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