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Course: AP®︎/College Chemistry > Unit 6
Lesson 1: Endothermic and exothermic processesRepresenting endothermic and exothermic processes using energy diagrams
A physical or chemical process can be represented using an energy diagram, which shows how the potential energy of the initial state relates to the potential energy of the final state. If the initial state has a lower potential energy than the final state, the process is endothermic. If the initial state has a higher potential energy than the final state, the process is exothermic. Created by Jay.
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what does the peak in the diagram of endothermic reaction meas?
As the energy is rising from 50kj/mol, it didn't stop when it reached 100.Instead, it went all the way up to 150 then dropback to 100.
Why is that?(4 votes)
The peaks in energy diagrams for both endothermic and exothermic reaction energy diagrams are known as the transition state or the activation complex.
In a reaction, any reaction, the same general trend occurs. First the bonds of the reactants are broken which requires an input of energy to be put into the reaction. This is why both diagrams have a large slope immediately following the reactants. This energy required to break all the reactant bonds is referred to as the activation energy and is represented as the height from the reactant's energy level to the transition state's energy. Transition states are highly energetic and unstable chemical species which are always higher than both the reactant and products in an energy diagram for any reaction. In the transition state chemical bonds are simultaneously being broken and formed. Essentially the reactants are transitioning to the products in this high energy state. Finally when the new product bonds are formed energy is released since the formation of bonds lowers energy. And this is represented as a large downward slope down to the products.
Hope that helps.(12 votes)
- From the last video it mentions that Delta H is equal to q only when pressure is constant. This example doesn't specify if pressure is constant. Therefore how can you say that Delta H equals Delta E. The first law equation, DeltaE=q+w becomes deltaE=q assuming no work is done. But it isn't q sub p. Without the sub p is it still change in enthalpy?(5 votes)
- If a reaction happens in an open-air situation, as we have here with an open beaker, then we assume the pressure is constant since the air pressure of a room remains more or less unchanged over the course of the reaction.
Hope that helps.(3 votes)
- what's a "mol". i already somewhat know what a joule is, but what's a mol?(2 votes)
- The mole (abbreviated mol) is a unit to quantify the amount of things there is. A mole is equal to an Avogadro's number of things (Avogadro's number being a mathematical constant equal to ~6.022 x 10^(23)). It's similar to how a dozen of a thing means 12 of that thing.
Sal does a video discussing this: https://www.khanacademy.org/science/ap-chemistry-beta/x2eef969c74e0d802:atomic-structure-and-properties/x2eef969c74e0d802:moles-and-molar-mass/v/the-mole-and-avogadro-s-number
Hope that helps.(3 votes)
- in this case, how do we know that the work done is 0?(1 vote)
- With constant pressure processes, like this one, work done by the system (the reaction) is not 0.
Through some math, we can show that the change in enthalpy for a reaction at constant pressure is equal to the heat exchanged with the surroundings, or ΔH = q. So enthalpy at constant pressure is the same as heat, but enthalpy is different from the change in internal energy, ΔU. The change in internal energy is the sum of heat and work, or ΔU = q + w. The type of work done by a reaction is pressure-volume work, or the force exerted by the volume change of gases, which is w = -PΔV. If the reaction produces gases which increases the volume of the surroundings, then it has negative work. If the reaction consumes gases, then the volume decreases and the work is positive.
For chemical reactions that do not exchanged much work with the surroundings-that is, those that do not cause a large change in volume as they occur-ΔH and ΔU are nearly identical in value. For chemical reactions that produce or consume large amounts of gas, and therefore result in large volume changes, ΔH and ΔU can be slightly different in value. So work is still being done in a constant pressure reaction, it’s just the amount is so minor compared to the heat exchanged. The change to the atmosphere from a single chemical reaction is insignificant most times. If the reaction was being done at constant volume though, work would be 0 because ΔV in w = -PΔV is 0.
For the reactions shown in the video, Jay states they occur in beakers so they are open to the atmosphere and therefore have constant pressure so the ΔH = q equation is valid. Here we’re just calculating ΔH, but it is making no mention of work. Again, it’s still present and we could consider it if we were calculating ΔU, but we’re not doing that because ΔH is the objective.
Hope that helps.(2 votes)
- but from the last video delta H was equals to heat at constant pressure. And now its equals to internal energy. and we use q to solve for the internal energy?.(1 vote)
I'm pretty sure this is because the work done on the system in this example is 0. Meaning that the delta H would be equal to delta E.(1 vote)
Video transcript
- [Instructor] Let's
say we run an experiment to determine if a reaction
is endo or exothermic. For our hypothetical reaction,
A reaction would B to form C. And let's say this reaction
takes place in aqueous solution in a beaker. We can define our system as
the reactants and products that make up our chemical reaction, and everything else is
part of the surroundings, for example, the water and also the beaker the reaction is taking place in. Let's say we run the reaction and we put our hand on the beaker and we feel that the beaker is warm. If the beaker is warm since the beaker is part
of the surroundings, energy must've been
transferred from the system to the surrounding. So heat flowed from the
system to the surroundings. And that's an example of
an exothermic reaction. So Delta H is negative. We can determine the amount
of energy that flowed from the system to the surroundings by looking at the energy profile for our hypothetical reaction. In an energy profile, potential
energy is in the y-axis in kilojoules per mole
and reaction to progress is on the x-axis. So as we move to the right on the x-axis, the reaction is occurring. Our reactants, which are A and B have a certain amount of potential energy. So that's right here
on our energy profile. So that part represents the
energy of our reactants. Our reactants react
together to form our product which is C and that's at the very end. So over here, this line
represents the energy the potential energy of our products. Notice how the potential
energy of our reactants is higher than the potential
energy of the product. So if we were to find
the change in energy, that would be the final minus the initial. So the energy of the
products minus the energy of the reactance for this energy profile, the energy of the products is
about 50 kilojoules per mole, and the potential energy
of our reactance is at 100. So this would be 50 minus 100 which is equal to negative
50 kilojoules per mole. So on our energy profile,
we could show Delta E which would be this difference right here. So that represents Delta E and the change in energy Delta
E is also equal to the change in the enthalpy Delta H for this reaction. So we know that the change in enthalpy is equal to negative 50
kilojoules per moles. Let's go ahead and plug that
in over here on the left. So by feeling the outside of the beaker, we knew that the reaction was exothermic, but the energy profile
allowed us to figure out how much energy was
transferred from the system to the surroundings. So for an energy profile, when
the energy of the reactants is higher than the energy of the products, this is the energy profile
for an exothermic reaction. Let's say we ran a similar reaction where A plus B turned into C, but this time when we felt the beaker, the beaker felt cool to the touch. If that's the case, it's because energy was being transferred from the
surroundings to the system. And since these surroundings
was losing energy, that's why the beaker felt cool. So heat flows from the
surroundings to the system, and this occurs in an endothermic reaction and the change in entropy,
Delta H is positive for an endothermic process. When we look at the energy profile for an endothermic reaction. The energy of the reactants. So go ahead and write reactants in here. The energy of the reactants is lower than the energy of the products. So this time, if we find Delta E that would be the energy of our products minus the energy of our reactants. And the energy of our products is about 100 kilojoules per mole and the energy of our
reactants is about 50. So let's say it's 100 minus 50, which would be positive
50 kilojoules per mole. So on our diagram if we represent Delta E, that would be this difference
here on the energy profile. And once again, Delta E
is equal to the change in the enthalpy Delta H for the reaction. So Delta H for this hypothetical reaction is positive 50 kilojoules per mole. Since Delta H is positive, we know that energy was
transferred from the surroundings to the system. And that's the reason why the products have a higher potential
energy than the reactants in our energy diagram.