unit ix solids, liquids heat problems chapter 16 part1 and chapter 14

UNIT IXSOLIDS, LIQUIDS HEAT PROBLEMS

CHAPTER 16 PART1 AND CHAPTER 14

INTERMOLECULAR FORCES

Forces between molecules

Not as strong as within molecules (covalent and ionic)

van der Waals Forces (Intramolecular Force)

Dispersion Forces (London Forces) Exists between non-polar molecules

weakest I.M.F. Due to temporary shifts in electron cloud density

ExamplesCH4O2

Dipole-Dipole Forces• slightly polar• Example:

CHCl3

HYDROGEN BONDING• VERY polar• Strongest• Examples

NH3 (N -- H)H2O (O -- H)HF (F-- H)HCl (Cl -- H)

SOLIDS AND LIQUIDS

SOLIDS

Orderly rigid and cohesive

Particles that vibrate around fixed points

SOLIDSCRYSTAL

• true solids• particles are arranged in an orderly repeating 3-D pattern

SOLIDS

CRYSTALS (cont)– consists of a MEMBER

o one particle (ion, atom, molecule

SOLIDS

several members together make up UNIT CELL

simplest repeating unitretains its shape

SOLIDS

several unit cells together make up a CRYSTAL LATTICE

3-D arrangement of unit cells repeated over and over

SOLIDSVocab-

ANHYDROUS (without water) - compound containing no water of hydration

HYDRATE-compound with water molecules attached(CuSO4 * 6H2O)

SOLIDS

AMORPHOUS –solid–no definite repeating pattern–no true melting point–no plateau

EXAMPLES: glass, butter, tar, plastic

LIQUIDS

DEFINITION– particles vibrate around a moving point

– non-orderly, non-rigid, cohesive– more space between particles than a solid

– exert a vapor pressure– Fluid – ability to flow

LIQUIDS

UNITSTemperature

average kinetic energy (KE) °C °F K (Kelvin)

LIQUIDS

VAPOR PRESSUREDefinition

pressure exerted by vapor molecules above a liquid when dynamic equilibrium is reached

LIQUIDS

Pressure measure of force with which gas molecules hit the side of container

normal atmospheric pressure at sea level

Standard Pressure Units =760 torrs = 760 mmHg = 101.3 kilopascals (kPa)

LIQUIDSVAPOR PRESSURE

Dynamic equilibrium - 2 opposite processes occurring at same time and same rate

VAPOR

LIQUID

LIQUIDSVAPOR PRESSURE

Dynamic Equilibrium depends upon:

Temperature - increase temperature, increase vapor pressure

T VP

VAPOR

LIQUID

LIQUIDS

• Strength of inter-molecular forces; hydrogen bonding(such as water) is strongest.

• increase forces; decrease vapor pressure

IMF VP

VAPOR

LIQUID

VAPOR VAPOR

LIQUID

WATER ALCOHOL

LIQUIDS

Viscosity - measure of resistance to flow (how thick)

Example – Molasses (syrup) has a high viscosity

Volatility - how easily a liquid evaporates

LIQUIDS

Very volatile:high vapor pressurelow IMFlow boiling pointEXAMPLES: alcohol, perfume

VAPORVAPOR

LIQUID

ALCOHOL

LIQUIDSNot volatile

low vapor pressurehigh IMFhigh boiling pointExamples: molasses, water

VAPORVAPOR

LIQUID

WATER

CHANGES IN STATE OR PHASES

Sublimation-– solid changes directly into gas without going through the liquid state

Examples: solid iodine, solid air fresheners, "dry" ice


Melting / Freezing– goes from solid to liquid or liquid to solid

Vaporization -• evaporation

occurs only on the surfaceat room temperaturecooling processSweat

• boiling occurs throughout the liquidrequires energy


Boiling Point:• vapor pressure = atmospheric (outside)

pressure (for any boiling point)• normal boiling point

vapor pressure = standard pressurestandard pressure = 1 atm, 760 torrs, 760 mm Hg,101.3 kPa


Boiling Point:• different altitudes

higher altitudes have lower air pressures

Denver has a lower boiling point 95 °C than Houston has (100 °C)

Foods take longer to cook in Denver than Houston.

VAPOR PRESSURE DIAGRAMS

1000

900

800

700

600

500

400

300

200

100

760

20 40 60 80 100

CHLOROFORM

ETHYL ALCOHOL

WATER

Temperature ( °C)

Vap

or p

ress

ure

(mm

Hg)

PHASE DIAGRAMS

Graphs that show conditions(temperature and pressure) under which a substance will exist as a solid, liquid, or gas.

PHASE DIAGRAMS

700

600

500

400

300

200

100

760

80 120 160Temperature (°C)

Pre

ssur

e (m

m H

g)

800

40 60 100 140 180

X

Z

X - Triple pointAll three states are in equilibrium at this temperature and pressure.

X-Y line - Theseare sublimation points.

Z - Critical temp. andpressure. A gas can'tbe liquified above thispoint.

PHASE DIAGRAMS

700

600

500

400

300

200

100

760


Pre

ssur

e (m

m H

g)

800

40 60 100 140 180

X

Z

SOLID

LIQUID

GAS

Lines represent 2 phases in equilibrium.

PHASE DIAGRAMS

700

600

500

400

300

200

100

760


Pre

ssur

e (m

m H

g)

800

40 60 100 140 180

X

Z

Normal boiling point(condensation) occurswhen standard pressure crosses liquid / gasline

Normal boiling point(condensation) occurshere.

PHASE DIAGRAMS

700

600

500

400

300

200

100

760


Pre

ssur

e (m

m H

g)

800

40 60 100 140 180

X

Z

Normal melting point(freezing) occurs wherestandard pressure crossesliquid / solid line.Normal melting point(freezing) occurs here

PHASE DIAGRAMS

700

600

500

400

300

200

100

760

80 120 160Temperature ( °C)

Pre

ssur

e (m

m H

g)

800

40 60 100 140 180

X

Z

Freezing or

melting point

Boiling or

condensa

tion

point

Deposition or sublimation point

UNIQUE PROPERTIES OF WATER

STRONG HYDROGEN BONDING CAUSES:

– high boiling point and melting point

– high specific heat capacity– high surface tension

needle floats– Water droplets are spherical

HEAT VS. TEMPERATURE

Energy transferred from one body to another because of a difference in temperature

Average Kinetic Energy

Written as KE

HEAT VS. TEMPERATURE UNITS

– calories (c)– kCal - C

• (1000 calories)– Joules - J

• energy for one heartbeat

– 1 cal = 4.18 J– 1 kCal = 4180 J

UNITS– °C - celsius

– °F -Fahrenheit

– K - kelvin (no degree sign!)


Measured by:– indirectly by a

calorimeter

Measured by:– thermometer


DEPENDS UPON– mass

more mass means more heat

– Cp (S) - specific heat type of matter some hold heat better than others

– T - change in temperature

DEPENDS UPON– amount of

movement of the particles in the substance


FORMULA q=energy (J) m=mass (g)

q = (m) (T) (Cp) q = (m) (T2-T1) (Cp)

Specific Heat or Heat Capacity

Amount of heat needed to raise 1 gram of a substance 1 degree Celsius

Units– (J/goC) – (cal/goC)

Examples– water --- 4.18 J/goC or 1 cal/goC– Au --- 0.129 cal/goC– alcohol --- 2.45 J/goC

Calorie

Amount of heat needed to raise one gram of water one degree of celsius

It takes one calorie to raise one gram of water one degree of Celsius

Heat of Fusion - Hf

Amount of heat needed to melt one gram of a substance at its melting point

Units(cal/g)

Examples– water (Hf) = 334 J/g or 76.4 cal/g– Ag = 88 J/g

HEAT OF VAPORIZATION - Hv

Amount of heat needed to vaporize one gram of a substance at its boiling point

Examples– water (Hv) = 2260 J/g or 539 cal/g– Pb = 858 J/g

PHASE CHANGE DIAGRAMS

SOLID

LIQUID

GAS

TE

MP

ER

AT

UR

E (

C)

o

HEAT (cal/g) OR TIME

100

0

Heat of fusion -Melting point - substance is becoming a liquid

Heat of vaporizationBoiling point- substance is becoming a liquid

WATER

Heat Calculations - Formulas

The state remains the same and there is no change in temperature.

q= joules m=grams Cp=J/g or J/c

q= (m) (Cp)q = Heat

Example of Non-Changing State

Melting/freezing at melting pointVaporizing/condensing at boiling point

How much energy does it take to melt55g of gold at its melting point?Cp = 64.5 J/gq= (m) (Cp) = (55g)(64.5 J/g) = 3547.5 J

HEAT EQUATION

One substance with a temperature change

q=joules (J)m= mass (g)Cp = specific heat capacity (J/g °C) (J/c °C)T2 = final temperatureT1 = initial temperature

q = (m) (Cp) (T2-T1)

HEAT EQUATION EXAMPLE

***Heating or cooling with no change in state***

How much energy is released as 33 gof solid silver cools from 95 °C to 60°C?

Cp of silver = 0.236 J/g °C

HEAT TRANSFER EQUATION

How a substance changes the temperature of another substance used in calorimeter calculations

(m1) (Cp1) (T2-T1) = (m2) (Cp2) (T2-T1)Warm substancelosing energy

Cool substancegaining energy

Energy LOST = Energy GAINED

HEAT TRANSFER EQUATION EXAMPLE

A piece of metal is dropped into a beaker ofboiling water whose temperature is 95 C. The 5g piece of metal is put into 100g of coldwater at 20 C. The temperature of the waterrises to 30 C. What is the specific heat of the metal?

Cp(water) = 4.18 J/g C

o

o

o

o

EQUATION FOR CHANGING TEMPERATURE AND STATESDraw the phase change diagram

CHANGING STATES AND TEMPERATURE

TE

MP

ER

AT

UR

E (

C)

o


100

0 Use the following equations:q = (m) (Cp)q = (m) (Cp) (T2-T1)


TE

MP

ER

AT

UR

E (

C)

o


100

0

1. Heat solid tomelting point

q = (m) (Cp) (T2-T1)


TE

MP

ER

AT

UR

E (

C)

o


100

0

2. Melting solidto liquid

q = (m) (Cp)


TE

MP

ER

AT

UR

E (

C)

o


100

0

3. Heat liquidto boiling point

q = (m) (Cp) (T2-T1)


TE

MP

ER

AT

UR

E (

C)

o


100

0

4. Change liquidto gas

q = (m) (Cp)


TE

MP

ER

AT

UR

E (

C)

o


100

0

5. Heating gas

q = (m) (Cp) (T2-T1)

CHANGING STATES AND TEMPERATURES

1. Heat solid to melting point : KE2. Melt solid to liquid: PE3. Heat liquid to boiling point: KE4. Change liquid to gas: PE5. Heat gas: KE

q = (m) (Cp) (T2-T1)

q = (m) (Cp)

q = (m) (Cp) (T2-T1)

q = (m) (Cp)

q = (m) (Cp) (T2-T1)

When to use which equations:

CHANGING TEMPERATURE AND CHANGING STATES EXAMPLE

How much energy is needed to change 30g of ice at -5 °C to steam at 120 °C?

unit ix solids, liquids heat problems chapter 16 part1 and chapter 14

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