Unit 3 [A] Section 1A

States of Matter

Properties / Solid / Liquid / Gas
Packing:
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

  • Kinetic Energy must be put INTO the substance in orderto increasethe ______of the molecules so as to break the ______forces holding the particles together

Melting: change of state from a solid to a ______

Vaporization (Boiling or Evaporation): change of state from a ______to a gas

Sublimation: direct change of state from a ______to a gas

Exothermic physical changes of state

  • 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

Freezing: change of state from a ______to a solid

Condensation: change of state from a gas to a ______

Deposition: direct change of state from a ______to a solid

Temperature of State Changes

Freezing point (fp) is the temperature at which the liquid turns into a solid

Melting point (mp) is the temperature at which the solid turns into a liquid

Freezing Point = Melting Point

Example: Water has a mp and a fp of 0C

Boiling point (bp)is the temperature at which the liquid turns into a gas

Condensation point (cp) is the temperature at which a gas turns into a liquid

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 Pressure
Force 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 boilingpoint 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

When external atmospheric pressure = vapor pressure of a liquid

At 90C, the water’s vapor pressure is not strong enough to push against the atmospheric pressure  starting to boil

At 100C, the water’s vapor pressure is equal to the atmospheric pressure boiling is occurring

Since atmospheric pressure changes at various altitudes, we use “______” boiling point to describe the temp at which a LG at 1 atm

Important Ideas

The higher in altitude the ______the atmospheric pressure

At higher altitudes, the boiling point is ______

It takes ______to cook foods at higher altitudes (lower atmospheric pressures)

Unit 3 [A] Section 1 C

Heating & Cooling Curves

A diagram that shows how solids, liquids & gases change state when ______is changed

Plateaus = the changes of state (freezing, melting, etc.)

Freezing Point & Melting Point are at the temperature or at the same plateau

Boiling Point &Condensation Point are at the same temperature or at the same plateau

Slopes= pure states (solid, liquid or gas)

At the plateaus, ______energy remains constant because temperature remains constant while potential energy changes

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

  1. What is the freezing point of this substance?
  1. How long does it take for the gas to completely liquefy?

Phase Diagram

A diagram that shows how solids, liquids & gases change state as both temperature and ______are changed

Crossing a line between states determines the change of state (boiling, melting, etc)

A point directly on a line will identify the pressure and temperature (boiling point, melting point, etc.) of the phase change

______is the temperature and pressure in which all 3 of the states coexist

______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 / Chemical
flammability / 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) H2O(g) /
  • Change in which the chemical structures of the substances ARE changed
  • Does produce new substances
H2O(l)H(g) + O2(g)
Examples: / Examples:

Another name for a chemical change is called a ______

Possible Signs of a Chemical Change

4.
5.

Confusing Changes

TERM / DEFINITION / TYPE OF CHANGE
Melting / 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

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.

Example1: 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 / mass
volume

Unit 3 Section 4 Part 1

Energy in Chemical & Physical Processes

Thermochemistry

Study of changes that accompany chemical reactions and phase changes

The Universe is considered to be made of 2 parts:

  1. System: part that contains the reaction or process
  2. Surroundings: everything else

Energy

Energy is defined as the ability to do ______or transfer ______energy.

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 ______.
  1. Kinetic energy (KE): Energy of the ______of particles in a substance and is ______proportional to temperature. As temperature increases, KE also ______.

Law of Conservation of Energy states that energy is neither ______nor destroyed, just changed in form

C8H18 + O2  H2O + CO2 + Energy

Stored PE converts to 25% work and 75% heat

Energy in chemical Reactions

HOT PACK

An exothermic reaction is when the system ______energy; heat flows ______and the surroundings get ______. They have a ______H

H products < Hreactants

4Fe + 3 O2  2 Fe2O3 + 1625 kJ or 4Fe + 3 O2  2 Fe2O3 H = - 1625 kJ

COLD PACK

An endothermic reaction is when the system ______energy; heat flows ______and the surroundings get ______. They have a ______H

H products > Hreactants

27kJ + NH4NO3(s)NH4(aq)+1+NO3(aq)-1 or NH4NO3(s) NH4(aq)+1+ NO3(aq)-1H = + 27 kJ

What is the difference between “Heat” and “Temperature”?

Temperature / Heat
Instrument 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

A calorie is defined as the amount of heat needed to raise the temperature of 1 g of water by 1 C

1 cal= 4.184 J

 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

Heat of Fusion/Solidification

Latent Heat of fusion (Hfus ) is the heat energy required to melt one gram of a solid at its melting point

q = Hfus x mass

For water, Hfus = 334 J/gOn reference sheet

Latent Heat of solidification (Hsolid ) is the heat energy lost when one gram of a liquid freezes to a solid at its freezing point

q = Hsolid x mass

 For water, Hsolid = -334 J/g

Heat of Vaporization/Condensation

Latent Heat of vaporization (Hvap) is the heat to vaporize one gram of a liquid at its normal boiling point

q = Hvapx mass

For water, Hvap= 2260 J/g On reference sheet

Latent Heat of condensation (Hcond ) is the heat energy released when one gram of a liquid forms from its vapor

 For water, Hcond= -2260 J/g

q = Hcondx mass

Examples

  1. How much heat is needed to melt 500.0g of ice at 0 C?
  1. How much heat is evolved when 1255 g of water condenses to a liquid at 100°C?