Thermochemistry

Definition: The study of the ________ __________ associated with

a physical and/or chemical ________ of a _________. Thermal Energy: arises from a particle’s __________ and _______. Systems: Chemical - _____________________________________________

_____________________________________________

- _____________________________________________

_____________________________________________ Surroundings - ___________________________________________

___________________________________________ Open - __________________________________________________

__________________________________________________ Closed - ________________________________________________

________________________________________________ Isolated - ________________________________________________

________________________________________________ Others: Heat (q) - _______________________________________________

_______________________________________________ Exothermic - _____________________________________________

_____________________________________________ Endothermic - ____________________________________________

____________________________________________ Temperature - ____________________________________________

____________________________________________

Calculation: q = mc∆T where m = ______________________________

c = ___________ _______ ____________

______________________________

∆T = __________________ ___________

= _____________________________

eg. Calculate the heat generated when 50.0 g of water (cH2O = 4.18

J/g•ºC) changes temperature from 12.5ºC to 47.0ºC? Solution: q = mc∆T

eg. Calculate the heat generated when 50.0 g of water and 125 g of

the Al can (cAl = 0.900 J/g•ºC) changes temperature from 12.5ºC

to 47.0ºC? Solution: q = mc∆T

= mH2O x cH2O x ∆T + mAl x cAl x ∆T

Heat Transfer and Energy Change

Chemical system energies, PE & KE, are sources of Enthalpy, H - interactions between ____________ and ______________

- electron __________

- vibration of _______ within ______________

- ____________ and _________________ of molecules

- nuclear _____________ energies

- electronic PE of atoms connected by chemical _________ and

molecules connected by _________ of attraction Enthalpy Change, ∆H Results in the ______________ in ____________ of the reactants and

the products during a ___________, due to:

- breakage of ______ or intermolecular ________ ∆H = ________

- formation of ______ or intermolecular ________ ∆H = ________ Determination of Enthalpy Change, ∆H - H and ∆H ___________ be directly _______________

- measured ___________ to the energy ____________________,

heat (q), during the _________ in the system

∆Hsys = - qsurr

= - mc∆Tsurr

eg. Calculate the enthalpy change for a combustion reaction which

results in the temperature of 500. mL of water in a 0.250 kg Al

(cAl = 0.900 J/g•°C) tea kettle rising from 20.0 °C to 95.0 °C. VH2O = mH2O as DH2O = 1.0 g/mL

∆H = - (mc∆TH2O + mc∆TAl)

= - (mcH2O + mcAl) ∆T

=

This is an ______________ reaction showing the change in _____ of

the ____________ is equal to the change in ______ of the

_______________________.

PE is that of _____________ or the __________ to work.

When objects that __________ each other:

Get _________ together, PE __

Move ___________ apart, PE __ KE is that of motion.

average KE __ as T __

Chemical change involves __ ______________ that alter the ___ of

atoms. Bond breaking ____________ energy and PE of the atoms ___. Bond forming ____________ energy and PE of the atoms ___. Depending on which _________ is _____________, the _____

change in ____ energy may be an ___________ or a ___________. The ___________________________________________________,

the __________________ the bond.

Molar Enthalpies

Molar Enthalpy, ∆Hrxn(X)

• the ∆H for ________ of any substance, (X), undergoing a

reaction.

• standard unit is ___________ eg. rxn solution _________ combustion _________

formation ________ neutralization _________

Determination:

H (x) = rxn

eg. a) Calculate the ∆Hcomb (C8H18) if 6.593 g causes 1.000 L of

H2O to increase its temperature by 75.2 °C.

( DH2O = 1 g/mL or 1kg/L)

mH2O =

Hcomb (C8H18) = -mH2OcH2OT

mC8H18

x MC8H18

b) Express this enthalpy change 4 ways:

i) ∆Hcomb(C8H18) =

ii) C8H18 + O2 CO2 + H2O +

OR C8H18 + O2 CO2 + H2O +

iii) C8H18 + O2 CO2 + H2O ; H =

iv) PE diagram

For comparison:

c) Calculate the molar enthalpy of combustion of octane, C8H18,

given that 6.593 g causes 150. mL of H2O in a 375 g Al can, to

increase its temperature by 75.2 °C.

M x m

T )cm c(m - = )H(CH

188

188

22

HCHC

AlAlOHOH188comb

Hess’ Law and Heats of Formation

A State Function: _____________________ Definition:

• When a reaction that can be expressed as the ____________ ____,

, of two or more __________ reactions, the enthalpy of reaction,

Hrxn, is the algebraic sum of the ___________________ rxn

enthalpies, Hx. Standard Enthalpies of formation , Hfº

• Are often used to calculate _______

• The enthalpy (_______ or _______ of heat energy) for the

_____________ of 1 mole of the ______________ from its

____________ in their standard state.

• For an _____________ in its standard state, Hfº = 0

• Standard state is SATP is 25ºC, 100. kPa

• These equations are created from its ________________

eg.

Rules for applying Hess’ Law:

1) Use the ________ reaction steps or _______ the _____________

equations of the overall reaction and find the value of Hx or

Hfº for each step. 2) ___________ intermediate steps (reactions) as needed and

remember to __________ the _____ of the Ho for that reaction

(multiply by -1).

eg.

3) ____________ intermediate reactions as necessary to _______

the coefficients in the overall equation.

Remember to multiply the Ho values by the same multiplier. eg. 4) ____________ _______ the elements/compounds that are the

________ on ________ sides. Add the _____________ terms

and determine the Horxn

from the ______________ ______ of

Ho values for all the ________________ steps. 5) If applicable use

Hrxnº = nHfº (products) - nHfº (reactants) eg.1 Determine the heat of reaction for the following: 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g) Hrxn = ?

Using Formation equations:

eg.2 Determine the heat of reaction for the following: C2H4 (g) + H2 (g) C2H6 (g) Hrxn = ?

Using the following GIVEN reactions:

(1) C2H4 + 3 O2 2 CO2 + 2 H2O (l) H1 = -1400.5 kJ

(2) C2H6 + 7/2 O2 2 CO2 + 3 H2O (l) H2 = -1550.0 kJ

(3) H2 + 1/2 O2 H2O (l) H3 = -285.8 kJ

eg.3 Repeat eg. 2, but use the Summation Formula of Hess’ Law: Hrxnº = nHfº (products) - nHfº (reactants)

eg.4 Determine the heat of reaction for the following: C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (g) Hrxn = ?

Using Formation equations:

Bond Energies at SATP

Bond kJ/mol Bond kJ/mol Bond kJ/mol Bond kJ/mol Bond kJ/mol Bond kJ/mol

H-H 432 C-N 305 N-H 391 F-F 154 C=C 614 N=O 607

H-F 565 C-O 358 N-N 160 Cl-Cl 239 CC 839 N=N 418

H-Cl 427 C-F 485 N-O 201 Br-Br 193 O=O 495 NN 941

H-I 295 C-Cl 339 O-H 467 I-I 149 C=O 745 CN 891

C-H 413 C-Br 276 O-O 146 S-H 347 CO 1072 C=N 615

C-C 347 C-S 259 S-S 266 S-Cl 253

Specific Heat Capacities, c at SATP

Substance J/(g•°C) Substance J/(g•°C)

aluminum 0.900 nickel 0.444

Copper 0.385 silver 0.237

Gold 0.129 tin 0.213

Iron 0.444 zinc 0.388

Lead 0.159 water 4.18

Thermodynamic Properties of Organics At SATP

Substance ∆H f° S°

(kJ•mol-1) (J•K-1•mol-)

benzene, C6H6(l) +49.0 172.8

bromoethane, CH3CH2Br(g) -90.5 -----

bromomethane, CH3Br (g) -37.2 246.3

butanal, CH3CH2CH2CHO(l) -241.2 ----

butane, n-C4H10 (g) -126.5 310.1

butan-1-ol, C4H9OH(l) -327.4 228.0

but-1-ene, C4H8 (g) -0.4 305.6

but-1-yne, C4H6 (g) +165.2 -----

carbon tetrachloride, CCl4 (l) -128.4 216.4

CCl4 (g) -96.0 309.9

chloroethene, CH3CH2Cl(g) -136.8 263.9

chloromethane, CH3Cl (g) -82.0 234.5

cyclopropane, (CH2)3 (g) +53.3 -----

1,2-dichloroethane, (CH2Cl)2(g) -165.0 -----

ethanal, CH3CHO(g) -191.5 160.2

ethane, C2H6 (g) -83.8 229.5

ethane-1,2-diol, (CH2OH)2(l) -454.8 166.9

ethanoic acid, CH3CO2H(l) -484.5 159.8

(acetic)

ethanol, C2H5OH(l) -277.1 160.7

ethene, C2H4 (g) +52.5 219.5

ethoxyethane, (CH3CH2)2O(g) -279.0 251.9

ethyne, C2H2 (g) +228.2 201.0

fluoromethane, CH3F (g) -247.0 -----

glucose, C6H12O6(s) -1260.0 212.1

Substance ∆H f° S°

(kJ•mol-1) (J•K-1•mol-1)

hexane, n-C6H14 (l) -198.6 296.1

iodoethane, CH3CH2I(g) -40.7 -----

iodomethane, CH3I (g) -15.5 163.2

methanal, HCHO(g) -108.7 218.7

(formaldahyde)

methane, CH4 (g) -74.8 186.2

methanoic acid, CH3OH(l) -425.1 129.0

(formic)

methanol, CH3OH(l) -239.1 239.7

methoxymethane,

CH3OCH3(g) -184.0 266.7

methylpropane, C4H10 (g) -134.5 294.6

napthalene, C10H8 (s) +77.7 -----

octane, n-C8H18 (l) -250.0 361.1

pentane, n-C5H12 (l) -173.1 262.7

phenylethene or styrene +103.8 345.1

phenol, C6H5OH(s) -165.0 ----

propanal, CH3CH2CHO(l) -217.1 ----

propane, C3H8 (g) -104.5 269.9

propanone(acetone) -248.0 198.8

propan-1-ol, C3H7OH(l) -302.7 196.6

propene, C3H6(g) +20.2 266.9

propyne, C3H4 (g) +186.6 248.1

sucrose, C12H22O11(s) -2221.0 360.2

2,2,4-trimethyl pentane -259.2 328.0

urea CO(NH2)2 -335.5 104.0

Thermodynamic Properties of Inorganics At SATP)

Substance ∆H f° S° ∆G f

°

(kJ•mol-1) (J•K-1•mol-1) (kJ•mol-1)

Al(s) 0 28.3 0

Al2O3 (s) -1675.7 50.9 -1582.3

Al2(SO4)3 (s) -3405.5 50.9 -1582.3

BaCO3 (s) -1216.3 112.1 -1142.0

BaCl2 (s) -860.2 124.1 -813.5

BCl3 (g) -404.0 291.1 -390.1

B2O3 (s) -1273.0 54.2 -1144.1

Br2 (l) 0 152.0 0

Ca(s) 0 41.4 0

CaCO3 (s) -1207.0 92.9 -1128.8

CaBr2 (s) -682.8 130.2 -1352.4

CaCl2 (s) -795.8 104.6 -748.1

CaO(s) -634.9 38.1 -566.5

Ca(OH)2 (s) -986.1 83.4 -901.7

Ca3(PO4)2 (s) -4119.0 236.9 -3897.7

CaSO4 (s) -1434.1 108.4 -1326.8

C(s) graphite 0 5.7 0

C(s) diamond +1.9 2.4 +2.9

CO (g) -110.5 197.66 -137.2

CO2 (g) -393.5 213.78 -394.4

Cl2 (g) 0 223.1 0

Cu (s) 0 +33.2 0

CuCl (s) -137.2 86.2 -119.9

CuCl2 (s) -220.1 108.1 -175.7

Cu2O (s) -168.6 93.1 -146.6

CuO (s) -157.3 42.6 -129.7

CuSO4 (s) -771.4 109.2 -663.6

CuSO4•5H2O s) -2279.0 301.6 -1887.1

F2 (g) 0 202.8 0

H2 (g) 0 130.7 0

H2O2 (l) -187.8 109.6 -120.4

HBr(g) -36.3 198.7 -53.5

HCl (g) -92.3 186.9 -95.3

HCl (aq) -167.2 56.7 -131.8

HCN(g) +135.1 201.8 +125.2

HF (g) -271.1 +173.8 -273.2

HI (g) +26.5 206.6 +1.8

HNO3 (l) -174.1 155.6 -80.7

HNO3 (aq) -207.0 ----- -----

H3PO4 (s) -1279.0 110.5 -1119.1

H2S (g) -20.6 205.8 -33.6

H2SO4 (l) -813.8 156.9 -690.0

H2SO4 (aq) -909.3 20.16 -743.4

I2 (s) 0 116.3 0

I2 (g) +62.4 180.79 -----

Fe (s) 0 27.8 0

FeO (s) -272.0 57.6 +245.1

Fe2O3 (s) -824.2 87.4 -742.2

FeCl2 (s) -341.8 118.0 -302.8

FeCl3 (s) -399.5 142.3 -344.0

Substance ∆H f° S° ∆G f

°

(kJ•mol-1) (J•K-1•mol-1) (kJ•mol-1)

Pb (s) 0 64.8 0

PbCl2 (s) -359.4 136.0 -314.1

PbO (s) -219.0 66.5 -188.6

PbO2 (s) -277.4 68.6 -----

Mg (s) 0 32.7 0

MgCO3 (s) -1095.8 65.7 -----

MgCl2 (s) -641.3 89.6 -591.8

Mg(OH)2 (s) -924.5 63.2 -----

MgO (s) -601.6 27.0 -569.4

N2 (g) 0 191.6 0

NH3 (g) -45.9 192.8 -16.5

N2H4 (l) +50.6 121.2 +149.3

N2H4 (g) +95.4 237.1 -----

NH4Cl (s) -314.4 94.6 -202.9

NH4NO3 (s) -365.6 151.1 -183.9

NO (g) +90.2 210.8 +86.6

NO2 (g) +33.2 240.1 +51.3

N2O (g) +82.1 219.9 +104.2

N2O4 (g) +9.2 304.3 +97.9

O2 (g) 0 205.1 0

O3 (g) +142.7 238.9 +163.2

PCl3 (g) -319.7 217.2 -----

PCl5 (g) -443.5 364.6 -----

K (s) 0 64.2 0

KCl (s) -436.7 82.6 -409.1

KClO3 (s) -397.7 143.1 -296.3

KOH (s) -424.8 78.9 -379.1

Ag (s) 0 42.6 0

AgBr (s) -100.4 107.1 -97.4

AgCl (s) -127.0 96.3 -109.8

AgNO3 (s) -124.4 140.9 -33.4

Ag2O (s) -31.1 121.8 -11.3

Na (s) 0 51.2 0

NaBr (s) -361.1 86.8 -350.2

Na2CO3 (s) -1130.7 135.0 -1044.0

NaCl (s) -411.2 115.5 ------

NaF (s) -571 51.7 -545.6

NaOH (s) -425.6 64.4 -379.5

NaI (s) -287.8 98.5 -287.3

S8 (s) rhombic 0 31.8 0

S (g) +278.8 167.8 +283.3

SCl2 (l) -49.0

SO2 (g) -296.8 248.2 -300.2

SO3 (g) -395.7 256.8 -371.1

SnO (s) -280.7 57.2 -----

SnO2 (s) -577.6 49.0 -----

H2O (l) -285.8 70.0 -237.1

H2O (g) -241.8 188.8 -228.6

ZnO (s) -350.5 43.7 -----

ZnS (s) -206.0 57.7 -----

Energy and Driving Forces:

The Laws of Thermodynamics

Spontaneous Reaction:

• given the required Ea, ________________ ___________the

reaction ________ to proceed by ________

• may be _____ or _____ Thermodynamics:

• Study of energy ______________________.

• There are ___ Laws of ________________________

– used to ____________ reaction _________________ 1st Law of Thermodynamics

• Law of Conservation of ______________

• ________ energy of the universe is ____________

• Energy ________ be ________ or ______________, just

_______________ from _____ form to ___________ For a chemical reaction

∆Hºuniverse = _________________________________________ In a chemical _____________, the PE of the reactants and products

___________ in the ______________ of __________ from the:

1) _________________ to the chemical __________ (ENDO)

2) chemical __________ to the _________________ (EXO)

Enthalpy changes and Spontaneity • Bond energy (BE): the minimum energy required to break one

mole of bonds between two atoms ______________

• Equals the energy released when 1 mole of bonds are formed,

_________

• The greater the value, the more stable the bond SO, the Enthalpy of reaction can also be calculated by:

∆Hreaction = __________________________________________ eg. C2H4 + 3 O2 2 CO2 + 2 H2O (g)

Using Hess’ Law Summation Formula:

The values from the methods _________ as the BE values are

____________ from several compounds with that bond type.

∆Hreaction __________

• indicates the formation of stronger bonds and more stable

compounds

• is most likely to be spontaneous reaction Entropy Changes (∆S) and Spontaneity

• the greater the # of ways particles can arrange themselves,

the less ordered they are

• the greater the # of ways a particular state can be achieved,

the more likely that state is going to exist Entropy , Sº • The measure of disorder or randomness

• increased disorder or entropy favours spontaneity,

____________

2nd Law of Thermodynamics ∆Sºuniverse = • disorder increases, since ____________

The following increases entropy:

1. The __________ of a gas ____________

2. The _________________ of the system ____________

3. Physical state _______ to ________ to ______

4. ___________ in the number of _________ produced

5. Breaking ___________ molecules into __________ ones

3rd Law of Thermodynamics

• At 0K all motion _________, the _________ of attraction have

___________ entropy to a ___________

• S = ___ at T = ____

• As a result as ___ ___________, S must ___________

• S is a _____________ of the __________ needed to achieve a

level of _____________ by __________________ the FA

• This is ________________ on the ______________ and

the _________________ reached

• _____________

• ∆Srxnº = ______________________________________

• ∆Srxnº ________________________________________

Entropy Calculations eg. Calculate ∆Sº for

2CO(g) + O2(g) 2CO2(g) ∆Srxnº = nSº products - nSº reactants

=

___________________________________________________

This would mean it is nonspontaneous, but we need to consider the

∆Hrxnº

∆Hrxnº = n∆H º products - n∆H º reactants

Gibb’s Free Energy – Enthalpy and Entropy

• Unites the 2 reaction driving forces • ∆G = ∆H – T∆S

If: ∆G < 0 ___________________________

∆G = 0 ___________________________

∆G > 0 ___________________________________________ For Spontaneity: ∆H ∆S ∆G spontaneity - + _____ ___________ + - _____ ___________ - - ___________ ___________ ___________ ___________ + + ___________ ___________ ___________ ___________

Calculating ∆Gº from ∆Hº and ∆Sº

eg. 2CO(g) + O2(g) 2CO2(g)

∆Hºrxn = -566.0 kJ (from before) ∆Sºrxn = -172.86 J/K = -0.17286 kJ/K (from before) ∆Gº = ∆Hº - T∆Sº =

_________________________________ Calculating ∆Gº from ∆Gfº

• these values are not in the textbook, but are in the reference

booklet

Use:

∆Gº = ____________________________________ Predicting Change in Spontaneity

• Temperature and equilibrium As ∆G = 0 at equilibrium and ∆Gº = ∆H – T∆S

then: 0 = ∆H – T∆S

• this gives the _____________ at which the system _________

________________

• need to examine ∆H and ∆S to determine if the reaction is

spontaneous ___________ or _________ the temperature

_______________ eg. For previous example: 2CO(g) + O2(g) 2CO2(g)

T = ∆H

∆S = as ∆H –ve and ∆S –ve , then spontaneous at low T

T >

T <

Chemical Kinetics

Definition: the ________ of reaction rates Rate of Reaction: The change in

concentration, ∆[ ], of a

_________ or a __________ per

unit of time • can be determined from a [ ] vs

time graph Average Rate of Rxn: from the

________ of the secant drawn

between __ points on the

________ over a ________

______ interval.

Instantaneous Rate of Rxn: from the _______ of the

___________ to the ________ at a __________ moment in

_____. Reaction Rates are given as: • as rates are __________ and as the rxn proceeds,

______________ • the rates of the different reactants and products are ________ by

the balanced chemical reaction. In General, for aA + bB cC + dD

Rrxn req’d =

eg. For 2 C2H6 + 7 O2 4 CO2 + 6 H2O

The rate of reaction with respect to C2H6 is 4.0 x 10-4 mol/L•s.

State the Rrxn with respect to each product and reactant.

smol/L 10 x 4.0 t

]H[C - )H(CR 4-62

62rxn

Measuring Raterxn

• requires an _______________, _____________ change that

doesn’t disturb the rxn:

1) ________________________________________________

2) ________________________________________________

3) ________________________________________________

Factors Affecting the Raterxn

Nature of the Reactant eg. aqueous, gas, liquid, reactivity

• depends on the _______ of the ______________ particle

forces.

1) Concentration ___________________________________

2) Temperature ___________________________________

3) Surface Area ___________________________________

4) Catalyst

Trial Form of

CaCO3 [HCl]

(mol/L) Temp

(°C) Time

Taken Conclusion

1 powdered 0.50 22 45 s Control

2 powdered 1.00 22 23 s

3 solid chips 0.50 22 > 4 h

4 powdered 0.50 40 14 s

Born Haber Cycle for the formation of SCl2

1/8 S8 (s) + Cl2 (g) SCl2 (l)

Born Haber Cycle for the combustion of Ethene. C2H4

C2H4 (g) + 3 O2 (g) 2 CO2 (g) + 2 H2O (g)

The Rate Law and the Rate Law Equation

As various factors (T, [ ], etc) affect the rate of the rxn, their

impact MUST be determined _________________ through

______________________. For aX + bY product Rate Law:

Raterxn _____________________ where m, n R determined experimentally

i = initial concentration Rate Law Equation:

Raterxn = ___________________

where k is the rate constant for a specific temp. Order of Rxn:

m = ________________________

n = ________________________

m + n = ________________________

eg. Raterxn =

∆Rate

∆[ ] 0th order 1st order 2nd order

2 (doubles) _______ _______ _______

3 (triples) _______ _______ _______

______________________________

For the following reaction, the data in the table was obtained:

2 NO (g) + 2 H2 (g) N2 (g) + 2 H2O (g)

Trial [NO] i

(mol / L)

[H2] i

(mol / L)

Initial Rate

(mol / L•s)

Conclusion

1 0.100 0.100 0.00123

2 0.100 0.200 0.00246

3 0.200 0.100 0.00492

k can be determined from any of the expt trials:

using Trial 3: [NO]i = 0.200 mol/L, [H2]i = 0.100 mol/L

R = 0.00492 mol/L•s and R= k [NO]2[H2]1

Reaction Rate and Time • often rates of ____________ are measured by the ________ that

it takes for a certain _________ _________ in a rxn

• usually measured by a ________ __________ – called a

________ ________

For A products

Graphically: Zeroth-Order

Reaction

First-Order

Reaction

Reaction with

Order Greater

Than One

Fixed Second-

Order Reaction

1/

t

1/

t

1/

t

1/

t

Reaction Mechanisms

• are the ______ of _______ that make up a ___________ .

• each one of these, called _______________ _______, involves

the collision of __, __ or at MOST __ particles.

• each Elementary step has a _____ law that _________ the

______________ of the ___________.

• the Rate lawrxn ________ the Rate ____ for the ________ step

eg. 2 NO(g) + 2 H2(g) N2(g) + 2 H2O(g)

Occurs in 3 elementary steps: 1)

2)

3)

where N2O2 and N2O are __________ __________________,

species that are ________________ and then ______________

• Step 2 is the ______ step, with the ___________ Ea

• this is called the _______ _______________________

___________, _____

• the overall Raterxn is proportional to the rate of this step only as

IT ______________________ THE _________ !!!!!!

• the Raterxn can be taken _________ from the ___________ in

the ______ .

as 2) N2O2 + H2 H2O + N2O slow

but N2O2 is not a ___________ in the __________ equation

(intermediate)

so the rate law equation is rewritten to include ___________ only. from 1) NO + NO N2O2

Developing the Reaction Mechanism from Experimental Data For the reaction

NO2 + CO NO + CO2

Trial [NO2]i

mol/L

[CO]i

mol/L Rate Conclusion

1 0.0100 0.0100 4.2 x 10-3

2 0.0200 0.0100 1.68 x 10-2

3 0.0100 0.0200 4.2 x10-3

Possible _______________ to create ___________ reaction NO2 + CO NO + CO2

• assume _______ step (1) – often true as reactants require bond

breakage

• Each _______, try to ________ a ___________

1)

2)

where _______ is a reaction intermediate, but _______ is a

________________ – it is _______ up and then

_______________.

Chain Reactions • are _____________ in which an _____________ formed in an

____________ step is _____________ and then acts as a

____________ . • often involve ___________ conditions due to the vast amount of

available ________, gases and the existence of _____ ________

which have an _________ ___ produced when ____________

molecules’ ______ are ______ and are ______________

reactive. • Are _______ to be confused with other ______________ .

eg. H2 + Cl2 UV

2 HCl Initiation

Propagation

Propagation

Termination

overall • These free radicals also cause big problems in the ozone layer

Normally:

Problem:

Initiation

Propagation

Propagation

Termination

overall

Collision Theory of Rates of Reactions The ___ requirements that ______________ whether a

____________ collision between molecules _________ in a

_____________ are:

1) Sufficient energy

2) Correct Orientation

If ______ of these ______________ are met then an ___________

species, called the ______________ ____________ is formed.

This can be:

_____________________________________________________

_____________________________________________________

Raterxn __________________________________________

Factors Affecting Raterxn and Collision Theory Raterxn = collision frequency x fraction effective

_________________ ___________________

_________________ ___________________

_________________ ___________________ I) Concentration

• ________________________________________________

________________________________________________

• ________________________________________________

________________________________________________

II) Surface area

• ________________________________________________

• ________________________________________________

________________________________________________

• ________________________________________________

III) Temperature

i) collision frequency

• ________________________________________________

• ________________________________________________

ii) fraction effective

• ________________________________________________

• ________________________________________________

IV) Nature of reactant

• ________________________________________________

• ________________________________________________

________________________________________________

________________________________________________

• ________________________________________________

• ________________________________________________

V) Catalyst with PE diagram

• ________________________________________________

• ________________________________________________

• ________________________________________________

________________________________________________

• ________________________________________________