The Flow of Energy

Think about this question for 15 seconds…What does a thermometer measure?Discuss with your neighbor what your answer isShare what your neighbor said

Suppose 2 identical candles are used to heat 2 samples of water. One sample is a cup of water, the other is 10 gallons in a large bucket.

1. How does the change in temperature of the samples compare?

2. How will the amount of heat received by each container compare?

Thermochemistry: the study of heat changes that occur during chemical reactions

Energy: the ability to do work or supply heat

energy is not considered to be matter- no mass and does not take up space

Potential Energy: stored energy, or energy of position-gas in a car has potential energy- it is stored, so that when you start the car, the car can use the energy of breaking bonds to make the car move

Kinetic Energy: energy of motionthe faster you move, the more kinetic energy you have- the potential energy of the gas is converted to kinetic energy when the car moves

Law of Conservation of Energy:Energy cannot be created or destroyed, but it can be transferred to another object, or changed from one form to another

Heat – (q) the flow of energy from one object to another due to temperature differencesTransferred from warmer object to cooler objectThe transfer occurs until an equilibrium is reached

When chemical reactions occur, there is either a release of heat or an absorption of heat

System: what is being studiedSurroundings: everything around the system in

natureUniverse: the system and the surroundings

together

It is the system that will gain energy from or lose energy to the surroundings

• 3 Types of Systems:1. Open- free exchange of matter and energy with

surroundings. - a saucepan of soup on a stove

2. Closed- no exchange of matter, but an exchange of energy with surroundings. - placing a lid on the saucepan so the soup

does not boil out3. Isolated- no exchange of matter or energy between

system and surroundings - placing soup in a thermos

• Determine the system and the surroundings in the following situations:

1. You are sitting around a campfire.

2. You place 2 solutions in a test tube.

3. A piece of hot copper is placed in a beaker of water.

Endothermic Process: when the system absorbs heat from the surroundings. q is a positive valueExample- when you sit around a campfire (surroundings) and you (system) start to become warm

Exothermic Process: system loses heat to the surroundingsq is negativeExample- After exercising, you perspire, you (system) is giving off heat to the surroundings

An ice cube is placed into a glass of room temperature lemonade in order to cool the temperature down.

1. Is the heat of the ice cube gained or lost?2. Is this endothermic or exothermic?3. Is q positive or negative?

Units of Energy:joule (J)- SI unit for heat or energy

may see kJ where 1kJ = 1000 Jcalorie(cal)- amount of energy needed to raise the temperature of 1g of pure water 1o CCalorie(Cal)- the nutritional unit

1Cal = 1kcal = 1000cal1J = 0.2390 cal4.184 J = 1 cal

Convert:1) 1656.70 J to cal2) 483.12 cal to J3) 0.56721 Cal to J

4) 395.951 cal5) 2021.4 J6) 2373.3 J

Heat Capacity: the amount of heat needed to change the temperature of a system 1o C

Specific Heat Capacity: the amount of heat needed to raise the temperature of 1g of a substance by 1o C

The higher the specific heat capacity, the more energy needed to raise the temperature

Metals tend to have a low specific heat because they heat up quickly

Water has a very high specific heat capacity- it takes a lot of energy to raise the temperature

Calculating the Specific Heat Capacity:

q = heat (J)m = mass (g)DT = temperature change Tfinal - Tinitial (o C)

The formula can be solved algebraically for an unknown

Do Problems 1 – 3 pg 299

Tm

qC

Measuring Heat Changes

• Calorimetry- the measurement of heat change in a chemical reaction

The heat released by the system is equal to the heat absorbed by the surroundings

For a system at constant pressure- the heat is the same as a property called Enthalpy (D H)So…

q = D H = C• m • D T

• When D H:-is positive- endothermic reaction- heat is absorbed into the system-is negative- exothermic reaction- heat is released from the system

When the reaction is in an aqueous environment, the mass, D T, and C of the water is used because the energy released or absorbed by the system is the same as the energy released or absorbed by the water

• Thermochemical Equations:-A balanced equation that includes the heat changes-Done at Standard Conditions:

physical states of the substances given at 25o C and pressure at 101.3kPa or 1 atmosphere

-From the balanced equation, you can determine the heat absorbed or released for a certain amount of substance

• Example:Baking soda is used in cooking to make cakes rise-2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g)

DH = 129 kJIf you start with 2.24 moles of NaHCO3, how much heat is needed to decompose?2.24molesNaHCO3 x 129 kJ

2 mol NaHCO3

= 144 kJ

• Heat of Combustion: DH for the complete combustion of 1 mole of a substance

Values are on Table 11.4 on page 305

Heat in Changes of State

• Heat of Fusion- DHfus- endothermic, melting

from a solid to a liquid• Heat of Solidification- DHsolid- exothermic,

liquid forms a solid

DHfus = - DHsolid

The heat absorbed as a substance melts is equal to the heat released as it freezes

• Heat of Vaporization- DHvap – endothermic,

heat absorbed as a substance vaporized from a liquid to a vapor• Heat of Condensation-DHcond – exothermic,

heat released as a substance condenses from a vapor to liquid

DHvap = -DHcond

• Heat of Solution- DHsoln – can be endothermic or exothermic, heat absorbed or released when a solute is dissolved into a solvent