Reaction Rate Chemistry Quiz

Front 1

chemical kinetics

Back 1

study of rates of reaction

Front 2

6 slow reactions

Back 2

rusting
decomposition
fermentation
yellowing of paper
setting of concrete
baking

Front 3

3 fast reactions

Back 3

combustion
neutralization
precipitation

Front 4

In a reaction we can measure (7):

Back 4

change temp over time
change pressure over time
change mass/moles of rctnts
volume product
pH
colour (spectural photometer)
conductivity

Front 5

Rate =

Back 5

change [product] divided by time

Front 6

concentration of product over time

Back 6

starts at bottom fast and gets slower bcs u run out of reactants

Front 7

Less reactants means

Back 7

less collisions and thus slower reaction

Front 8

concentration of reactant over time

Back 8

starts at top big and gets smaller. 1st fast then slow

Front 9

reactants with a 2,

Back 9

used up twice as fast

Front 10

6 factors affecting reaction rate

Back 10

1 temp. up up
2 agitation up up
3 surface area down down
4 catalysts (faster), inhibitors (slow)
5 [ ] increase of reactants
6 Nature of reaction (some naturally fast, some naturally slow)
increase rate

Front 11

Why would u ever want to slow down a reaction?

Back 11

ex: to slow down food expiration

Front 12

HOW does temp increase reaction rate?

Back 12

When heated, reactants gain energy (meaning there are more molecules with sufficient Kinetic Energy). Hence, they move faster, and there is a bigger chance of collision and reaction. Bigger reaction rate.

Front 13

HOW does concentration increase reaction rate?

Back 13

There are more reactants available, and hence there is a bigger chance they're collide and react.

Front 14

HOW do catalysts increase reaction rate?

Back 14

Change the pathway of the reaction?

Front 15

The rate law relates:

Back 15

reactant concentrations and rate quantitatively.
r=k[A]^m[B]^n

Front 16

in rate law, what depends on temp?

Back 16

k = proportionality constant/rate constant
NOT rate law exponents

Front 17

RATE LAW: m and n are usually

Back 17

0,1,2,3 but can b fractions

Front 18

Rate constant:

Back 18

indicated speed of reaction.
k large: fast: 10^2 (.1 s)
k small: slow: 10^ -3 (2h)

Front 19

k depends on

Back 19

not reactant. constant thruout reaction
temp and pressure

Front 20

2nd order rate law units

Back 20

L/ mol x s

Front 21

1 L how many cm^3

Back 21

1000 cm^3 in

Front 22

systems consist of

Back 22

atoms, ions, molecules, in constant motion with varying KE's.

Front 23

temp =

Back 23

avg KE = .5mv^2

Front 24

Reactions occur when particles collide with:

Back 24

a) sufficient energy
b) correct geometry (alignment, position)

Front 25

Reaction rate depends on

Back 25

the frequency of collisions
amount of energy

Front 26

homogeneous rxns

Back 26

reactants in the same state
ex:
N2(g) + 3H2(g) --> 2NH3(g)

Front 27

Heterogeneous rxn

Back 27

reactants in diff states

Front 28

bimolecular rxn

Back 28

2 reactants

Front 29

trimolecular rxn

Back 29

3 reactants

Front 30

unimolecular rxn

Back 30

1 reactant

Front 31

Rate = x * y

Back 31

collision frequency times fraction of effective (sufficient) collision

Front 32

3 things that affect collision frequency

Back 32

concentration
surface area (only 4 heterogeneous)
temperature

Front 33

3 things that affect fraction of effective (sufficient) collision

Back 33

temp
catalysts
nature of reactant

Front 34

the higher the temp, the more

Back 34

molecules with higher KE, the temp

Front 35

in "# molecules w/particular KE" vs "KE" graph (y,x),

Back 35

the graph righter is higher temp

Front 36

low rate even when collision frequency high, bcs

Back 36

effective collisions can b low

Front 37

activation energy

Back 37

the min increase in PE of a system for molecules to react (begin bond braking)

Front 38

Maxwell-Boltzmann distribution

Back 38

KE vs # molecules (x,y)
With lower activation energy, more molecules react

Front 39

Lower activation energy when (2)

Back 39

weak bonds
simple geometry
thrfr faster rate

Front 40

# molecules with min energy required increases with

Back 40

increasing temp (increase #molecules . . .)

Front 41

when asked: y high temp=more collisions?

Back 41

say:there are more molecules w/min energy

Front 42

catalysts provide a

Back 42

lower energy pathway for the rxn to occur. Diff mechanism w/intermediates.

Front 43

where y is PE,

Back 43

you point out activation energy and delta H.
x is reaction progress

Front 44

2 other word 4 activation energy

Back 44

energy hill
activated complex

Front 45

During climb up energy hill (2)

Back 45

repulsive forces increase causing molecules to slow down
and PE increases as KE decreases

Front 46

Products that have higher PE but lower KE. Molecules they come in contact with,

Back 46

slow down. Hence, lower KE means feels cooler

Front 47

After energy hill there is a

Back 47

transition state. Unstable molecules

Front 48

A half life (t .5) is

Back 48

the time for .5 of the nuclei in a radioactive sample to decay OR in a 1st order rxn the time 4 .5 the amount of a reactant 2 b used up.

Front 49

In living organisms there are 15.3

Back 49

disintegrations of C-14 per minute per gram.

Front 50

When organism dies, rate of disintegration

Back 50

decreases w/a t.5 = 5730 years

Front 51

Radiodating, they

Back 51

look at ration of c-14 to c-12 of organism and compare to recently decayed organism

Front 52

relative half-life to rate constant

Back 52

kt.5=ln2=.693

Front 53

half life to rate constant come from

Back 53

simple integration of rate law expression as it relates to t.5

Front 54

spontaneity of a reaction

Back 54

inherent tendency not speed

Front 55

to find reaction rate, we must choose conditions in which

Back 55

reverse reaction neglected

Front 56

differential rate law

Back 56

rate divided by concentration

Front 57

integral rate law

Back 57

rate divided by time

Front 58

half life (txtbk)

Back 58

ln([a]intl divided by [a]) = kt