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Thermochemistry

Gives off heat (emits)

exothermic

$1

Absorbs heat

endothermic

$1

It flows from hot to cold objects and is known by the letter “q”

heat

$1

The study of energy changes that occur during chemical reactions

thermochemistry

$1

What is a calorie or joule?

Unit of heat (q)

$1

Defined as the amount of heat needed to increase the temperature of an object by 1oC.

$2

Heat capacity

Defined as the amount of heat needed to increase 1 gram of an object by 1oC.

Specific heat or specific heat capacity

$1

The formula for q?

q = C x m x DT

$2

Solve the previous equation for the other three variables.

C = q/mDTm = q/CDTDT = q/Cm

$3

A balance bar has 200 Calories. How many kilojoules is this? How

many joules is this?

836.8 kj836,800 joules

$3

True or False. cal/oC is an acceptable unit for specific heat.

False, cal/goC

$1

True or False. Metals generally have a higher specific heat

capacity than water.

False

$1

True or False. Metals generally have a higher specific heat capacity

than molecular compounds.

False

$1

You measure 1200 joules of heat during a 30oC temperature change with a

substance that weighs 100 g. What is the specific heat of the substance.

0.40 J/goC

$3

During a phase change, the temperature of a substance ________.

Remains constant

$1

A calorimeter can effectively measure the heat of another

substance because of the ____________?

Law of Conservation of Energy

$2

DH of fusion involves which phase change?

Melting

$1

DH of solidification involves which more commonly known phase change?

Freezing

$1

If the percent mass of a solution weighing 300 g is 6%, what is the

mass of the solute?

.06 = x/300gx = 18 g

$2

The heat content of a system at a constant pressure is known as the

________ of that system.

enthalpy

$2

What is the enthalpy change in a chemical reaction known as?

Heat of reaction, DH

$2

If the DH of a reaction is negative then the reaction is _______ .

$1

exothermic

The heat of the reaction for the complete burning of one mole of a substance.

$1

Heat of combustion

The enthalpy change when a mole of solute is dissolved in a solvent.

$1

Heat of solution

True or False. The quantity of heat absorbed when a solid melts is the

same as the quantity released when the substance freezes.

True. DHfus = -DHsolid.

$1

What makes Hess’s law useful?

It allows you to determine heats of reactions indirectly

$2

The change in enthalpy that involves the formation of one mole of a compound from its

elements (at 25oC) is known as?

Standard heat of formation, (DHf

o)

$2

True or False. The standard heat of a reaction can be calculated by

the following equation …DHo = DHf

o(products) - DHf

o(reactants)

True

$1

How to solve a phase change problem

You just multiply DH of fusion/vaporization by mass

Water is vaporizing DHvap= 2260 J/g of H2Oq = m x DHvap

Water is melting DHfus = 334 J/g of H2Oq = m x DHfus

(J)

Heating/Cooling Curve: No Phase Changes

What happens at A, C, and E??

(J)

We use our old formula q = m x C x DT

q = m x Cliq x DT

q = m x Csolid x DT

q = m x Cgas x DT

Heating/Cooling Curve: Putting it all Together

(J)

You should be able to calculate the total heat going all the way from heating a substance from its solid to its gas

q = m x Cliq x DT

q = m x Csolid x DT

q = m x Cgas x DT

q = m x DHfus

q = m x DHvap

Sample Problem

You have a 4.30 grams of ice at -13.2oC. You heat it until it completely vaporizes. How much heat was needed to complete this process? Here are some numbers you might need.

(Cice = 2.10 J/goC)(Cwater = 4.18 J/goC) (Csteam = 1.70 J/goC)(DHfus = 334 J/g)(DHvap = 2260 J/g)

q = m x Cice x DT (4.30)(2.10)(0.00 - -13.2)= 119.2 J

q = m x DHfus (4.30)(334)= 1436.2 J q = m x Cwat x DT (4.30)(4.18)(100. - 0.00)= 1797.4 J

q = m x DHvap = (4.30)(2260)= 9718 J

To get answer you simplyadd these 4 numbers together:13071 J or 13.1 kJ (3 sig. figs)