Unit 3 [A] Section 1A
States of Matter
Properties / Solid / Liquid / GasPacking:
closeness of particles
Attractive Forces between particles
Movement / Vibrate in place
Can’t switch places
Low ______of particles / Rapid Random chaotic motion
High ______of particles
Shape
Volume:
how much space is taken up
Compressibility
ability for particles to move closer together
Changes in State
Endothermic physical changes of state
o Kinetic Energy must be put INTO the substance in order to increase the ______of the molecules so as to break the ______forces holding the particles together
ü Melting: change of state from a solid to a ______
ü Vaporization (Boiling or Evaporation): change of state from a ______to a gas
ü Sublimation: direct change of state from a ______to a gas
Exothermic physical changes of state
o Kinetic Energy must be taken OUT (removed) the substance in order for the molecules to ______down so that the ______forces can begin to hold the particles together
ü Freezing: change of state from a ______to a solid
ü Condensation: change of state from a gas to a ______
ü Deposition: direct change of state from a ______to a solid
Temperature of State Changes
ü Freezing point (fp) is the temperature at which the liquid turns into a solid
ü Melting point (mp) is the temperature at which the solid turns into a liquid
ü Freezing Point = Melting Point
Example: Water has a mp and a fp of 0 °C
ü Boiling point (bp) is the temperature at which the liquid turns into a gas
ü Condensation point (cp) is the temperature at which a gas turns into a liquid
ü Boiling Point = Condensation Point
Example: Water has a bp and a cp of 100 °C
****All substances have their own specific freezing and boiling point which makes this physical property a great way to identify an unknown substance. ****
Unit 3 [A] Section 1 B
Atmospheric Pressure vs Vapor Pressure
Pressure / Atmospheric Pressure / Vapor PressureForce per unit area created as gas molecules collide with objects / force per unit area exerted against a surface by the weight of the air molecules above the surface / Force per unit area of the gas molecules above a liquid colliding
Usually measured in newtons/meter 2 but in chemistry we use atmospheres (atm) or millimeters of mercury
( ______) / The more air molecules above a surface, the more molecules to exert a force and thus a ______air pressure / The ______the attractive forces, the higher the vapor pressure
At sea level, atmospheric pressure equals ______atm or 760 mm Hg / Substances with high vapor pressure are called ______
There are only 2 factors that control VAPOR PRESSURE!
1. Temperature
2. Attractive forces of the liquid
® When the vapor pressure of a liquid = atmospheric pressure, boiling will occur.
® Normal boiling point is @ 760 mmHg
Vaporization: Difference between Evaporation and Boiling
® Evaporation occurs spontaneously at all temperatures at the ______of the liquid
® Boiling occurs when extra ______known as heat is added and takes places within the body of the liquid
® Boiling occurs at only 1 temperature dependent on pressure
Another definition of Boiling Point
ü When external atmospheric pressure = vapor pressure of a liquid
Ø At 90°C, the water’s vapor pressure is not strong enough to push against the atmospheric pressure à starting to boil
Ø At 100°C, the water’s vapor pressure is equal to the atmospheric pressureà boiling is occurring
ü Since atmospheric pressure changes at various altitudes, we use “______” boiling point to describe the temp at which a LàG at 1 atm
Important Ideas
ü The higher in altitude the ______the atmospheric pressure
ü At higher altitudes, the boiling point is ______
ü It takes ______to cook foods at higher altitudes (lower atmospheric pressures)
Unit 3 [A] Section 1 C
Heating & Cooling Curves
Ø A diagram that shows how solids, liquids & gases change state when ______is changed
ü Plateaus = the changes of state (freezing, melting, etc.)
ü Freezing Point & Melting Point are at the temperature or at the same plateau
ü Boiling Point &Condensation Point are at the same temperature or at the same plateau
ü Slopes= pure states (solid, liquid or gas)
ü At the plateaus, ______energy remains constant because temperature remains constant while potential energy changes
ü At the slopes, kinetic energy ______because temperature changes while potential energy remains constant
**DANGER!!** Notice that a gas can get higher than boiling point!
SELF CHECK:
1. What is the boiling point of the substance?
2. What letter represents the solid state only?
3. What letter represents the melting process?
SELF CHECK:
1. While the substance is cooling during the liquid phase, the average kinetic energy of the molecules of the substance:
Increases, Decreases or Stay the Same
2. What is the freezing point of this substance?
3. How long does it take for the gas to completely liquefy?
Phase Diagram
Ø A diagram that shows how solids, liquids & gases change state as both temperature and ______are changed
ü Crossing a line between states determines the change of state (boiling, melting, etc)
ü A point directly on a line will identify the pressure and temperature (boiling point, melting point, etc.) of the phase change
ü ______is the temperature and pressure in which all 3 of the states coexist
ü ______is the temperature & pressure at which a gas can no longer liquefy
Important information regarding the Phase Diagram of Water:
Important Information regarding the phase diagrams of water and carbon dioxide
SELF CHECK: See diagram
[1] What is the temperature (freezing point) of line B at 1 atm?
[2] What is the temperature (boiling point) of line Cat 1 atm?
[3] What is point D?
[4] What is point E?
[5] What change of state happens when you cross line B at a constant pressure of 10 atm and increase temperature?
[6] What change of state occurs when you cross line A at constant pressure of .001 atm?
[7] What change of state happens when you cross line C at 400 K to 300 k at approximately 5 atm?
Unit 3 [A] Section 2
Properties of Matter
Physical Property / Chemical Property· Can be determined or measured ______changing the atoms or molecules of a substance’s identity / · Can only be determined or measured as the substance ______into different substances
Examples:
Malleable: ability of a substance to mold into different shapes / Examples:
Intensive Property / Extensive Property
· Size of the sample DOES NOT matter
· A small piece & a large piece are the ______with respect to the property. / · Size of the sample DOES matter
· A big piece and a small piece would be ______with respect to this property
Examples: / Examples:
SELF CHECK: Determine if each property is Physical or Chemical [check the box]
Physical / Chemical / Physical / Chemicalflammability / malleability
boiling point / reactivity with oxygen (Combustion)
solubility
Physical and Chemical Changes
Physical Change / Chemical Change· The chemical structure of the substances IS NOT changed but it will look different
· Do not produce new substances
H2O(l)à H2O(g) / · Change in which the chemical structures of the substances ARE changed
· Does produce new substances
H2O(l)àH(g) + O2(g)
Examples: / Examples:
Ø Another name for a chemical change is called a ______
Possible Signs of a Chemical Change
1. / 4.2. / 5.
3.
Confusing Changes
TERM / DEFINITION / TYPE OF CHANGEMelting / Changes a solid to a liquid
Burning / Reacting with ______to produce CO2 and H2O
Dissolving / Adding 1 substance to another to form a ______mixture (solution)
Drying / Heating a sample to ______the water
Unit 3 [A] Section 3
Density
Do you want high or low density in your airbag?
® Density is defined as the ______of mass to volume of a sample.
How Heavy is it for its size:
LEAD = ______= small size is very heavy
AIR = ______= large sample has very little mass
Density Equation: /® Substances ______when they are less dense than the substance they are in. Using density values, Is water more or less dense than vegetable oil? ______
® Look at the density values to compare the various densities of substances. The larger the density, the more dense!
® Density does vary with ______.
Why? Most substances will expand when heated, increasing the volume and decreasing the density.
But ______is an exception. As water cools, it expands, increasing the volume and decreasing the density…….. Ice Floats on water
Calculating Volume using Water Displacement
v You can measure the volume of an object by water displacement. The volume is the difference between the ______and initial volume of the water after the object has been added to the water.
Example 1: What is the density of a sample with a mass of 2.50g and a volume of 1.7 ml?
Example 2: What is the mass of a 2.34 ml sample with a density of 2.78 g/ml?
Example 3: A sample is 45.4 g and has a density of .87 g/ml. What is the volume?
Self-Check:
Is it Aluminum? The metal has a mass of 612 g and a volume of 345 cm3? Aluminum’s density =2.70g/cm3
Graphing Density
Slope = ______so
Slope = Density / massvolume
Unit 3 Section 4 Part 1
Energy in Chemical & Physical Processes
Thermochemistry
v Study of changes that accompany chemical reactions and phase changes
v The Universe is considered to be made of 2 parts:
1. System: part that contains the reaction or process
2. Surroundings: everything else
Energy
v Energy is defined as the ability to do ______or transfer ______energy.
v There are 2 forms of energy. Chemical systems contain both Potential Energy and Kinetic Energy.
1. Potential Energy (PE): Energy at rest due to the ______of an object; chemical potential energy is the energy stored in a substance’s ______.
2. Kinetic energy (KE): Energy of the ______of particles in a substance and is ______proportional to temperature. As temperature increases, KE also ______.
v Law of Conservation of Energy states that energy is neither ______nor destroyed, just changed in form
C8H18 + O2 à H2O + CO2 + Energy
Stored PE converts to 25% work and 75% heat
Energy in chemical Reactions
HOT PACK
v An exothermic reaction is when the system ______energy; heat flows ______and the surroundings get ______. They have a ______DH
v H products < H reactants
4Fe + 3 O2 à 2 Fe2O3 + 1625 kJ or 4Fe + 3 O2 à 2 Fe2O3 DH = - 1625 kJ
COLD PACK
v An endothermic reaction is when the system ______energy; heat flows ______and the surroundings get ______. They have a ______DH
v H products > Hreactants
27kJ + NH4NO3(s)à NH4(aq)+1+NO3(aq)-1 or NH4NO3(s)à NH4(aq)+1 + NO3(aq)-1DH = + 27 kJ
What is the difference between “Heat” and “Temperature”?
Temperature / HeatInstrument used to measure this
Unit used to measure this
Definition / A measure of the average ______
of the molecules in a substance.
A measure of the
______of the molecules.
A measure of how or cold something is. / The total amount of energy in a substance. A form of ______that is transferred between objects because one is warmer than the other.
It depends on 3 things: ______
______
Units of Heat Energy
v A calorie is defined as the amount of heat needed to raise the temperature of 1 g of water by 1 °C
v 1 cal= 4.184 J
v Food “Calories” are kilocalories. 1kcal = ______calories.
Unit 3 [A] Section 4 Part 2
Calculating Heat
Specific Heat
· Amount of heat required to raise the ______of 1 g of a substance by 1 °C
· Different substances have different specific heats.
· Water has a specific heat of ______. Iron (Fe) has a specific heat of .449 J/g°C. Gold (Au) has a specific heat of .129 J/g°C.
· The higher the ______the more energy it takes to change its temperature.
Calculating Heat
Example:
A 155 g sample of an unknown substance was heated from 25.0 °C to 40.0 °C. The substance absorbed 5696 J of energy. What is the specific heat?
Example:
How much heat is needed to change the temperature of 12.0 g of silver with a specific heat of 0.057 cal/g°C from 25.0°C to 83.0 °C?
Unit 3 [A] Section 4 Part 3
Calorimetry: Measuring Heat (q)
® A coffee cup calorimeter measures heat at constant pressure; works on the premise that the amount of heat ______in a reaction or physical change is equal to the amount of heat ______by the water - q = + q
® Rearrange the specific heat equation: q = m x c x ∆T
Example:
A piece of unknown metal with mass 17.19 g is heated to an initial temperature of 92.50 °C and dropped into 25.00 g of water (with an initial temperature of 24.50 °C) in a calorimeter. The final temperature of the system is 30.05°C. What is the specific heat of the metal? Specific heat of water = 4.184 J/g° C
Example:
A 32.07 gram sample of vanadium was heated to 75.00 °C (its initial temperature). It was then dumped into a calorimeter. The initial temperature of the calorimeter’s water was 22.50 °C. After the metal was allowed to release all its heat to the calorimeter’s water, 26.30 °C was the final temperature. What mass of distilled water was in the calorimeter?
Specific heat of vanadium = .4886 J/g°C Specific heat of water = 4.184 J/g°C
Unit 3 [A] Section 4 Part 4
Measuring Heat during Phase Changes