Thermochem.pdf

Thermochemistry

Law of Conservation of Energy

 Energy is neither created nor destroyed it

What is Heat?

 Form of Kinetic Energy

 Moves from things with high temperature

to things with a low temperature

 What units are used to measure it?

Units for Measuring Heat

The Joule is the SI system unit for measuring heat:

The calorie is the heat required to raise the temperature of 1 gram of water by 1 Celsius degree 2 2

1

1

1

s

m

kg

meter

newton

Joule









Joules

calorie

4

.

18

Internal Energy vs. Heat (Q)

 Internal Energy - KE in matter causing

its particles to move around.

 Heat - KE that moves between pieces of

matter.

 As heat flows out of an object, what

happens to its internal energy?

 As heat flows to an object, what happens

Thermal Equilibrium

 When will heat stop flowing?

Temperature = Avg. KE

 In a glass of hot water all molecules move

at different speeds.

 Average kinetic energy of the

Temperature ain’t Heat

 Hot cup of coffee vs. Big mountain lake  Which one has more internal energy?

 If I put the cup of hot coffee in the lake,

which way does the heat go?

Heat Capacity

 Heat capacity - Amount of heat energy

required to raise the temperature of a portion of matter by 1o_C.

 What has a greater heat capacity? 500 kg

Specific Heat Capacity

 Specific Heat Capacity - amount of heat

energy required to raise one gram of a certain substance by one 1o_C.

 Which has a greater Specific Heat

Q = mC

D

_T

 Q is heat energy flowing into or out of a

system. (Endothermic vs. Exothermic)

 m = mass

 C = specific heat capacity

Specific Heat Formula

 C = Q/ m X DT

 Specific Heat = Heat / [mass X change in

A problem

 Temperature of a 95.4 g piece of copper

increases from 25o_{C to 48}o_{C when the}

Two Trends in Nature



Entropy - Order



Disorder

 



Thermodynamics –

High energy



Low energy

System v. Surroundings



System

- actual chemical reaction

taking place



Surroundings

- rest of the universe

Heat Flows



Heat always flows from hot to

cold

◦

Heat flows (not cold)

Exothermic process - process that gives off heat –

transfers thermal energy from the system to the surroundings.

2H₂ (g) + O₂ (g) 2H₂O (l) + energy

H₂O (g) H₂O (l) + energy

Exothermic Processes

Endothermic process

- process in which

heat has to be supplied to the system from

the surroundings.

energy + 2HgO (s) 2Hg (l) + O₂ (g)

energy + H₂O (s) H₂O (l)

Endothermic Processes

Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.

DH = H (products) – H (reactants)

DH = heat given off or absorbed during a reaction at constant pressure

H_products < H_reactants

DH < 0

H_products > H_reactants

Thermochemical Equations

H₂O (s) H₂O (l) DH = 6.01 kJ

Is DH negative or positive?

System absorbs heat

Endothermic

DH > 0

6.01 kJ are absorbed for every 1 mole of ice that melts at 00_{C and 1} atm.

Thermochemical Equations

CH₄ (g) + 2O₂ (g) CO₂ (g) + 2H₂O (l) DH = -890.4 kJ

Is DH negative or positive?

System gives off heat

Exothermic

DH < 0

890.4 kJ are released for every 1 mole of methane that is combusted at 250_{C and 1 atm.}

Calorimetry

The amount of heat absorbed or released during a

Calorimetry

 When measuring heat lost or gained by

the reaction, we measure it using water as the basis for the calculations.

Q

_w

=

m x C x

D

t

m of water Q of water

C of water

Calorimetry

 Heat gained or lost by reaction is equal to, but opposite in sign, to the heat gained or lost by the water

-q_water = q_rxn

Sample problem

 When 25.0 mL of water containing 0.025 mol

HCL at 25o_{C is added to 25.0 mL of water} containing 0.025 mol NaOH in a foam cup calorimeter, a reaction occurs. Calculate the enthalpy change during this reaction if the

Sample problem

What do we know?

C _water = 4.18 J/(go_C)

V _final = 25.0 mL HCl+ 25.0 mL NaOH= 50 mL T_i = 25.0 o_C

T_f = 32.0 o_C

Density solution = 1.00 g/ml

Sample problem

Solution

1. Calculate the mass of water

m = 50 mL x 1.00g/mL = 50.0 g

2. Calculate DT

32.0 – 25.0 = 7.0 o_{C (Heat gained by water}

is lost by reaction)

Endothermic or Exothermic?

 Water absorbed heat : heat gained by

water is lost in the reaction

This was an exothermic reaction

Heat (Enthalpy) Change,

ΔH

Definition: The amount of heat energy released or absorbed during a process.

Energy and Heat Review

Energy is the capacity to do work, and can take many forms

 Potential energy is stored energy or the energy of position

Heat in Change of States

 Heat energy is involved in chemical

reactions, but it’s also involved in physical changes as well.

H₂O _(s) + DH  H₂O _(l)

Heat of Fusion

 When a substance begins to melt, which

direction is heat energy going?

 What type of process is this?

 Heat needed to melt the solid substance is

called Heat of Fusion

 Molar Heat of Fusion is the heat to melt

one mole of substance

Heat of Solidification

 When a substance freezes, which direction

is the heat energy going?

 What type of process is this?

 The heat you need to draw out of a

substance to freeze is the Heat of Solidification

 Molar Heat of Solidification is heat to

freeze one mole of substance

Heat of Vaporization and

Condensation

 Heat required to vaporize 1 mole of a

substance is the Molar Heat of Vaporization

 Heat you must remove from 1 mole of a

substance to condense it is the Molar Heat of Condensation

Water phase changes

Phase Change

Diagram

Processes occur by addition of energy 

Phase Diagram

•

Represents phases as a function of temperature and pressure.

•

Critical temperature: temperature above which the vapor can not be liquefied.

•

Critical pressure: pressure required to liquefy AT the critical temperature.