___________ - substances involved in the chemical and physical changes under investigation, part of the universe we are

focusing on.(Ex. Your experiment – what is happening inside the beaker)

___________ - rest of the universe(outside the beaker)

___________ - system plus surroundings

___________ ___________ ___________ - set of conditions that describe and define the system

Ex. number of moles of each substance; temperature; pressure; physical states of each substance

___________ - properties of a system that depend only on the ___________ of the system

Normally capital letters

Heat is a state function: it is a variable of a system that is independent of the path it took to get there. Most

thermodynamic quantities are state functions.

Properties that depend only on values of state functions are also state functions

Ex: T, P, V

___________ (heat changes) H

___________ (order and disorder) S

___________ _________(thermodynamically favored or not) G

We use thermo to control rxnsEx. raw food + heat cooked

STP – the T in STP is 0°C and humans are not happy lab workers when it is that cold!

So, think of standard conditions as standard lab conditions which are 1 atm of pressure, 25°C /298 K (much more comfy!)

All of this information is shown by adding the symbol ° to G, H or S. So, so if you see ΔH°, then you automatically know the

pressure, temperature and conditions that apply!

___________ ___________ ___________ ___________ T = 298.15 KP = 1.00 atm

___________ ___________ ___________ pure substances in their liquid or solid phase

- the standard state for the pure liquid or solid will be zero for ΔH°and ΔG°

gases – the standard state is the gas at 1.00 atm of pressure will be zero for ΔH° and ΔG°

gaseous mixtures - partial pressure must be 1.00 atm

aqueous solutions - 1.00 M concentration

___________ ___________ ___________ ___________ __________ΔErxn = 0

Energy involved in a chemical rxn is neither created nor destroyed. Energy can be converted from one form to another but cannot be

created.

This means any change in energy of a system must be balanced by the transfer of energy either into or out of the system.

Also known as ___________ ___________ ___________

I. systems tend toward a state of ___________ ___________ ___________

( )

a. H O flows downhillb. objects fall when droppedc. E mghd. E mg h

2

potential

potential

=

=∆ ∆

II. systems tend toward a state of ___________ ___________

a. mirror shatters when droppedb. easy to scramble an eggc. food coloring disperses in water

What is energy? The ___________ to do work or transfer heat;

hot stuff which changes other stuff.

Energy causes changes in stuff. Examples:

___________ – the ability to do work or produce heat ; the sum of all potential and kinetic energy in a system is known as the internal energy of the system

___________ – stored energy P.E. = mghIn chemistry this is usually the energy stored in bonds When bonded atoms are separated, the PE is raised because energy must be

added to overcome the coulombic attraction between each nucleus and the shared electrons.

When atoms bond, the same coulombic attraction results in energy being released and a subsequently lower PE.

___________– energy of moving objects K.E. = ½ mv2

Energy depends on mass

___________– the amount of energy needed to break one mole of bonds in a covalent gas substance to form new gaseous products

In gas phase reactions ∆Ho values may be related to bond energies of all species in the reaction.

( ) ( ) ( )

( ) ( ) ( )

A - B bond energy A + B

H - Cl H + Clg g g

gkJ

mol g g

+ →

+ →432

∆H BE BE298o

reactants products= ∑ − ∑

Use the bond energies listed in Table B.14 estimate the heat of reaction for

1) CCl2F2 + F2 CF4 + Cl2

Use the bond energies listed to estimate the heat of reaction for

2) CH4 + O2 CO2 + H2O

To calculate the amount of heat gained or lost by an object

q = m × c × ∆T

q = heat (J)

m = mass (g)

c = specific heat J/g°C

∆T = Temp. final – Temp. initial, in °C

3) What is the heat absorbed by 177 g of water in order to raise it from 25 °C to 75 °C?

Specific heats for water: for steam = 2.03 J/goC for liquid = 4.184 J/goC for ice = 2.06 J/goC

The Law of Conservation of Energy also tells us that the amount of heat ___________ by the hot object must be equal to the amount of heat ___________ by the cold object.

20°C 10°C

15°C

Hot becomes cooler, and cold becomes warmer.

-qlost = qgained

- (mlost × clost × ∆Tlost ) = (mgained × cgained × ∆Tgained )

4) A piece of iron is heated to 135oC and dropped into 135 mLof water that was 52.0oC, the water’s final temperature was 77.0o C. What was the mass of the iron if the specific heat of iron is 0.449 J/gK and the specific heat of water is 4.18 J/gK.

+q: heat is added or absorbed – Endothermic–q: heat is lost or released – Exothermic

Phase changes occur when sufficient heat energy is added or removed

s ℓ melting (or fusion)ℓ s freezingℓ g boiling (or vaporization) g ℓ condensation (liquefaction if forced to occur by pressure)s g sublimationg s deposition

From this heat curve (for water) we can see how the temperatures change and the phases change. Notice that the temperature does not change during a phase change (the two flat lines).

q = mHfHf = heat of fusion, going between the solid and liquid phasesMelting (+) s ℓ freezing (-) ℓ s

q = mHvHv = heat of vaporization, going between the liquid and gas phasesvaporizing (+) ℓ g condensing (-) g ℓ

q = mHsHs = heat of sublimation, going between the solid and gas phases

5) Calculate the heat needed for the heat curve shown below for 277 g of water cooled from 140. °C to 40.0 °C

Hfwater = 334 J/g

Hvwater = 2260 J/g

c ice = 2.06 J/g °C

c water = 4.18 J/g °C

c steam = 2.03 J/g °C

___________ ___________– show amount of activation energy needed for the rxn to occur.

___________reactions release specific amounts of heat as products

Potential energies of products are ___________potential energies of reactants.

___________reactions absorb specific amounts of heat in order to make products

Potential energies of products are ___________potential energies of reactants.

Chemistry is done at constant pressureopen beakers on a desk top are at atmospheric pressure

___________ - enthalpy changechange in heat content at constant pressure

∆H = qp

___________ - heat of reaction – unit kJ/molrxn∆Hrxn = Hproducts - Hreactants

∆Hrxn = Hsubstances produced - Hsubstances consumed

Change in enthalpy, ∆H, or heat of reaction is amount of heat absorbed or released when a reaction occurs at constant pressure.

When:∆H is > 0, (positive)the reaction is ___________

(heat is a reactant) heat energy is added into the system

∆H is < 0, (negative)the reaction is ___________(heat is a product) heat energy is lost from the system

Standard molar enthalpy of formationsymbol is ∆Hf

o

unit is kJ/molrxn

defined as the enthalpy for the reaction in which one mole of a substance is formed from its constituent elements

for example:

( ) ( ) ( )

( )

Mg Cl MgCl kJ

H kJ / mol

s 2 g 2 s

f MgClo

2 s

+ → +

= −

6418

6418

.

.∆

Standard molar enthalpies of formation have been determined for many substances and are tabulated in Table B.13 in

Appendix B in the text.

Standard molar enthalpies of elements in their most stable forms at 298.15 K and 1.000 atm are zero.

The standard molar enthalpy of formation for phosphoric acid is -1281 kJ/molrxn . Write the equation for the reaction for

which ∆Horxn = -1281 kJ. P in standard state is P4

( ) ( ) ( ) ( )

( )

32

142 1281

1281

H O P H PO kJ

H kJ / mol

2 g 2 g 4 s 3 4 s

f H PO o

3 4 s

+ + → +

= −∆

___________ ___________ ___________– the overall enthalpy change in a rxn is equal to the sum of the

enthalpy changes for the individual steps in the process.

enthalpy change for a reaction is the same whether it occurs by one step or by any (hypothetical) series of steps ~ true because

∆H is a state function

6) Using Hess’s law find ΔHrxn for 2HCl(g) + F2(g) 2HF(l) + Cl2(g)

from the following rxns

ΔH4HCl(g) + O2(g) 2H2O(l) + 2Cl2(g) -202.4 kJ/molrxn

½ H2(g) + ½ F2(g) HF(l) -600.0 kJ/molrxn

H2(g) + ½ O2(g) H2O(l) -285.8 kJ/molrxn

7) ΔH for 4FeO(s) + O2(g) 2Fe2O3(s) is -560 kJ/mol. Use the ΔH for the following two rxns to verify.

ΔH in kJ/molrxn

2Fe(s) + O2(g) 2FeO -544 4Fe(s) + 3O2(g) 2Fe2O3(s) -1648

___________ ___________ in a more useful formany chemical reaction at standard conditions, the standard enthalpy change is the sum of the standard molar enthalpies of formation of the products (each multiplied by its coefficient in the balanced chemical equation) minus the corresponding sum for the reactants

∆ ∆ ∆H H H298o

f productso

f reactantso= ∑ − ∑n n

n n

___________ ___________ ___________ - The universe tends toward a state of greater disorder and low energy in a

thermodynamically favored (spontaneous) reaction…exothermic rxns tend to lower energy states and

order requires much more energy than disorder.

___________ ~ continues on its own after given activation energy to start the rxn. Happens without any continuing outside

influences.

Spontaneous processes require:~ free energy change of system must be negative~ entropy (disorder) of universe must increase

rusting of iron - thermodynamically favored so it ___________ ___________

Have you ever seen rust turn into iron metal without help?

melting of ice at room temperature - thermodynamically favored, ___________ ___________

Will water spontaneously freeze at room temperature?

* Exothermicity does not ensure spontaneity___________ ___________

exothermic and thermodynamically favored (spontaneous) only below 0oC

* Increase in disorder of the system also does not ensure spontaneity

Entropy is a measure of the ___________ ___________ of a system. ___________ ___________ ___________.

When:∆S is ___________ disorder increases (favors spontaneity)

∆S is ___________ disorder decreases (disfavors spontaneity)

symbol is ∆So

unit is J/molrxn K

Notice that entropy is temp based and it is in J not kJ

From 2nd Law of Thermodynamics, for a spontaneous process

Gases are more disordered than solutionsSolutions are more disordered than pure liquids

Pure liquids are more disordered than solids

In general for increase in entropy:

∆ ∆ ∆S S Suniverse system surroundings= + > 0

Determine whether the entropy value increases or decreases for the following situations

Melting of ice at room temperature

Determine whether the entropy value increases or decreases for the following situations

Condensation of water vapor on a window

Determine whether the entropy value increases or decreases for the following situations

An iron rusting4Fe(s) + 3O2(g) 2Fe2O3(s)

Determine whether the entropy value increases or decreases for the following situations

Baking soda reacting with vinegarNaHCO3(s) + CH3COOH(l) NaCH3COO(aq) + H2O(l) +

CO2(g)

Which of the following rxns creates the most disorder?

a) H3BO3(aq) → HBO2(aq) + H2O(l)

b) 4 HBO2(aq) → H2B4O7(aq) + H2O(l)

c) H2B4O7(aq) → 2 B2O3(s) + H2O(l)

___________ ___________ ___________ ___________ The entropy of a hypothetical pure, perfect, crystalline substance at

absolute zero temperature is zero.

allows us to measure absolute values of entropy for substancescool them down to O K, or as close as possible, then measure

entropy increase as substance warms up

Entropy changes for reactions can be determined similarly to ∆H for reactions. As with ∆H, entropies have been measured and

tabulated in Table B.16 in Appendix B as So298. When:

∆S n S n S298o

productso

reactantso= ∑ − ∑

J. Willard Gibbs determined the relationship of enthalpy and entropy that best describes the maximum useful energy which can be obtained in the form of work from a process at constant

Temperature & Pressure.

The relationship also describes the ___________ ___________. Whether the reaction is ___________ ___________ or

___________ ___________ ___________

∆ ∆ ∆G = n G n G298o

productso

reactantso∑ − ∑

The change in the Gibbs Free Energy is a reliable indicator of spontaneity of a physical process or chemical reaction.

___________ tell us the speed of the process kinetics deals with rates of rxns –coming soon…

When:∆G is > 0 reaction is not thermodynamically favored

(___________ ___________) reactant favored

∆G is = 0 system is at equilibrium (more later)

∆G is < 0 reaction is thermodynamically favored (___________) product favored

Changes in free energy obey the same type of relationship we have described for enthalpy and entropy changes.

Enthalpy and entropy can sometimes ___________ each other –this makes the reaction really go or really not go.

Ex. Dynamite has a neg Δ H & a pos Δ S so the rxn really goes once started…

If the signs don’t reinforce does a rxn occur? This is where Gibbs Free energy addresses the spontaneity of rxns.

Ex. Liquid water to water vapor Δ H is + and Δ S is +

Δ G is -, rxn is thermodynamically favored (spontaneous) –reaction will go on its own once started to make the products.

Ex. ___________.

Δ G is +, rxn is not thermodynamically favored (nonspontaneous) – rxn won’t go on it’s own, wants to stay as reactants.

Ex. ___________

8) Find the following for the rxn belowa) enthalpy, is the rxn exothermic or endothermic b) entropy, is the rxn more or less ordered c) Gibbs free energy, is the rxn thermodynamically favored

(spontaneous/product favored), or not thermodynamically favored, (nonspontaneous/reactant favored)?

Use data from Appendix B, table B.16 C2H5OH(l) + 3O2(g) 2CO2(g) + 3H2O(g)

Sometimes the ΔHfo for a compound will be missing

9) Given that the ΔHfo for O2 = 0; ΔHf

o for SO2 = -296.8 kJ/molerxn ; ΔHf

o for H2O = -285.8 kJ/molerxn and the ΔHrxnfor the following balanced equation is = -1124 kJ/molerxn , what is the Δ Hf

o for H2S? 2H2S + 3O2 2SO2 + 2H2O

10) Upon adding solid potassium hydroxide pellets to water the following reaction takes place:

KOH(s) → KOH(aq) + 43 kJ/mol Answer the following questions regarding the addition of

14.0 g of KOH to water: a) Does the beaker get warmer or colder? b) Is the reaction endothermic or exothermic? c) What is the enthalpy change for the dissolution of the

14.0 grams of KOH?

11) Calculate the enthalpy change for the reaction in which 15.0 g of aluminum reacts with oxygen to form Al2O3 at 25oC and one atmosphere. ΔH rxn is -3352 kJ/molerxn (hint: use stoichiometry)

The relationship between ∆G0 and E0cell is a simple one.

e ofnumber ne mol J/V 96,500 F where

E F-n G

-

-

0cell

0

=

=

=∆

53

For Electrolytic cells:(needs a battery/electricity to occur)

Eocell is negative – does not happen naturally

so ∆G is positive – not thermodynamically favored (___________ ___________)

For Voltaic (Galvanic) cells:(produces voltages)

Eocell is positive - occurs as expected

so ∆G is negative - thermodynamically favored (___________ ___________)

12) Using the table of standard reduction potentials, calculate ΔG° for the following reaction. Explain whether or not this reaction is thermodynamically favorable.

Cu2+ (aq) + Fe (s) → Cu(s) + Fe2+

(aq)

1st Calculate E0cell using the appropriate half-reactions.

2nd – After you have E0cell , you can calculate ∆G0

56

The general relationship of ∆G, ∆H, and ∆S is (at constant pressure)

Which gives us 4 possibilities among the signs

∆H ∆S ∆G Therefore- + - forward rxn spontaneous ___________- - ? forward rxn spontaneous ___________+ + ? forward rxn spontaneous ___________+ - + forward rxn nonspontaneous ___________

∆ ∆ ∆G = H - T S

13) Calculate ΔS298 for the combustion of propane. Δ H298= -2219.9 kJ, and ΔG298= -2108.5 kJ.

State function + -Δ H Endothermic

(taking in heat)ice melting

Exothermic (giving off heat)

dynamite Δ S Towards disorder Towards orderΔ G not thermodynamically

favoredthermodynamically

favored

∆H ∆S ∆G Therefore- + - forward rxn spontaneous at all T’s- - ? forward rxn spontaneous at low T’s+ + ? forward rxn spontaneous at high T’s+ - + forward rxn nonspontaneous at all T’s

If Eocell is negative – ∆G is positive

If Eocell is positive – ∆G is negative

______________________ - used to measure the amount of heat produced (or absorbed) in a reaction at constant Pmeasures qP

___________ ___________ - heat evolved by reaction is determined from the temperature rise of the solution

Amount of heat gained by calorimeter is the heat capacity of the calorimeter or ___________ ___________

value determined by adding a specific amount of heat to calorimeter and measuring T rise

Amount of heatreleased by reaction

Amount of heatgained by calorimeter

Amount of heatgained by solution

=

+

14) When 3.425 kJ of heat is added to a calorimeter containing 50.00 g of water the temperature rises from 24.000oC to 36.540oC. Calculate the heat capacity of the calorimeter in J/oC. The specific heat of water is 4.184 J/goC.

heat capacity: h.c. = qcal

∆Tqrxn = qcal + qsoln (in this case the soln is just water)

15) A coffee-cup calorimeter is used to determine the heat of reaction for the acid-base neutralization of acetic acid and sodium hydroxide. When we add 25.00 mL of 0.500 M NaOH at 23.0000C to 25.00 mL of 0.600 MCH3COOH already in the calorimeter at the same temperature, the resulting temperature is observed to be 25.9470C. The heat capacity of the calorimeter had previously been determined to be 27.8 J/0C. Assume that the specific heat of the mixture is the same as that of water, 4.18 J/g0C and that the density of the mixture is 1.02 g/mL.

A) Calculate the amount of heat given off in the reaction. B) Determine ΔH for the reaction under the conditions of

the experiment. (must determine the number of moles of reactants consumed; use limiting reactant)