaporating from the sample and becoming gas particles.

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Vapor Pressure of a LiquidSolvent particles on the very top layer of the sample can evaporate.

Looking down on the top of a solution in a beaker:

Side viewGas particles now cause pressure-vapor pressure.

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Definition: Equilibrium

Equilibrium – Rate of change equals the rate of the opposite change.rate of the opposite change.

Rate forward = Rate reverse

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Vapor Pressure EquilibriumVapor Pressure equilibrium can be achieved in a closed system.

The rate of evaporation of a liquid is constantThe rate of evaporation of a liquid is constant.

evaporation

condensation

rate

time

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As the liquid evaporates, gas particles are formed which can then begin to condense down to the liquid form again.

As more gas particles are created, the rate of condensationincreases.

Equilibrium is reached when rate evaporation = rate condensation.



Volume of the Liquid and Equilibrium

The volume change as equilibrium is established.

Liquid volume decreases as initial gas particles are formed.

Liquid volume begins to increase as gas particles begin to re-condense.

Liquid volume is constant once equilibrium has been reached.

It’s less than the initial volume.

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Vapor Pressure and TemperatureSolvent particles on the very top layer of the sample can evaporate.

In order to evaporate, the particle must have enoughIn order to evaporate, the particle must have enough energy to break the intermolecular forces connecting it with the other liquid particles.

Higher temperature means the average kinetic

More molecules have the minimum

As temperature increases,

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average kinetic energy of the molecules is higher.

minimum energy needed to vaporize.

vapor pressure increases.



Solids

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Properties of SolidsSome general properties of solids:

Property Example

Definite volume A solid cannot “grow or shrink”

Definite shape

Molecules are not free to

Solids cannot spontaneously change shape

A dot of ink doesn’t travel

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Molecules are not free to move past each other

Not compressible

A dot of ink doesn t travel across the top of a desk

A piece of ice cannot be compressed



Types of Solids Structures - Part 1

Solids

Amorphous Solids

Crystalline Solids

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Definition: Amorphous and Crystalline

Amorphous Solid – Has a

Crystalline Solid – Has a

fair amount of disorder in the structure.

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Crystalline Solid Has a highly regular, repeating arrangement of particles.



Amorphous SolidsAmorphous solids are still rigid—they are still solid…

But they’re particles are not trapped in a repeating y p pp p gpattern as in crystalline solids.

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An example is glass.

The particles are trapped in a position before they have a chance to arrange themselves in a repeating pattern.

Crystalline SolidsCrystalline structures are lattice structures, composed of unit cells.

LatticeOverall crystalline

structure

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Unit CellSmallest repeating unit



Types of Unit CellsThese are the most simple unit cells are:

Cubic8 particles create a cube

Body-centered cubicA cube with a particle in the center

of the “body” of the cube

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Face-centered cubicA cube with a particle in the center of

each “face” of the cube

Types of Crystalline Structures

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Types of Solids Structures - the Rest

Solids

Amorphous Solids

Crystalline Solids

Atomic Solids

Molecular Solids

Ionic Solids

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Solids Solids Solids

Metallic Network

Definition: Atomic Solids

Atomic solids –Atoms are the particles in the unit cell.

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Metallic SolidsMetallic bonding contains metal atoms packed closely together and bonded to the atom in each direction equally.

This is called “closest packing”

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This is called “closest packing”.

The electrons form a large “pool” that are free to move throughout the structure.This allows metals to conduct electricity!

Network SolidsNetwork solids can be thought of as one giant molecule.

Graphite

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All of the atoms in the network solids are covalently bonded to their neighbors.

Examples of network solids include: graphite, diamonds, silica (sand)



Definition: Molecular Solids

SMolecular Solids – Molecules are the particles in the unit cell.

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Molecular SolidsStrong covalent forces within the molecule with weaker forces between the molecules.

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An example of a molecular solid is ice.



Definition: Ionic Solids

Ionic Solids – Structure containing positive and negative ions (cations and anions).

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Binary Ionic SolidsBinary ionic solids have a variation of the closest packing structures.

The larger ion is packed in the closest packing structure.

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The smaller ion fits in between the holes created by the larger ion.

This minimizes repulsions from like-charged ions.



Solid Structure Summary

Type Unit cell Bonding

SThere is no unit cell

Metal atoms

Non-metal atoms

Strong intramolecular, weaker intermolecular

Pool of electrons that are free to move

Covalent bonding throughout

Amorphous

Metallic

Network

allin

e

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Molecules

Ions

Covalent within molecule, weaker between

Electrostatic attraction between ions

Molecular

Ionic

Cry

sta

Solid Structure Properties

Type Properties

General properties of the different structure types

Disorder in their structure—no repeating pattern

Excellent conductors of heat and electricity, malleable and ductile

Brittle, poor conductors of heat and electricity

Amorphous

Metallic

Network

lline

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Strong bonds within but weak between (not much energy needed to melt, but a lot to break the chemical bond)

Stable, high melting points, brittle

Molecular

Ionic

Cry

stal



Phase Changes

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Definition: Phase Changes

Phase Change – Changes between solids, liquids and gases.liquids and gases.

Solid Liquid Gas

melting boiling

sublimation

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Solid Liquid Gas

freezing condensing

Deposition



IMF’s and Phase ChangesWhat role to intermolecular forces play in phase changes?

Break all IMF’s

Solid Liquid Gas

Break some IMF’s

Break rest of IMF’s

Form more IMF’s

Form some IMF’s

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Form IMF’s

Breaking IMF’s requires energy Forming IMF’s releases energy

Definition: Boiling & Melting Points

Boiling/Condensation Point – When liquid and gas phases are atliquid and gas phases are at equilibrium with one another.

Melting/Freezing Point – When solid

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Melting/Freezing Point When solid and liquid are at equilibrium.



How is Boiling Point Determined?Boiling point is where …

vapor pressure = atmospheric pressure.

When the vapor pressure of the liquid is great enough to form “bubbles” boiling will occur.

Atmospheric pressure

As temperature increases, the vapor pressure of the liquid increases

Bubbles rise (less dense)…boiling!

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increases.

As vapor pressure within the liquid increases, bubbles can begin to form—pushing against atmospheric pressure.

dense)…boiling!

How is Melting Point Determined? - 1

Melting point is where…vapor pressure liquid = vapor pressure solid.

Molecules escape the solid faster than they

The solid particles escape and join the liquid…but

After time…end up with all liquid.

Temperature is

Temperature 1: Solid vapor pressure > Liquid vapor pressure

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than they escape the liquid.

qthe liquid molecules aren’t crossing over as fast.

Temperature is above melting/freezing point.





Temperature 2: Solid vapor pressure < Liquid vapor pressure

Molecules escape the liquid faster th th

The liquid particles escape and join the solid but

After time…end up with all solid.

T t i

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than they escape the solid

the solid…but the solid molecules aren’t crossing over as fast

Temperature is below melting/freezing point



Molecules escape the liquid at the same pace as

The liquid particles escape and join the solid at the

At equilibrium. Amount of solid and liquid don’t h

Temperature 3: Solid vapor pressure = Liquid vapor pressure

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sa e pace asthey escape the solid.

the solid at the same rate the solid particles escape and join the liquid.

change over time.

Temperature is freezing point.



Why Do Some Solids Sublime?Why do some substances (such as dry ice) go straight from solid to gas?

The intermolecular forces are very weak.

The solid particles have enough energy to break all of the IMF’s and go straight to a gas rather than

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of the IMF s and go straight to a gas, rather than only breaking some of them and going to a liquid.

Phase Diagrams - 1Phase diagrams show what state of matter a substance would exist as at various temperatures and pressures.

re (a

tm)

Phase Diagram for H2O

s

l

melting Triple PointAll 3 states exist together

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0.0099 100Temperature (°C)

Pre

ssur

0.006

1s

g

boiling

sublimation

Critical PointPoint above which it cannot exist as a liquid



Phase Diagrams - 2Phase diagrams show what state of matter a substance would exist as at various temperatures and pressures.

freezing

re (a

tm)

Phase Diagram for H2O

s

l

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condensing

deposition

0.0099 100Temperature (°C)

Pre

ssur

0.006

1s

g

Energy of Phase Changes

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Definition: Enthalpy of Fusion

Enthalpy of fusion (Hfus) – energy py ( fus) gynecessary to break enough IMF’s to turn a solid into a liquid.

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Energy released when turning a liquid into a solid = - Hfus

Enthalpy of Fusion CalculationsEnthlapy of fusion is used in calculating energy needed to melt or released when freezing.

HmH ×=Δ H = enthalpy (energy)fusHmH ×=Δ m = mass of sample

Hfus = enthalpy of fusion (melting)(use –Hfus for freezing)

Example: Find the enthalpy of fusion of water if it takes 4175 J to melt 12.5 g of water

HgJ ×= )512(4175

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Hfus = ?m = 12.5 g H2OΔH = 4175 J

fusHgJ ×= )5.12(4175

fusHgJ=

5.124175

Hfus = 334 J/g



Definition: Enthalpy of Vaporization

Enthalpy of vaporization (Hvap) –py p ( vap)energy necessary to break the rest of the IMF’s and turn a liquid into a gas.

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Energy released when turning a gas into a liquid = - Hvap

Enthalpy of Vaporization Calculations

Enthalpy of vaporization is used in calculating energy needed to vaporize or released when condensing.

H th l ( )vapHmH ×=Δ H = enthalpy (energy)

m = mass of sampleHvap = enthalpy of vaporization (vaporizing)

(use –Hvap for condensation)

Example: If the heat of vaporization of water is 2287 J/g, how much energy is released when 15.75 g of water is condensed?

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Hvap = - 2287 J/gm = 15.75 g H2OΔH = ? J

⎟⎠⎞⎜

⎝⎛−×=Δ g

JgH 2287)75.15(

ΔH = -36020 J

Energy is released because it’s condensing.



Solids, Liquids & The AP Exam

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Liquids & Solids in the Exam

Using intermolecular forces to describe properties of d

Common Liquid & Solid problems:

compounds

What types of particles are at lattice points in different types of solids

The properties different types of solid bonding exhibit

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Multiple Choice Questions

Example: Which of the following has the highest melting point?

A. FB. ClB. ClC. BrD. IE. Same

Melting point is largely determined by intermolecular forces.

All of these are pure non-metal elements.

They have London Dispersion Forces only.

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The larger the atom, the more electrons it has, the more London Dispersion Forces it has.

The more London Dispersion Forces it has, the more energy is needed to break the IMF’s to melt the solid.

Answer: D

Free Response Questions

Example: Use chemical bonding and/or intermolecular forces to explain the following observations:

A. At STP, propane (C3H8) is a gas, while octane (C8H18) is , p p ( 3 8) g , ( 8 18)a liquid.

B. MgO melts at a much higher temperature than CO2C. Ethanol (CH3CH2OH) dissolves in water much more

easily than ethane (CH3CH3).

These are the two sub-questions you can

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answer after this tutorial.



Answering Free Response Questions

A. At STP, propane (C3H8) is a gas, while octane (C8H18) is a liquid.

Both molecules only have London Dispersion Forces. Propane is a larger molecule, therefore it has greater LDF. The greater

B. MgO melts at a much higher temperature than CO2

g , g gthe IMF’s, the more energy is needed to melt or boil a substance. C8H18 is larger, and therefore has more IMF’s, and therefore needs more energy to boil—which it hasn’t gotten at STP and is therefore a liquid still.

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MgO is an ionic compound while CO2 is a non-polar covalent compound. MgO has London Dispersion Forces and Ion-Ion interactions while CO2 only has LDF. The more IMF’s, the more energy is needed to pull apart molecules (melt) and therefore, the MgO melts at a higher temperature.

Matter undergoes phase changes,

Matter undergoes phase changes,

Solids and liquids are

condensed states of matters that

Solids and liquids are

condensed states of matters that

The energy changes during phase changes

d b

The energy changes during phase changes

d b

Learning Summary

Li id d lidLi id d lid

shown by phase diagrams.

shown by phase diagrams.

of matters that have

intermolecular forces.

of matters that have

intermolecular forces.

are governed by enthalpies of

fusion and vaporization.

are governed by enthalpies of

fusion and vaporization.

Solids have t t ith

Solids have t t ith

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Liquids and solids have vapor pressuresas particles escape to

the gas form.

Liquids and solids have vapor pressuresas particles escape to

the gas form.

structures—either amorphous or

crystalline (which can be atomic, molecular

or ionic).

structures—either amorphous or

crystalline (which can be atomic, molecular

or ionic).



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You have successfully completed the core tutorial

Liquids and Solids

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