Living organisms rely on energy to stay alive and to carry out day-to-day activities. Potential energy in the form of chemical bonds is converted to kinetic energy which produces movement. Two laws about energy are important concepts to grasp in biochemistry, i.e., the first and the second law of thermodynamics.

First Law of Thermodynamics

The first law of thermodynamics says that energy cannot be created nor destroyed but simply moves from one state to the next. For example, chemical energy in ATP results in motion (kinetic energy) of an organism.

Second Law of Thermodynamics

The second law says that when energy is transferred from one form to another, some of the usable energy (energy available to do work) is lost to the surroundings. Therefore the amount of useable energy in a system decreases with time. Energy that is not usable is random and disorganized. Entropy (S) is the measure of this disorder. Disorganized energy has a high entropy. The universe gradually tends toward higher entropy. With this in mind, it is important to note that no system is 100% efficient. Energy will always being lost. This typically occurs in the form of heat.

Free Energy

Chemical reactions that take place in living organisms is referred to as their metabolism. There are two types of metabolic pathways, anabolic and catabolic,

  1. Anabolic: Pathways involving synthesis of biomolecules
  2. Catabolic: Pathways involving breakdown of biomolecules

Metabolic reactions involve the breakdown and forming of chemical bonds. The total energy in the chemical bonds in a system is know as enthalpy (H).

Enthalpy and entropy are related to a third type of energy called Gibbs free energy (G). G is the the total amount of energy available to do useful work.

Enthalpy (H) = Usable energy (G) + unusable energy (S)

H = G + TS, where T = temperature in kelvin (T is added since the problem of entropy gets worse with increase in temperature)

This equation can be rearranged to, G = H – TS

The equation G = H-TS can be used to predict whether a chemical reaction will require input of energy or will release energy. To predict changes that occur between the initial and final stage of a chemical reaction, we use the expression,

ΔG = ΔH – TΔS

Where ΔG and ΔH are expressed in kilojoules/mole and ΔS is expressed in kilojoules/degree

An exergonic reaction is one that releases energy and is said to be spontaneous. In this reaction, free energy moves from high to low, producing a negative ΔG.

An endergonic reaction is one that gains free energy and is non-spontaneous. In this reaction, free energy moves from low to high, producing a positive ΔG.

Many biochemical reactions are endergonic which means they cannot take place without the input of energy. To ensure that they are able to proceed, endergonic reactions are coupled with exergonic reactions. Together, they have a an overall ΔG that is negative. Here is an example,

Coupled reaction

Energy Transfer in Redox Reactions

Energy can be transferred through the transfer of phosphate groups as you saw in the above equation. However, energy can also be transferred through the transfer of electrons. This occurs during oxidation-reduction, or redox reactions. Substances that are oxidized, lose electrons or hydrogen atoms and hence give up energy. Substances that are reduced, accept electrons or hydrogen atoms, and hence gain energy. Redox reactions often occurs in a series where electrons or hydrogen pass from one molecule to another. Examples of hydrogen acceptor molecules that you will learn about in biochemistry are, nicotinamide adenine dinucleotide (NAD+), nicotinamide adenine dinucleotide phosphate (NADP+), and flavin adenine dinucleotide (FAD).

Courtney Simons
Administrator
Courtney Simons is a food science professor. He holds a BS degree in food science and a Ph.D. in cereal science from North Dakota State University.
Courtney Simons on EmailCourtney Simons on FacebookCourtney Simons on Linkedin