CHAPTER ELEVEN&TWELVE

UNIT A: THERMODYNAMICS

Energy Demands & Resources

Personal Use of Chemical Energy Food (energy from photosynthesis)

Access through cellular respiration

Fossil Fuels (combustion reactions) Heat home, cook, transportation

Alternative forms: try to be more eco-friendly Solar Wind Water

THE LAW OF CONSERVATION OF ENERGYTHE LAW OF CONSERVATION OF ENERGY

During physical and chemical processes, energy may change form, but it may never be created nor destroyed.

If a chemical system gains energy, the surroundings lose energy

If a chemical system loses energy, the surroundings gain energy

Examples: When octane (C3H8, the main

component of gasoline) is burned in your car engine, chemical bond energy (potential energy) is converted into mechanical energy (pistons moving in the car engine; kinetic energy) and heat.

When we turn on a light switch, electrical energy is converted into light energy and, you guessed it, heat energy.

DO YOU REMEMBER??

EXOTHERMICEXOTHERMIC ENDOTHERMICENDOTHERMIC

A change in a chemical energy where energy/heat EXITS the chemical system

Results in a decrease in chemical potential energy

A change in chemical energy where energy/heat ENTERS the chemical system

Results in an increase in chemical potential energy

DO YOU REMEMBER??

SYSTEMSSYSTEMS Types of SystemsTypes of Systems

System: the part of the universe we are interested in studying

Surroundings: the rest of the universe

Open: exchange matter & energy

Closed: doesn’t exchange matter, but does energy

Isolated: doesn’t exchange matter or energy

DO YOU REMEMBER??

open closed isolated

An Introduction to Thermodynamics

Kinetic Energy (Ek) is related to the motion of an entity

Molecular motion can by translational (straight-line), rotational and vibrational

Chemical Potential Energy (Ep) is energy stored in the bonds of a substance and relative intermolecular forces

Thermal Energy is the total kinetic energy of all of the particles of a system. Increases with temperature.

Symbol (Q), Units (J), Formula used (Q=mcΔT)

Temperature is a measure of the average kinetic energy of the particles in a system

Heat is a transfer of thermal energy. Heat is not possessed by a system. Heat is energy flowing between systems.

An Introduction to Thermodynamics

Which has more thermal energy, a hot cup of coffee or an iceberg?

An iceberg!Put very roughly, thermal energy is related to the amount of something you have multiplied by the temperature.

Let's assume your iceberg is at the freezing point of water - 0 degrees Celsius (~273 Kelvin). Now your cup of coffee might be 75 degrees Celsius (~350 Kelvin).

350 isn't a whole lot more than 270, but an iceberg is thousands of times larger than a cup of coffee. Even though the iceberg is at a lower temperature, it contains more thermal energy because the particles are moving and it's much larger than the cup of coffee.

An Introduction to Thermodynamics

Heat energy transferred will be related to the temperature change of the system.

It takes different amounts of heat energy to raise the temperature of 1 g of different substances by 1 oC.

This number is called the specific heat capacity (c), and is measured in units of J

g C

Thermal Energy Calculations

There are three factors that affect thermal energy

Q = mcΔtMass (m) – unit is g or kg

Type of substance (c) – unit is J/goC or kJ/kgoC

Temperature change (Δt)

Water’s Specific Heat Capacity

Water has a c value of

(can also use 4.19kJ/kgoC) This means that it takes 4.19 J of heat to raise the temperature of 1 g of water by 1oC

Water has a very large c compared to most other common substances.

Consider a bathtub and a teacup of water! They both contain water with the same heat capacity, but it would take much longer to heat up the water in the bathtub!

4.19 J g C

Thermal Energy Calculations

After coming in from outside, a student makes a cup of instant hot chocolate by heating water in a microwave. What is the gain in thermal energy if a cup (250 mL) of tap water is increased in temperature from 15oC to 95oC? (1mL of water = 1g)

Thermal Energy Calculations

A backpacker uses an uncovered pot to heat lake water on a single-burner stove. If the water temperature rises from 5.0oC to 97oC, and the water gains 385 kJ of thermal energy, what is the mass of water heated?

Thermal Energy Calculations #2

A 750 g sample of ethanol absorbs 23.4 kJ of energy. Its initial temperature was 35oC. The specific heat capacity of ethanol is 2.44 J/goC . What is the final temperature of the ethanol sample?

Thermal Energy Calculations

Experiments have shown that a thermal energy change is affected by mass, specific heat capacity, and change in temperature. What happens to the thermal energy if

a) The mass is doubled?b) The specific heat capacity is divided by two?c) The change in temperature is tripled?d) All three variables are doubled?

Thermal Energy Calculations

If the same quantity of energy were added to individual 100g samples of water, aluminum, and iron, which substance would undergo the greatest temperature change? Explain. (see page 3 of data booklet)

How do we measure Q?

With a simple laboratory calorimeter, which consists of an insulated container made of three nested polystyrene cups, a measured quantity of water, and a thermometer.

The chemical is placed in or dissolved in the water of the calorimeter.

Energy transfers between the chemical system and the surrounding water is monitored by measuring changes in the water temperature.

“Calorimetry is the technological process of measuring energy changes of an isolated system called a calorimeter”

Includes: Thermometer, stirring rod, stopper or inverted cup, two Styrofoam cups nested

together containing reactants in solution

Comparing Q’s

Negative Q value

An exothermic change

Heat is lost by the system

The temperature of the surroundings increases and the temperature of the system decreases

Example: Hot Pack

Question Tips: “How much energy is released?”

Positive Q value

An endothermic change

Heat is gained by the system

The temperature of the system increases and the temperature of the surroundings decreases

Example: Cold Pack

Question Tips: “What heat is required?”

Other Calorimetry assumptions. . .

• All the energy lost or gained by the chemical system is gained or lost (respectively) by the calorimeter; that is, the total system is isolated.

• All the material of the system is conserved; that is, the total system is isolated.

• The specific heat capacity of water over the temperature range is 4.19 J/(g•°C). (** IN YOUR DATA BOOK)

• The specific heat capacity of dilute aqueous solutions is 4.19 J/(g•°C).

• The thermal energy gained or lost by the rest of the calorimeter (other than water) is negligible; that is, the container, lid, thermometer, and stirrer do not gain or lose thermal energy.

Calorimetry Questions

•Remember to keep the system and the surroundings separate in your calculations.

Because we assume the system is isolated, then we assume that:

1. the energy lost by the system = energy gained by the surroundings (water)

OR

2. the energy gained by the system = the energy lost by the surroundings.

Calorimetry Questions

A 15.0 g piece of copper at 100oC is placed in a calorimeter with 100.0mL of water at 25.00oC. The final temperature of the water and copper is 25.35oC. What is the specific heat capacity of copper?

Calorimetry Questions

Suppose a piece of iron with a mass of 21.5 g at a temperature of 100.0oC is dropped into an insulated container of water. The mass of the water is 132.0 g and its temperature before adding the iron is 20.0oC. What will be the final temperature of the system? (ciron = 0.449 J/goC)

Calorimetry Questions

A student collected the following data from a calorimetry lab:

What is the specific heat capacity of the metal?

Mass of empty cup 2.31 g

Mass of cup + water 180.89 g

Mass of cup + water + metal

780.89 g

Initial temperature of water

17.0oC

Initial temperatuer of metal

52.0oC

Final temperature of system

27.0oC

Calorimetry Questions

Suppose we mixed 40.0 g of hot water at 60oC with 100.0 g of cold water at 30.0oC. What is the final temperature?