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AP®︎/College Chemistry
Course: AP®︎/College Chemistry > Unit 5
Lesson 1: Reaction ratesFactors affecting reaction rates
The rate of a chemical reaction is influenced by many different factors, including reactant concentration, surface area, temperature, and catalysts. In general, increasing the concentration of a reactant in solution, increasing the surface area of a solid reactant, and increasing the temperature of the reaction system will all increase the rate of a reaction. A reaction can also be sped up by adding a catalyst to the reaction mixture. Created by Jay.
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how does a catalyst speed up the reaction? I don't understand what it does.(2 votes)
So every reaction begins by having the reactant's bonds broken somehow. This requires an input of energy which we quantify with the activation energy. In a simplified explanation a catalyst works by reducing this activation energy barrier. Requiring less energy to begin the reaction means more reactant molecules possess the minimum energy required and so the reaction rate speeds up by being able to process a higher reactant concentration in the same amount of time compared to the non-catalyzed reaction.
A more complete explanation would explain, why does a catalyst reduce the activation energy barrier? This requires us to be familiar with reaction mechanisms. A reaction mechanism is a step-by-step process of how a chemical reaction occurs using simpler elementary reactions. Elementary reactions are simple reactions where single collisions or decompositions occur. Most chemical reactions are actually a combination of these smaller elementary reactions and the rate of a reaction is constructed using the slowest of these elementary reactions. Anyhow a catalyst works by providing an alternative reaction mechanism which is faster than the uncatalyzed reaction mechanism.
An important point to note is that catalysts are not reactants in that they are not consumed by the reaction and turn into products. Instead they show up as intermediates in the elementary reactions and are regenerated by the end of the reaction. This allows them to be reused many times over and participate in catalyzing multiple molecules. When we add the several elementary reactions to create the overall reaction the catalyst is eliminated from the reaction because it an intermediate (showing that it is neither a reactant or product). This also means that a catalyzed and uncatlayzed reaction will have the same chemical equation; except the rates will differ.
A common example of a catalyzed reaction is the decomposition of ozone in the upper atmosphere by chlorofluorocarbons (CFCs). Ozone is a molecule with the chemical formula O3 which is a toxic gas to humans, but when present in Earth's upper atmosphere help shield us from the sun's more damaging high energy light. Ozone can degrade into diatomic oxygen by the reaction: O3 + O → 2 O2, but the activation energy for this reaction is high and this reaction is fairly slow so ozone does not decompose easily. However the reaction can be sped up with the addition of a chlorine atom catalyst. Chlorine atoms can be found in the upper atmosphere too when CFCs (collectively called freon) such as dichlorodifluoromethane degrade because of absorption of high energy light from the sun. The liberated chlorine can react with ozone in a two-step catalyzed reaction:
Cl + O3 → ClO + O2
ClO + O → Cl + O2.
And if we add these two elementary reaction together they give the original uncatalyzed reaction and the chlorine disappears from the equation as an intermediate (which is what a catalyst should do). This catalyzed reaction mechanism has a lower activation energy which causes ozone to decompose much faster. This depletes our atmosphere's ozone layer and allows a greater percentage of damaging high energy light to make it to the Earth's surface. These CFCs were commonly used as refrigerants in the latter half of the 20th century and created the famous "ozone hole" which caused them to be gradually phased out. But being catalysts which don't get consumed in the reaction which decomposes ozone, these CFCs still linger in the atmosphere causing damage and take many years to finally stop being a threat even after they were banned. So that's an example of an unwanted catalyzed reaction.
Hope that helps.(16 votes)
What is the rationale behind how surface area increases when the substance is broken down into smaller parts?
I understand that the surface area to volume ratio would be higher with a smaller fragment (and how that would increase the rate of the reaction), but I find the terminology to be confusing as the fragments do occupy less space than the larger whole.(3 votes)
When you break something in half, the surfaces where those two pieces were connected are now new surface areas for the two objects. As you continue breaking down objects into smaller pieces, this continues exposing new surface area.(5 votes)
when does the concentration just effect the rate of reaction, and when does it also effect the volume of products?(2 votes)- For most reactions increasing the concentration of the reactants will increase the rate of the reaction. The exception to this are zero order reactions where the concentrations of the reactants have no bearing on the reaction rate.
The volume of products depends on how much of the products we produce in a reaction. The extent of product creation is governed by its equilibrium, not its reaction rate. Reaction rate only tells us how fast a reaction progresses, but equilibrium tells us how much of the reactants have been consumed and converted to products.
Hope that helps.(3 votes)
Video transcript
- [Instructor] There are several factors that can affect the rate of a reaction. One factor is the
concentration of a reactant. Most chemical reactions proceed faster when the concentration of one
of the reactants is increased. For example, let's look at
the reaction of solid zinc with hydrochloric acid to
form an aqueous solution of zinc chloride and hydrogen gas. Let's say we put a piece of zinc metal. So go ahead and draw in the
piece of zinc metal in here in a flask that contains
three molar hydrochloric acid. So in our flask here, we have three molar, a solution of three
molar hydrochloric acid. As the reaction proceeds,
hydrogen gas is formed. So we could monitor the
rate of this reaction by observing the amount
of hydrogen gas bubbles that are coming out of the flask. Let's say that we repeat the experiment. This time instead of using
three molar hydrochloric acid we're gonna use six
molar hydrochloric acid. So we've increased the concentration of our hydrochloric acid solution. This time, when we add
our piece of solid zinc to our six molar
hydrochloric acid solution, we would observe more hydrogen gas bubbles coming out of our flask. So we have increased the concentration of one of our reactants hydrochloric acid. And we've observed an increase
in the rate of the reaction. As the concentration of
hydrochloric acid increases, there are more acid particles to collide with the piece of zinc. And therefore, as the concentration of hydrochloric acid goes up, the frequency of collisions increase and the rate of the reaction increases. Let's use the same reaction to talk about another factor that affects the rate of reaction. And that factor is surface area. You've already reacted
a piece of solid zinc with hydrochloric acid. We could use three molar
hydrochloric acid again. And we saw some bubbles come
off of the piece of zinc indicating that hydrogen gas was produced. This time let's try
breaking the piece of zinc into smaller pieces. So instead of using one
large piece of zinc, here we have a bunch of
small pieces of zinc. And if we were to do the experiment again, with three molar hydrochloric acid, this time we would see more
bubbles of hydrogen gas coming off of those small pieces of zinc. So we have increased the
surface area of the solid and we observed an increase
in the rate of the reaction. When we had only one piece of zinc, the rate of the reaction was limited by the surface area of this one piece. Therefore by breaking it
up into smaller pieces, we were able to the rate of the reaction. Temperature is another factor that can affect the rate of a reaction. So let's say on the left,
we have a glow stick in a flask that contains some cold water. And the glow stick, let's say the glow stick
is already glowing here. So let's draw in this color
here for our glow stick. The glow from a glow stick comes from a chemical reaction. And if we were to
increase the temperature, so let's say we heated
this flask on the left. So we increase the temperature. We would observe the glow to get stronger. So let's go ahead and draw in here, a more vibrant glow coming
from our glow stick. So increasing the temperature must have increased the
rate of the reaction. The reason increasing the temperature increases the rate of
the reaction in general, is because increase in temperature means the molecules are moving faster. And therefore the molecules
are colliding with each other with greater frequency
and with greater force which increases the rate of the reaction. The catalyst is another factor that can affect the rate of a reaction. Let's look at the balanced
equation for the decomposition of hydrogen peroxide, which
turns into water and oxygen. And let's say in our flask on the left, we have a solution of hydrogen peroxide. The hydrogen peroxide is
decomposing at room temperature, but the reaction proceeds so slowly that we don't see it even happening. We can speed up the reaction
by adding a catalyst. Let's say we have an aqueous
solution of potassium iodide in our beaker here. And we pour the solution
of potassium iodide into our flask containing
the hydrogen peroxide. The addition of the
iodide ion as a catalyst, causes the decomposition
of hydrogen peroxide to occur very quickly. And we would see a huge plume of gas come out of the reaction flask. So the addition of a catalyst, in this case it was the iodide, an ion, increased the rate of the reaction. A catalyst increases the rate of reaction by effecting the kinds of collisions that occur between particles. And a catalyst increases
the rate of reaction without being used up.