Energy

Energy is a quantitative property that, when transferred to an object, can perform work. In other words, if you transfer energy to an object, that object now has the ability to move against a force (such as gravity).

There are many types of energy including (but not limited to): Chemical, electrical, mechanical, light, radiation, and thermal energy.

For the sake of this course, we will divide energy into two categories:

Kinetic Energy- the energy of motion (recall: Kinetic Molecular Theory, KMT)
Potential Energy- stored energy/energy that is available, but not yet released.

The classic illustration of potential energy is seen in this image. Above, the picture on top has rock at the top of a hill. This rock has potential energy; with a slight amount of encouragement, it has the ability to change that stored energy into kinetic energy (movement), and roll down the hill.

In comparative terms, the picture above has more potential energy than the picture below.

To get a better grasp of this concept, we must contemplate The First Law of Thermodynamics, which states:

Energy cannot be created nor destroyed, only converted from one type to another or transferred from one object to another.

Chemical bonds are a major type of potential energy in this course. Breaking a bond releases chemical energy that can be harnessed to do work. Bond Energy is the minimum amount of energy that is required to break a particular type of bond (essentially, how much energy is stored in a particular bond). To the right is a table of average bond energies of common bonds found in organic molecules.

Potential Energy Diagrams

Used to illustrate the changes in potential energy that take place during a chemical reaction
The activation energy (E_A) is the amount of energy needed to strain and break the reactant’s bonds
The transition state is when the reactant bonds are breaking and the bonds between products are forming

Recall: Enzymes lower activation energy (E_A)

The Second Law of Thermodynamics states that in every energy transfer or conversion, some of the useful energy (called Gibbs Free Energy, G) is lost. A great example of this is an incandescent light bulb- while the electrical energy is transformed into light energy, the majority of the energy is released as heat!

As we learned in the First Law of Thermodynamics, the energy is not destroyed, but rather increases the entropy of the surroundings.

Entropy (S) is a measure of the tendency of a system to become unorganized or disordered. Entropy increases when disorder increases.

Examples:

throwing a deck of cards into the air and letting them land scattered
a campfire: The solid wood burns and becomes ash, smoke and gases, all of which spread energy outwards more easily than the solid fuel

Endergonic Reactions

If the amount of energy being absorbed in breaking reactant bonds is greater than the energy released in the formation of product bonds, then the reaction is called endothermic or endergonic (ΔG is positive)
The products contain more potential energy than the reactants and are considered less stable

Exergonic Reactions

If the amount of energy being absorbed in breaking reactant bonds is less than the energy released in the formation of product bonds, then the reaction is called exothermic or exergonic (ΔG is negative)
The products contain less potential energy than the reactants and are considered more stable

Textbook Practice

p.126-140 #1-4, 7

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