unit ix solids, liquids heat problems chapter 16 part1 and chapter 14
TRANSCRIPT
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
CHANGES IN STATE OR PHASES
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
CHANGES IN STATE OR PHASES
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
CHANGES IN STATE OR PHASES
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
80 120 160Temperature (°C)
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
80 120 160Temperature (°C)
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
80 120 160Temperature (°C)
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!)
HEAT VS. TEMPERATURE
Measured by:– indirectly by a
calorimeter
Measured by:– thermometer
HEAT VS. TEMPERATURE
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
HEAT VS. TEMPERATURE
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
HEAT (cal/g) OR TIME
100
0 Use the following equations:q = (m) (Cp)q = (m) (Cp) (T2-T1)
CHANGING STATES AND TEMPERATURE
TE
MP
ER
AT
UR
E (
C)
o
HEAT (cal/g) OR TIME
100
0
1. Heat solid tomelting point
q = (m) (Cp) (T2-T1)
CHANGING STATES AND TEMPERATURE
TE
MP
ER
AT
UR
E (
C)
o
HEAT (cal/g) OR TIME
100
0
2. Melting solidto liquid
q = (m) (Cp)
CHANGING STATES AND TEMPERATURE
TE
MP
ER
AT
UR
E (
C)
o
HEAT (cal/g) OR TIME
100
0
3. Heat liquidto boiling point
q = (m) (Cp) (T2-T1)
CHANGING STATES AND TEMPERATURE
TE
MP
ER
AT
UR
E (
C)
o
HEAT (cal/g) OR TIME
100
0
4. Change liquidto gas
q = (m) (Cp)
CHANGING STATES AND TEMPERATURE
TE
MP
ER
AT
UR
E (
C)
o
HEAT (cal/g) OR TIME
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?