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Unit 2 - Reaction Kinetics/Speed
What is a chemical reaction? It’s when new compounds form via bond rearrangements
Basically it is a successful molecular collision.
Molecules buzz about and bang into each other all the time, if they do so with enough force they rearrange and
that is a reaction or a molecular collision. Take 50 lego bricks and make a truck, now smash it, ok take the same
bricks and now make and SUV. The SUV is a Nnew thing made (product) from the same bricks that made the
truck (reactants). Now how do these molecules get the energy to do all this rearranging.
Kinetic Molecular Theory (K.M.T):
Other than at -273 degress celius, all molecules (or atoms) are moving and so possess energy.
1.
2.
3.
4.
5.
Collisions between molecules are elastic (no energy is gained or lost during the collision).
There are no attractive or repulsive forces between the molecules (ignore intermolecular attractions).
The average kinetic energy of a molecule is proportional to temp.
The molecules in a gas occupy no volume (that is, they are points).
Gases are small particles traveling in a straight-line
The collision theory states that molecules must collide to react, and the must collide with the correct goemetry
Collision Geometry: A2
A
+
A
+

B2
B
2AB:
No go!
B
the above collision has unfavourable alignment, Next the geomerty is favorable
A
B
B
A
+
+
A
B
A
B
Reactant Molecules (A2 & B2)
APPROACH EACH OTHER
A
B
ACTIVATED
COMPLEX
(A2B2)
B
A
breaks apart to form
the PRODUCTS
(2AB)
Bond breaking (above where A-A & B-B is broken) is endothermic process (requires the input of energy). Bond
forming (above where the 2 A-B’s are forming) is exothermic, it releases of energy.
Yet, not all collisions result in bond reformations (successful collisions) because particles must collide with
sufficient intensity (activation energy), not just the correct geometry.
For example, at 20°C oxygen in air moves about 1650 km/h, and they are colliding with this paper (assuming
you downloaded). So why is it not on fire, they are colliding madly? Well, 1650 km/h is not madly enough, is is
actually too slow for successful collisions. Wow, what more does it need, why, more speed of course! And how
is that done? By energy, usually in the form of heat. Let’s look at a reaction by looking at methane burning:
Ex: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l)
Bond Energies.
Bonds being broken (ABSORBS ENERGY):
REACTANT BONDS ONLY are:
CH4(g) + 2 O2(g)  Products
4 C-H & 2 (O=O) bonds bonds are broken (absorbing energy)
Bonds formed (RELEASE ENERGY):
Reactants  CO2(g) + 2 H2O(l)
2 C=O & 4 H-O bonds of water form (releases energy)
As the type and number of bonds being broken and formed change, so will the change in enthalpy.
H = Hbond energy reactants - H bond energy products
Calculating Bond Energies: Bond breaking is endothermic (energy is absorbed to break the bonds). The potential
energy of the separated particles is greater than the bonded particles. Reforming bonds requires an energy input.
The amount of energy needed to separate one mole of a bonded species is called the bond energy.
Example: Using average bond energies (always given and written in blue above each bond), calculate the
enthalpy change for the combustion of ethanol.
So now we know what a reaction is, breaking bonds then remaking them, we need to now to speed up
theses reactions. How would you break things more easily? Let us sped things up.
1)
2)
3)
4)
5)
The 5 Factors that affect reaction rates are:
surface area of a solid reactant
concentration or pressure of a reactant (pressure and volume wil be important here)
temperature of the reactants
nature of the reactants
a catalyst introduced to the reactants
A change in one or more of these factors may alter the rate of a reaction.
1) Surface Area: is the exposed matter of a solid substance. Imagine that you are holding a perfect cube of
magnesium. The surface area is the sum of the area of all six sides of the cube, like dice.
The surface area of the cube is increased by chopping the cube into 2 smaller cubes. The surface area is now like
2 die, 12 sides in total, the number of collisions per second and increases the reaction rate.
In general, solid reactants are slow (except when powdered to increase the surface area). Gaseous reactants are
real fast, BUT so are ionic solutions fast, why, because their molecules are spread out
eg. (precipitate reactions) Ag+(aq) + Cl-(aq)  AgCl (s)
(aqueous ions are mobile and more concentrated than molecules in a gas and no bonds to break)
2A
Concentration:
2B
Pressure (which is really a change in concerntraytion as c = n /v)
Aqueous solutions and gases have their concentrations expressed in mol/L.
In terms of the collision theory, increasing the concentration of a reactant increases in the number of collisions
For an example, consider the reaction between hydrochloric acid and zinc metal.
In one beaker, 6.0 mol/L HCl is reacted with 2.0 g of Zn. In another, 1.0 mol/L HCl is reacted with 2.0 g of Zn.
Which reaction should occur at the faster rate?
In terms of the collision theory, collisions between zinc atoms and hydrochloric acid are more frequent in the
beaker containing 6.0 M HCl - there is more acid per unit of volume, more frequent collisions.
You can change the concentration of an aqueous species by simply adding more solute (to make it more
concentrated, solute is the solid) or adding more solvent (the liquid/water, to make it more dilute).
You can change the concentration of a gas by adding more gas to a fixed volume or by decreasing the volume.
CHANGING THE PRESSURE, AS WE SEE BELOW IS CHANNGING THE CONCENTARTION.
Pressure
If you have gas molecules, you can increase the pressure by forcing them into a smaller volume.
Under higher pressure (which is a higher concentration) gas molecules collide more frequently and react at a
faster rate. Also note that solids and liquids are not affected by pressure changes.
3.
Temperature
With the exception of some precipitation reactions involving ionic compounds in solution, just about all
chemical reactions take place at a faster rate at higher temperatures. The question is why?
Temperature (in Kelvin degrees – JUST ADD 273 TO ANY CELCIUS TEMP) is proportional to the kinetic
energy of the particles. For example, if the Kelvin temperature of a substance is doubled, then the average
kinetic energy of the particles in that substance is doubled. A general rule is that a 10°C temperature increase can
double a reaction rate. At higher temperatures, particles collide more frequently and with greater intensity.
Increasing the temperature by say 10°C causes some of the intermediate speed molecules to move faster. The
result is more molecules with sufficient kinetic energy to form an activated complex upon collision!
Thus there are two effects of increasing temperature: greater collision intensity and more frequent collisions.
Activation Energy -minimum energy needed for a reaction to take place. A higher temp, a greater fraction of
the molecules have KE > = the Ea. So this just says to have a reaction you need t have enough energy to be put
into reactants in order for their bonds to break, that minium enegy is called activation energy.
The Relationship Between Kinetic (K.E.) and Potential Energy (P.E)
Potential energy is defined as energy of position. Energy is conserved in chemical and physical changes. That
means kinetic energy (energy of movement) can be converted to PE and vice versa. Moving particles possess
KE. As two molecules move towards each other they lose their KE as the repulsive forces of their electron
clouds slow the molecules. At collision, the molecules come to a complete momentary stop, and the KE is zero,
at that moment the PE is at the max. All the KE is converted to the PE of bonds. If collisions are hard enough to
cause old bonds to break and new bonds to form. This critical kinetic energy level is called the activation
energy.
e
e
e
e
+
e
negatively charged
e
e
e
e
+
e
e
Repulsion
e
e
e
negatively charged
as molecules approach the electron repulsion slows molecules down so kinetic energy decreases, as they push
against the repulsive force potential energy increases; KE + PE = Total E (stays constant)
If colliding molecules don’t have
enough KE to convert to PE to make
it “over the Energy Barrier”, there is
NO reaction.
Complex
ENERGY BARRIER
Potential
Energy
(kJ)
REACTANTS
PRODUCTS
Reaction Proceeds
The Activated Complex, a short-lived, unstable combination of reactant atoms
100
90
80
70
Ea = 85 - 50
= 35 kJ
60
50
REACTANTS
40
30
20
PRODUCTS
10
Progress of Reaction
Temperature’s role the temperature determines how many molecules will have energy > Ea (to make it over the
barrier & have a successful collision)
Most of us don’t
have the energy to
make it over
Yaaay! A few of us
made it!!!
Energy
REACTANTS
PRODUCTS
Progress of Reaction
Notice in the diagrams above, only a small fraction of the molecules had enough energy.
At the higher temp, more of the molecules have the energy to “make it over”, so Rx faster.
some don’t
have the
energy
Yaaay! A greater
fraction made it
Energy
REACTANTS
PRODUCTS
Progress of Reaction
Increasing the temperature increases the fraction of molecules which have sufficient energy to form the
Activated Complex (or to “make it over” the activation energy barrier.)
Consider two reactions AT THE SAME TEMPERATURE:
100
100
90
80
70
Ea
60
50
40 REACTANT
30 S
20
10
PRODUCTS
Progress of Reaction
Reaction A
90
80
Ea
70
60
50
40
REACTANTS
30
20
10
PRODUCTS
Progress of Reaction
Reaction B
Which reaction is above faster? Well, A is, because the activation energy barrier is smaller to overcome
We have just delat with 3 ways to increase a reaction rate and the last one, temperature was quite a lot to explain.
The next 2 ought not be so bad!
4)
Nature of Reactants
Individual properties of substances also affect reaction rates. The scope of these properties is broad and there are
few generalizations that you can apply consistently. Some of the properties in this category are state of matter,
molecular size, bond type and bond strength. There are several subcategories in 4.
4 a) State of Matter (subcategory)
Gases tend to react faster than solids or liquids: It takes energy to separate particles from each other. In order to
burn candle wax, the solid wax has to be melted and then vaporized before it reacts with oxygen. Methane gas is
already in the gas state so it burns faster than wax.
fast
slow
Although being dissolved in water is not a state, aqueous ions tend to react faster than the 3 states of matter:
Solid Pb(NO3)2will react with solid KI, but the reaction is really slow. That's because the ionic bonds are hard to
separate. However, when aqueous solutions of these compounds are mixed, the ionic bonds are already broken,
the formation of PbI2 is rapid. In aqueous solutions, the ions are dissociated and reactions are fast
please note the states of matter
4 b) Bond Type: Reactions of ionics go faster than reactions involving molecular compounds.
4 c) Bond Strength
Reactions involving the breaking of weaker bonds proceed faster than reactions involving the breaking of
stronger bonds. For example, double carbon to carbon bonds are stronger than single C-C bonds.
4 d) Number of Bonds and Molecular Size
Breaking fewer bonds goes faster than breaking lots of bonds.
The simple ion Fe2+ reacts faster than the larger oxalate (OOCCOO2-) molecule.
Kerosene burns more slowly than CH4 because there are more bonds to be broken, kerosene is a larger molecule
fast
slow
5.
Catalyst: speeds up a chemical reaction without being chemically changed. In other words, the mass of
a catalyst is the same before and after a reaction occurs. The number of successful collisions cab be increased by
lowering the threshold energy (or activation energy)
Basicallly, a catalyst provides an alternative pathway that has a lower activation energy.
The catalyzed pathway (shown as a dotted green line above) has lower activation energy.
Activity: Textbook Readings : page 466 and 467 and 470-471
Textbook Practice Items: pages 467-468: #’s 1, 2, 4, 5, and 6; pg 486-487: 1, 2, 3, 4, 5, 6, 9, 12
Ea, H for forward and reverse reactions
Try this question: Using the graph , find the eats of the forward, revervse and overall reactions:
ACTIVATED
COMPLEX
100
90
POTENTIA
L ENERGY
80
70
60
Ea – Energy
needed for
Reactants to
form the
Activated
Complex
50
40
REACTANTS
- Energy Difference
between Reactants and
Products. (In this case  is
negative so Rx. is
Exothermic
30
20
PRODUCTS
10
Progress of Reaction
Answers:
Ea (forward rx.) = 85-50= 35 kJ
Considering reverse rx.
Ea (reverse rx.) = 85 - 10 = 75 kJ
___________________
Now try another one:
H (forward rx. ) =
35 – 75 = -40 kJ
H (reverse rx. ) =
75 - 35 = 40 kJ
A2 + B2
50
POTENTIAL
ENERGY

2AB
A2B2
45
40
35
30
25
A2 + B2
20
15
10
2 AB
5
Progress of Reaction
ANSWERS
a)
Ea (forward) =
20
KJ
b)
Energy needed to break bonds in A2 & B2
c)
Ea (reverse) = 45 - 5 =
40 KJ
d)
Energy needed to break bonds in AB (A-B)
20
40
KJ
KJ
Now for some FINAL and general ideas about reactions…….
Reaction Mechanisms: Most reactions take place in a series of simple steps, each step depends on the others
before it and the of steps of a reaction (you will not be asked to come up with a mechanism)
Example:
4HBr + O2  2H2O + 2Br2 (5 reactant particles can’t colloide in 1 step)
The Acutal Mechanism is:
step 1:
HBr + O2
 HOOBr (found to be slow elementarty step)
step 2:
HBr + HOOBr  2HOBr (fast elementarty step)
step 3:
HOBr + HBr  H2O + Br2 (very fast elementarty step)
Rate determining step (RDS) - is the slowest step in the reaction. The only way to speed up a reaction is via
the RDS; speeding the fast step has no effect. Each step is an Elementary Step
Determining overall reaction mechanism cancel stuff identical on both sides - add up what’s left.
eg.) 1.)
HBr + O2  HOOBr
2.)
HBr + HOOBr  2HOBr
3.)
2HBr + 2HOBr 2H2O + 2Br2
_________________________________________________________________
overall rx:
4HBr + O2  2H2O + 2Br2
This reaction just occured in 3 steps, and the crossed out species (chemicals) are called reaction
intermediates -a species (atom, molecule or ion) produced in one step and used up in a later step. (appears on
right & on left). Intermediates are: HOOBr
& HOBr
Another example:
2A4+ + B+  2A3+ + B3+
__________________________________
step 1: A4+ + C2+  C3+ + A3+ fast
step 2: A4+ + C3+  C4+ + A3+ slow
step 3: ?
find step 3.
Ans:
C4+ + B+  B 3+ + C2+
PE diagram for a Multiple Reaction Mechanism
Each “bump” is a step, the higher the bump, (greater Ea) and therefore the slower the step
and it is called the RDS (rate determinig step)
PE
Ea (Overall Rx.)
REACTION PROCEEDS
In each of the reactions in the diagram above, the Ea for the overall forward reaction is the difference in
energy between the reactants and the top of the highest peak.
How catalysts work: “to avoid a hill, build a tunnel.” A Catalyst- is an introduced substance which has an
alternate mechanism with a lower activation energy…and
- a catalyst NEVER changes the PE of reactants or of products - only the route between them.
- also there is no change in H of the reaction
Catalysts sometimes work by... providing a surface whose spacing of atoms is just right to break a reactant
molecule and hold it for an attack from another reactant.
eg.) 2H2O2(l)  2H2O(l) + O2(g)
(very slow uncatalyzed)
Catalyzed Mechanism using iodide ions:
step 1) H2O2 + I- ---> H2O + OI- (The catalyst I- is put in.)
step 2) H2O2 + OI- ---> H2O + O2 + I- (The catalyst I- is regenerated.)
overall rx.
2H2O2 ---> 2H2O + O2
IMPORTANT: CATAYLSTS DO NOT MAKE MORE PRODUCT, the make things faster, ok, but they
cannot make the Law of Conseravation of mass disappear, they are good, but NOT that good.
Reaction Rates: STSE: Smog, Catalytic Converters:
All cars today must have catalytic converters to minimize pollutants or smog. The word "smog" means "smoke
and fog". Engine exhaust contains NO, NO2, unburned hydrocarbons and CO. This is smog is created thusly:
½ N2
+
O2

NO2
The high CO is poisonous as it combines with hemoglobin in our red blood cells better than oxygen.
Compounding the CO problem, the NO and NO2 components of smog can combine with O2 in the air to create
ground level ozone, O3. At ground level, ozone O3, irritates our respiratory tract and eyes.
Ozone Formation: Reaction Mechanism: The currently accepted mechanism follows:
Step 1:
2 NO(g) + O2(g)  2 NO2(g)
fast
Step 2:
2 NO2(g)
 2 NO(g) + 2O(g)
slow (rate determining step. RDS)
Step 3:
2 O(g) + 2 O2(g) 2O3(g)
fast
The Catalytic Converter
To reduce smog, NO must be converted to harmless N2 and O2, and CO must be converted to CO2. A catalytic
converter is having a convoluted surface providing a large surface area to hold the catalyst, which is palladium
(Pd) and or platinum (Pt) and rhodium(Rh) catalytic surfaces.
On a catalytic surface NO & CO from the exhaust are adsorbed (called chemisorption). This physically weakens
the N-O chemical bonds and they are easily broken (Steps 1 & 2 below).
(ad) means absorbed on catalyst surface)
Step 1:
2 CO(g)  2 CO(ad)
Step 2:
NO(g)  N(ad) + O(ad)
Step 3:
N(ad) + NO(g)  N2(g) + O(ad)
Step 4:
2 CO(ad) + 2 O(ad)  2 CO2(g)
One potential problem with the operation of a catalytic converter is that it may lose its effectiveness when
impurities bind strongly to the catalyst surface and accumulate and inactivate it. When this occurs, the
ineffective catalytic converter is said to be poisoned, and must be replaced.
One last issue, up until the end of the 1980s, it was possible to buy "leaded gasoline" Pb(C2H5)4. The addition of
tetraethyl lead allowed a car's engine to run more smoothly with less engine knock. Unfortunately, this
compound poisons catalytic converter surfaces, and in 1990 the sale of leaded gasoline was banned in Canada.
(MTBE) was selected as a fuel additive in North America. Here's an interesting twist in the story - more recent
research has implicated MTBE as a carcinogen (cancer-causing agent). MTBE has been found in the
groundwater wells. Way to go environmentalists.
Reaction Kinetics - reaction speed, how long does it take and how to make reaction go faster. We need a way
to see how fast reactions go, to measure their speed or rate.
Expressing Rates:
= amount (a reactant or product) or Concentration final – Concentration initial
time
Time final - Time initial
EX.
Zn(s) + 2HCl(aq)  H2(g) + ZnCl2(aq)
r = - mass of Zn(s)
time
r = - [ HCl ]
time
(note: [ ] = molar concentration)
r = volume H2
r = [ZnCl2 ]
time
time
Notice the negative in the reactants? Rates of reactions are always positive, however, We do use the
NEGATIVE sign to show reactants as being lost in the equation:
Rate = amount (g, mol, L)
time
(s, min)
Comparing rates using balanced equations: use coefficient ratios – for mol /s (not to grams/s), remember only
MOLES (mole ratio). So figure out one rate and you have the rest with the mole ratio
EX.
2 C2H6 + 7O2

4 CO2 + 6 H2O
- if ethane is consumed at a rate of 0.066 mol /s, calculate the rate of consumption of O2 in mol /s
Solution: 0.066 mol/s C2H6 x 7 mol O2 = 0.23 mol / s O2
2 mol C2H6
if ethane is consumed at a rate of 0.066 mol /s calculate rate of production of CO2
Solution: 0.066 mol C2H6 x 4 mol CO2 = 0.13 mol CO2
s
2 mol C2H6
s
- when other units used – you must use moles to (ride the “mole” bus)
Another example:
Given: 2Al + 3Br2  2AlBr3
If 67.5 g of Al are consumed per second - calculate the rate of consumption of Br2 in g/s.
So convert to MOLES and then the mole ratio.
Solution: 67.5 g Al x
s
1 mol Al x
27.0 g Al
3 mol Br2
2 mol Al
x 159.8 g Br2 = 599 g Br2
1 mol Br2
s
Problem Example
For the oxidation of ammonia 4 NH3 + 3O2 => 2 N2 + 6 H2O
it was found that the rate of formation of N2 was 0.27 mol L–1 s–1.
a) At what rate was water being formed?
b) At what rate was ammonia being consumed?
Solution:
a) From the equation stoichiometry, Δ[H2O] = 6/2 Δ[N2], so the rate of formation of H2O is 3 ×
(0.27 mol L–1 s–1) = 0.81 mol L–1 s–1.
b) 4 moles of NH3 are consumed for every 2 moles of N2 formed, so the rate of disappearance of
ammonia is 2 × (0.27 mol L–1 s–1) = 0.54 mol L–1 s–1.
Comment: it would be acceptable to express this last value as a negative number.
Measuring Reaction Rates
eg. CaCO3(s) + 2HCl(aq) 
H2O(l) + CO2(g) + CaCl2 (aq)
ionic form: CaCO3(s) + 2H+(aq) + 2Cl-(aq)  H2O(l) + CO2(g) +
net ionic form:
CaCO3(s) + 2H+(aq)  H2O(l) + CO2(g) + Ca2+(aq)
HCl(aq)
Ca2+(aq) + 2Cl-(aq)
note, the 2Cl-(aq)ion would not be used for rate expressions, also note that as CO2 escapes, mass of the rest of
the system will decrease, so rate could be expressed as..
r = - mass of container and contents
(open system)
time
Note: rate = slope of graph
Slope will be negative if something is being consumed and positive if something is being produced.
Slope = rise (g)
run (s)
Mass of
Container and
Contents (g)
Rate = Slope (made +)
rise (g)
run (s)
Time (s)
100.
Average slope between 0 and 30 s is shown by a
straight line drawn from time “0” to time “30 s” on
the graph.
Mass
(g)
Slope = rise = 100.g – 20.g = - 80. g = - 2.7 g
run 0.00 s – 30.0 s 30. s
s
20.
0.0
10.
20.
30.
Time (s)
or rate at a certain point in time is the slope of the tangent at that point. INSTANTANEOUS RATE
100.
tangent
Mass
(g)
Rate at t =22 seconds is the slope of the tangent to
the curve at time = 22 seconds.
20.
0.0
10.
20.
30.
Time (s)
6 Ways to Monitoring Reaction Rates
1.) Colour changes
- only in rx where coloured reactant is consumed or new coloured product is formed.
eg.) Cu(NO3)2(aq) + Zn(s)  Cu(s) +
Zn(NO3)2(aq)
blue
grey
reddish brown
colourless
- in this case could measure - intensity of blue fading OR the intensity of reddish brown gas being made
Cu2+(aq) + Zn(s)  Cu(s) + Zn2+(aq)
[Cu2+ is blue!]
colour intensity can be measured quantitatively using a spectrophotometer
net ionic:
-
2.) Temp changes: reactions produce or consume heat, so monitor the temp to show Rx progression
3.) Pressure changes (constant volume or sealed container)
- if more moles of gas (m.o.g.) in products pressure will go up
Zn(s) + 2HCl(aq)  H2(g) + ZnCl2(aq)
None (0.00) m.o.g (moles of gas)
 1 m.o.g. (hydrogen)
(And if more MOG in reactants - pressure will decrease)
- If equal MOG, pressure will not change: NO2(g) + CO(g)  CO2(g) + NO(g)
2 m.o.g.
2 m.o.g.
4.) Volume change (constant pressure)
eg.) if more gas is produced, volume of balloon will increase
5.) Mass changes: if only one solid is used up (could remove periodically and weigh it)
Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq)(periodically remove Mg(s) and weigh what is left)
- or if one gas is produced and escapes, measure mass of what’s left in container
eg)
heat + CaCO3(s)  CaO(s) + CO2(g)
Only CO2 gas escapes. So as
CO2 escapes, the mass of the
container decreases.
rate = mass of container & contents
time
Note: it’s not practical to measure masses of (aq) substances since they are mostly water.
6.) Changes in molar concentration of specific ions
eg)
Mg(s) + 2HBr(aq)  H2(g) + MgBr2(aq)
ionic form: Mg(s) + 2H+(aq) + 2Br -(aq)  H2(g) + Mg2+(aq) + 2Br -(aq)
- could monitor [ H+], using pH meter - it will decrease as rx proceeds
zed
TIAL
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atalyzed
on
ATED
pdverse
GRAM
erse
yzed
2)1)
3)
TANTS
O
e
Br
2route
2 of
TION
erted
123UCTS
ep
ess
goes
NG
on
LEX
ds
EEDS
ion
her
TALYZ
D THE
up.
YZED
ION
Note: Does the [Br -] change? No, it is a spectator ion found both as a reactant and product.
Chemical reactions proceed at different rates. And we ONLY consider the reactants!!!
.
SUMMARIZING WAYS TO MEASURE A REACTION
Methods of Measuring Reaction Rates
Property
Mass
Observation
Mass reactant decreases/product increases
Exothermic reaction /Heat is released/Temperature increases
Temperature
Endothermic reaction /Heat is absorbed. Temperature decreases
Colour
If a reactant or product absorbs light, then the color will vary as the
amount of this species changes
Pressure
If there is a change in number of moles of gas as the reaction
proceeds pressure will change
Conductivity
If the reaction involves ions in solution then the conductivity of the
solution will change as the reaction proceeds.
Concentration
The concentration of reactant and products changes as the reaction
proceeds.
Rate