Chapters 13 & 17 Phases and Heat
Mar 26, 2015
Chapters 13 & 17
Phases and Heat
Phases of Matter
Chapter 13
Phases
There are three phases, or states, that we will discuss
Solid Liquid Gas
Solids
form of matter that has a definite shape and definite volume.
Use (s) to denote solids in chemical reactions
Solids
In most solids the atoms, ions, or molecules are packed tightly together
The particles in solids tend to vibrate around fixed points
When you heat a solid, its particles vibrate more rapidly, eventually the solid breaks down and melts.
Types of Solids
Crystalline Solids In a crystal the particles are arranged in
an orderly, repeating, three-dimensional pattern called a crystal lattice. There are many different shapes of crystalline solids, pg 397
Types of Solids
Non-Crystalline Solids Amorphous solids lack an orderly internal
structure. Ex – Rubber, plastic, and asphalt.
Glass – transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. Sometimes called super-cooled liquids.
Allotropes
Two or more different molecular forms of the same element in the same physical state Different properties because they have
different structures
Allotropes of Carbon
Liquids
form of matter that has a definite volume, indefinite shape, and flows.
Use (l) to denote liquids in chemical reactions
Liquids
In liquids the atoms or molecules are able to slide past each other.
In liquids there are intermolecular attractions between the atoms or molecules, which determine the liquid’s physical properties.
When you heat a liquid the particles vibrate more rapidly and start moving past each other faster.
Gases
form of matter that takes both the shape and volume of its container
Use (g) to denote gases in chemical reactions
Phase Changes
Six Changes Solid Liquid Melting Liquid Solid Freezing Liquid Gas Vaporization Gas Liquid Condensation Solid Gas Sublimation Gas Solid Deposition
Phase Changes
During any given phase change, both phases can exist together in equilibrium
Example At 0°C, water can exist in both the liquid
and solid phases in equilibrium
Energy
When energy is added to a reaction, or phase change, it is called Endothermic
When energy is released during a reaction, or phase change, it is called Exothermic
Phase Changes
Which phase changes are endothermic, requiring the addition of energy?
Melting Vaporization Sublimation
Phase Changes
Which phase changes are exothermic, releasing energy?
Freezing Condensation Deposition
Phase Diagram of CO2
Energy
What is energy? Capacity to do work Ability to do work
Two main types Kinetic Potential
Types of Energy
Kinetic Energy Energy of motion Related to the speed and mass of
molecules
Potential Energy Stored energy
Temperature
How is energy related to Temperature?
What happens to the temperature of a substance when you add energy? Particles move faster Temperature increases
Temperature
Relationship between energy, particle speed, and temperature
Temperature Definition Average Kinetic Energy
Temperature Scales
Kelvin (K) and Celsius (°C) scales Kelvin scale is called the absolute
scale Related to the kinetic energy of a
substance Celsius scale is a relative scale
based on the boiling and freezing points of water
Temperature Conversion
K = °C + 273
Pressure
What is pressure?
Physics – Force per unit area
Chemistry – related to the number of collisions between particles and container walls
Pressure Conversion
1 atm = 101.3 kPa
Vapor Pressure
Pressure exerted by vapor that has evaporated and remains above a liquid
Related to temperature As temperature increases, vapor
pressure increases
Boiling vs. Evaporation
Boiling Vapor pressure equals external, or
atmospheric pressure
Evaporation Some molecules gain enough energy to
escape the liquid phase At temp. less than boiling point
Normal Boiling Point
Boiling Point at Standard Pressure
1 atm or 101.3 kPa
Evaporation
Why is evaporation considered a cooling process?
As the molecules with higher kinetic energy evaporate, the average kinetic energy of the substance decreases
Table H
Shows the relationship between temperature and vapor pressure for four specific substances
Thermochemistry
Chapter 17
Thermochemistry
Heat involved with chemical reactions and phase changes
Heat
Energy transferred from one object to another, usually because of a temperature difference
Measured in Joules (J) or calories (cal)
Heat flows from hot to cold
Heat Transfer
Endothermic Energy being added
Exothermic Energy being released
Specific Heat Capacity
Amount of heat needed to raise the temperature of 1 g of a substance by 1°C Unique for each phase of each substance
4.18 J/(g*°C) for liquid water Listed in Table B of Reference Tables
Heat
What factors affect the amount of heat transferred?
Specific Heat Capacity Mass Temperature difference between objects
Heat Equation
Heat, Q Mass, m Specific Heat Capacity, c Change in Temperature, ΔT
Q=m*c*ΔT
Example
200g of water is heated from 20°C to 40°C, how much heat is needed?
Q = m*c*ΔT Q = (200g) * (4.18J/g°C) * (20°C) Q = 16720 J
Example
How much energy is required to raise the temperature of 50g of water from 5°C to 50°C?
Q = m*c*ΔT Q = (50g) * (4.18J/g°C) * (45°C) Q = 9405 J
Another Example
What is the Specific Heat Capacity of Fe, if it takes 180J of energy to raise 10g of Fe from 10°C to 50°C?
Q = m*c*ΔT 180J = (10g) * c * (40°C) c = 0.45 J/(g*°C)
Phase Change
At what temperature does ice melt? 0°C
At what temperature does water freeze? 0°C
Melting point and freezing point are the same
Phase Change
What happens to temperature during phase changes? Temperature remains constant
Temperature remains CONSTANT during a phase change
Phase Change
If energy is being added, what kind of energy is it? Energy being added is potential energy,
not kinetic energy Potential energy is being used to
separate or spread the particles apart
Heat of Vaporization, Hv
Amount of energy needed to vaporize 1g of a substance
Water = 2260 J/g
Q=mHvUse for Liquid Gas orGas Liquid
Heat of Fusion, Hf
Amount of energy needed to melt 1g of a substance
Water = 334 J/g
Q=mHfUse for Solid Liquid orLiquid Solid
Examples
How much energy is needed to melt 10g of ice at 0°C? Q = m*Hf
Q = (10g) * (334J/g) Q = 3340 J
Example
How much energy is needed to vaporize 10g of water at 100°C? Q = m*Hv
Q = (10g) * (2260J/g) Q = 22600 J
Phase Change
Which requires more energy melting or vaporization? Vaporization
Why? Molecules are spread farther apart as a
gas It takes more energy to get gas particles
spread apart
Heating (Cooling) Curves
Shows relationship between temperature and time during constant heating or cooling.
Also shows phases, and the phase changes between them.
Heating Curves
Diagonal lines are phases Horizontal lines are phase changes
Time (s)
Tem
p (˚
C)
Gas
Liquid
Solid
Heating Curves
Diagonal lines are phases Horizontal lines are phase changes
Time (s)
Tem
p (˚
C)
VaporizationCondensation
MeltingFreezing
Conservation of Energy
Energy can not be created or destroyed, only transferred or converted from one form to another.
Energy lost by one object must be gained by another object or the environment Qlost = Qgained
Example
A chunk of iron at 80°C is dropped into a bucket of water at 20°C.
What direction will heat flow? From the iron to the water Hot to cold
Example
A chunk of iron at 80°C is dropped into a bucket of water at 20°C.
What could be the final temperature, when they both come to equilibrium? Between 20°C and 80°C
Example
A 100g block of aluminum, c=0.90J/g*°C at 100°C is placed into 50g of water at 20°C, what will be the final temperature when the aluminum and water reach equilibrium? Qlost = Qgained
m*c*ΔT = m*c*ΔT 100g*0.90J/g°C*(100°C-Tf) = 50g*4.18J/g°C*(Tf-20°C) 90*(100-Tf) = 209*(Tf-20) 9000-90Tf = 209Tf-4180 13180 = 299Tf
Tf = 44°C
Conservation of Energy
SystemEnergy In Work Done (Energy)
Energy Out
Conservation of Energy
Energy In Work Done (Energy)
Energy Out
Conservation of Energy
Food In Metabolism
Waste Out