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1.4.6 Reaction Energy and Heat

Learning Objectives

By the end of this section, you will be able to:

  • Describe endergonic, exergonic, endothermic, and exothermic chemical reactions
  • Relate enthalpy change to the abovementioned reactions
  • Calculate enthalpy change in a reaction using bond energy
  • Describe the concept of Gradient
Some chemical reactions require an input of energy. The formation of carbonic acid (H2CO3) from CO2 and H2O requires an input of energy and is an endergonic reaction. However, the reverse reaction, where CO2 and H2O are formed from the breakdown of H2CO3 releases energy. Chemical reactions that release energy are called exergonic reactions.Every chemical reaction will transform energy. The first law of thermodynamics states that energy can neither be created nor destroyed. The energy required or released from chemical reactions is not created nor lost by cells. The energy already existed in the universe before the reaction occurred. So, where does the energy come from for endergonic chemical reactions to proceed?Outside of cells, the ambient heat provides the energy for chemical reactions to occur. The formation of carbonic acid under standard conditions (25°C, 1 atm) takes 60–90 seconds using ambient heat. That may seem like a short time, but cells require reactions to occur much faster to support life. Cellular reactions make use of catalysts: chemicals that increase the rate of chemical reactions without being changed by the reaction. Proteins called enzymes are cellular catalysts: they increase the rate of chemical reactions in cells.

Enzyme-substrate interactions resemble a lock and key. B. Enzymes decrease the activation energy of chemical reactions to increase reaction rate.
Figure 1.4.6.1 Enzyme-substrate interactions resemble a lock and key.

How do enzymes speed up the rate of chemical reactions in cells? The reactants in an enzyme-catalyzed reaction have a special name: substrates. The conversion of substrates to products requires an energy investment. This energy is called activation energy. Activation energy is the energy required to weaken and rearrange the bonds of the substrates to form the products. Enzymes bind substrates at a specific location on the protein called the active site and the active site precisely fits the shape of the substrate(s), like a lock fits a key. The ability of enzymes to only bind its substrates is called specificity. At the active site, the enzyme positions the substrate(s) optimally for the conversion into product(s), thereby decreasing the time needed for the chemical reaction to occur.

Enzyme-catalyzed reactions follow the first law of thermodynamics, like all matter in the universe. The energy for chemical reactions performed by enzymes must come from somewhere. In cells, many chemical reactions are powered by the consumption of chemical sources of energy such as ATP. The energy invested is used to rearrange the chemical bonds within the substrates to form the products. Some of the energy will be converted and dissipated as heat energy (more in the next paragraph). The point is that the energy for chemical reactions already exists in the universe before it is converted to another form of energy. The energy is never destroyed: it simply takes a different form.

For example, when you eat breakfast and run for the bus to come to class, you are transforming the potential energy in the food that you have eaten into kinetic energy while running for the bus. Inside of your cells, chemical reactions rearrange the matter in the food you eat to release the potential energy from those molecules and your muscle cells convert the energy released into kinetic energy (running). However, energy transformations are never 100% efficient. This means that you can never recover 100% of the potential energy stored in the food you eat to convert it to kinetic energy. This is an application of the second law of thermodynamics. A precise statement of the second law of thermodynamics is beyond the scope of this course. However, we can apply it to an understanding of energy transformation by stating that transformations of energy are never 100% efficient because some energy will be dissipated as heat energy. Every single energy transformation or chemical reaction within and outside of cells will result in some of the energy invested or released to be converted into heat energy. The heat energy is not “lost” but it cannot be recovered or used to do additional work. This is not a good nor a bad thing: it is simply the way that our universe works!

Animals that maintain body temperatures different than their surroundings benefit from the second law of thermodynamics, including humans. When we convert the potential energy of food we have eaten into kinetic energy during muscle contraction, every single one of those chemical reactions dissipates some of the energy as heat. The released heat helps to maintain our internal body temperatures at a constant 37°C, regardless of the temperatures of our surroundings.

Thermal Energy, Temperature, and Heat

Thermal energy is kinetic energy associated with the random motion of atoms and molecules. Temperature is a quantitative measure of “hot” or “cold.” When the atoms and molecules in an object are moving or vibrating quickly, they have a higher average kinetic energy (KE), and we say that the object is “hot.” When the atoms and molecules are moving slowly, they have lower average KE, and we say that the object is “cold” (Figure 1.4.6.2). Assuming that no chemical reaction or phase change (such as melting or vaporizing) occurs, increasing the amount of thermal energy in a sample of matter will cause its temperature to increase. And, assuming that no chemical reaction or phase change (such as condensation or freezing) occurs, decreasing the amount of thermal energy in a sample of matter will cause its temperature to decrease.

Two molecular drawings are shown and labeled a and b. Drawing a is a box containing fourteen red spheres that are surrounded by lines indicating that the particles are moving rapidly. This drawing has a label that reads “Hot liquid.” Drawing b depicts another box of equal size that also contains fourteen spheres, but these are blue. They are all surrounded by smaller lines that depict some particle motion, but not as much as in drawing a. This drawing has a label that reads “Cold liquid.”
Figure 1.4.6.2 (a) The molecules in a sample of hot water move more rapidly than (b) those in a sample of cold water.

LINK TO LEARNING

Click on this interactive simulation to view the effects of temperature on molecular motion.

Most substances expand as their temperature increases and contract as their temperature decreases. This property can be used to measure temperature changes, as shown in Figure 1.4.6.3. The operation of many thermometers depends on the expansion and contraction of substances in response to temperature changes.

A picture labeled a is shown as well as a pair of drawings labeled b. Picture a shows the lower portion of an alcohol thermometer. The thermometer has a printed scale to the left of the tube in the center that reads from negative forty degrees at the bottom to forty degrees at the top. It also has a scale printed to the right of the tube that reads from negative thirty degrees at the bottom to thirty five degrees at the top. On both scales, the volume of the alcohol in the tube reads between nine and ten degrees. The two images labeled b both depict a metal strip coiled into a spiral and composed of brass and steel. The left coil, which is loosely coiled, is labeled along its upper edge with the 30 degrees C and 10 degrees C. The end of the coil is near the 30 degrees C label. The right hand coil is much more tightly wound and the end is near the 10 degree C label.
Figure 1.4.6.3 (a) In an alcohol or mercury thermometer, the liquid (dyed red for visibility) expands when heated and contracts when cooled, much more so than the glass tube that contains the liquid. (b) In a bimetallic thermometer, two different metals (such as brass and steel) form a two-layered strip. When heated or cooled, one of the metals (brass) expands or contracts more than the other metal (steel), causing the strip to coil or uncoil. Both types of thermometers have a calibrated scale that indicates the temperature. (credit a: modification of work by “dwstucke”/Flickr)

LINK TO LEARNING

The following demonstration allows one to view the effects of heating and cooling a coiled bimetallic strip.

 

Heat (q) is the transfer of thermal energy between two bodies at different temperatures. Heat flow (a redundant term, but one commonly used) increases the thermal energy of one body and decreases the thermal energy of the other. Suppose we initially have a high temperature (and high thermal energy) substance (H) and a low temperature (and low thermal energy) substance (L). The atoms and molecules in H have a higher average KE than those in L. If we place substance H in contact with substance L, the thermal energy will flow spontaneously from substance H to substance L. The temperature of substance H will decrease, as will the average KE of its molecules; the temperature of substance L will increase, along with the average KE of its molecules. Heat flow will continue until the two substances are at the same temperature (Figure 1.4.6.4).

Three drawings are shown and labeled a, b, and c, respectively. The first drawing labeled a depicts two boxes, with a space in between and the pair is captioned “Different temperatures.” The left hand box is labeled H and holds fourteen well-spaced red spheres with lines drawn around them to indicate rapid motion. The right hand box is labeled L and depicts fourteen blue spheres that are closer together than the red spheres and have smaller lines around them showing less particle motion. The second drawing labeled b depicts two boxes that are touching one another. The left box is labeled H and contains fourteen maroon spheres that are spaced evenly apart. There are tiny lines around each sphere depicting particle movement. The right box is labeled L and holds fourteen purple spheres that are slightly closer together than the maroon spheres. There are also tiny lines around each sphere depicting particle movement. A black arrow points from the left box to the right box and the pair of diagrams is captioned “Contact.” The third drawing labeled c, is labeled “Thermal equilibrium.” There are two boxes shown in contact with one another. Both boxes contain fourteen purple spheres with small lines around them depicting moderate movement. The left box is labeled H and the right box is labeled L.
Figure 1.4.6.4 (a) Substances H and L are initially at different temperatures, and their atoms have different average kinetic energies. (b) When they contact each other, collisions between the molecules result in the transfer of kinetic (thermal) energy from the hotter to the cooler matter. (c) The two objects reach “thermal equilibrium” when both substances are at the same temperature and their molecules have the same average kinetic energy.

LINK TO LEARNING

Click on the PhET simulation to explore energy forms and changes. Visit the Energy Systems tab to create combinations of energy sources, transformation methods, and outputs. Click on Energy Symbols to visualize the transfer of energy.

Matter undergoing chemical reactions and physical changes can release or absorb heat. A change that releases heat is called an exothermic process. For example, the combustion reaction that occurs when using an oxyacetylene torch is an exothermic process—this process also releases energy in the form of light as evidenced by the torch’s flame (Figure 1.4.6.5). A reaction or change that absorbs heat is an endothermic process. A cold pack used to treat muscle strains provides an example of an endothermic process. When the substances in the cold pack (water and a salt like ammonium nitrate) are brought together, the resulting process absorbs heat, leading to the sensation of cold.

Two pictures are shown and labeled a and b. Picture a shows a metal railroad tie being cut with the flame of an acetylene torch. Picture b shows a chemical cold pack containing ammonium nitrate.
Figure 1.4.6.5 (a) An oxyacetylene torch produces heat by the combustion of acetylene in oxygen. The energy released by this exothermic reaction heats and then melts the metal being cut. The sparks are tiny bits of the molten metal flying away. (b) A cold pack uses an endothermic process to create the sensation of cold. (credit a: modification of work by “Skatebiker”/Wikimedia commons)
Historically, energy was measured in units of calories (cal). A calorie is the amount of energy required to raise one gram of water by 1 degree C (1 kelvin). However, this quantity depends on the atmospheric pressure and the starting temperature of the water. The ease of measurement of energy changes in calories has meant that the calorie is still frequently used. The Calorie (with a capital C), or large calorie, commonly used in quantifying food energy content, is a kilocalorie. The SI unit of heat, work, and energy is the joule. A joule (J) is defined as the amount of energy used when a force of 1 newton moves an object 1 meter. It is named in honor of the English physicist James Prescott Joule. One joule is equivalent to 1 kg m2/s2, which is also called 1 newton–meter. A kilojoule (kJ) is 1000 joules. To standardize its definition, 1 calorie has been set to equal 4.184 joules.

Enthalpy

Chemists ordinarily use a property known as enthalpy (H) to describe the thermodynamics of chemical and physical processes. Enthalpy is also a state function. Enthalpy values for specific substances cannot be measured directly; only enthalpy changes (ΔH) for chemical or physical processes can be determined.

The following conventions apply when using ΔH:

  • A negative value of an enthalpy change, ΔH < 0, indicates an exothermic reaction; a positive value, ΔH > 0, indicates an endothermic reaction. If the direction of a chemical equation is reversed, the arithmetic sign of its ΔH is changed (a process that is endothermic in one direction is exothermic in the opposite direction).

  • Chemists use a thermochemical equation to represent the changes in both matter and energy. In a thermochemical equation, the enthalpy change of a reaction is shown as a ΔH value following the equation for the reaction. This ΔH value indicates the amount of heat associated with the reaction involving the number of moles of reactants and products as shown in the chemical equation. For example, consider this equation:

     

H2(g)+12O2(g)⟶H2O(l)ΔH=−286kJ

This equation indicates that when 1 mole of hydrogen gas and 1/2 mole of oxygen gas at some temperature and pressure change to 1 mole of liquid water at the same temperature and pressure, 286 kJ of heat are released to the surroundings. If the coefficients of the chemical equation are multiplied by some factor, the enthalpy change must be multiplied by that same factor (ΔH is an extensive property):

(two-fold increase in amounts)2H2(g)+O2(g)⟶2H2O(l)ΔH=2×(−286kJ)=−572kJ(two-fold decrease in amounts)12H2(g)+14O2(g)⟶12H2O(l)ΔH=12×(−286kJ)=−143kJThe enthalpy change of a reaction depends on the physical states of the reactants and products, so these must be shown. For example, when 1 mole of hydrogen gas and 1/2 mole of oxygen gas change to 1 mole of liquid water at the same temperature and pressure, 286 kJ of heat are released. If gaseous water forms, only 242 kJ of heat are released.

 

H2(g)+12O2(g)⟶H2O(g)ΔH=−242kJ 

Bond Strength

A bond’s strength describes how strongly each atom is joined to another atom, and therefore how much energy is required to break the bond between the two atoms.

It is essential to remember that energy must be added to break chemical bonds (an endothermic process), whereas forming chemical bonds releases energy (an exothermic process). In the case of H2, the covalent bond is very strong; a large amount of energy, 436 kJ, must be added to break the bonds in one mole of hydrogen molecules and cause the atoms to separate:

H2(g)⟶2H(g)bond energy=436kJConversely, the same amount of energy is released when one mole of H2 molecules forms from two moles of H atoms:

2H(g)⟶H2(g)bond energy=−436kJ 

Bond Strength: Covalent Bonds

Stable molecules exist because covalent bonds hold the atoms together. We measure the strength of a covalent bond by the energy required to break it, that is, the energy necessary to separate the bonded atoms. Separating any pair of bonded atoms requires energy. The stronger a bond, the greater the energy required to break it.

The energy required to break a specific covalent bond in one mole of gaseous molecules is called the bond energy or the bond dissociation energy. The bond energy for a diatomic molecule, DX–Y, is defined as the standard enthalpy change for the endothermic reaction:

XY(g)⟶X(g)+Y(g)DX−Y=ΔH°For example, the bond energy of the pure covalent H–H bond, DH–H, is 436 kJ per mole of H–H bonds broken:

H2(g)⟶2H(g)DH−H=ΔH°=436kJMolecules with three or more atoms have two or more bonds. The sum of all bond energies in such a molecule is equal to the standard enthalpy change for the endothermic reaction that breaks all the bonds in the molecule. For example, the sum of the four C–H bond energies in CH4, 1660 kJ, is equal to the standard enthalpy change of the reaction:

A reaction is shown with Lewis structures. The first structure shows a carbon atom single bonded to four hydrogen atoms with the symbol, “( g )” written next to it. A right-facing arrow points to the letter “C” and the symbol “( g ),” which is followed by a plus sign. Next is the number 4, the letter “H” and the symbol, “( g ).” To the right of this equation is another equation: capital delta H superscript degree symbol equals 1660 k J.

The average C–H bond energy, DC–H, is 1660/4 = 415 kJ/mol because there are four moles of C–H bonds broken per mole of the reaction. Although the four C–H bonds are equivalent in the original molecule, they do not each require the same energy to break; once the first bond is broken (which requires 439 kJ/mol), the remaining bonds are easier to break. The 415 kJ/mol value is the average, not the exact value required to break any one bond.

The strength of a bond between two atoms increases as the number of electron pairs in the bond increases. Generally, as the bond strength increases, the bond length decreases. Thus, we find that triple bonds are stronger and shorter than double bonds between the same two atoms; likewise, double bonds are stronger and shorter than single bonds between the same two atoms. Average bond energies for some common bonds appear in Table 1.4.6.1, and a comparison of bond lengths and bond strengths for some common bonds appears in Table 1.4.6.2. When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group. For example, C–F is 439 kJ/mol, C–Cl is 330 kJ/mol, and C–Br is 275 kJ/mol.

Table 1.4.6.1 Bond Energies

Bond Bond Energy Bond Bond Energy Bond Bond Energy
H–H 436 C–S 260 F–Cl 255
H–C 415 C–Cl 330 F–Br 235
H–N 390 C–Br 275 Si–Si 230
H–O 464 C–I 240 Si–P 215
H–F 569 N–N 160 Si–S 225
H–Si 395 N=N 418 Si–Cl 359
H–P 320 N≡N 946 Si–Br 290
H–S 340 N–O 200 Si–I 215
H–Cl 432 N–F 270 P–P 215
H–Br 370 N–P 210 P–S 230
H–I 295 N–Cl 200 P–Cl 330
C–C 345 N–Br 245 P–Br 270
C=C 611 O–O 140 P–I 215
C≡C 837 O=O 498 S–S 215
C–N 290 O–F 160 S–Cl 250
C=N 615 O–Si 370 S–Br 215
C≡N 891 O–P 350 Cl–Cl 243
C–O 350 O–Cl 205 Cl–Br 220
C=O 741 O–I 200 Cl–I 210
C≡O 1080 F–F 160 Br–Br 190
C–F 439 F–Si 540 Br–I 180
C–Si 360 F–P 489 I–I 150
C–P 265 F–S 285
Table 1.4.6.2 Average Bond Lengths and Bond Energies for Some Common Bonds
Bond Bond Length (Å) Bond Energy (kJ/mol)
C–C 1.54 345
C=C 1.34 611
C≡C 1.20 837
C–N 1.43 290
C=N 1.38 615
C≡N 1.16 891
C–O 1.43 350
C=O 1.23 741
C≡O 1.13 1080
The bond energy is the difference between the energy minimum (which occurs at the bond distance) and the energy of the two separated atoms. This is the quantity of energy released when the bond is formed. Conversely, the same amount of energy is required to break the bond. For a H2 molecule shown, at the bond distance of 74 pm the system is 7.24 × 10−19 J lower in energy than the two separated hydrogen atoms. This may seem like a small number. However, as we will learn in more detail later, bond energies are often discussed on a per-mole basis. For example, it requires 7.24 × 10−19 J to break one H–H bond, but it takes 4.36 × 105 J to break 1 mole of H–H bonds. We can find many of these bonds in a variety of molecules, and this table provides average values. For example, breaking the first C–H bond in CH4 requires 439.3 kJ/mol, while breaking the first C–H bond in H–CH2C6H5 (a common paint thinner) requires 375.5 kJ/mol.

We can use bond energies to calculate approximate enthalpy changes for reactions where enthalpies of formation are not available. Calculations of this type will also tell us whether a reaction is exothermic or endothermic. An exothermic reaction (ΔH negative, heat produced) results when the bonds in the products are stronger than the bonds in the reactants. An endothermic reaction (ΔH positive, heat absorbed) results when the bonds in the products are weaker than those in the reactants.

The enthalpy change, ΔH, for a chemical reaction is approximately equal to the sum of the energy required to break all bonds in the reactants (energy “in”, positive sign) plus the energy released when all bonds are formed in the products (energy “out,” negative sign). This can be expressed mathematically in the following way:

ΔH=ƩDbonds broken−ƩDbonds formedIn this expression, the symbol Ʃ means “the sum of” and D represents the bond energy in kilojoules per mole, which is always a positive number. The bond energy is obtained from a table (like Table 1.4.6.2) and will depend on whether the particular bond is a single, double, or triple bond. Thus, in calculating enthalpies in this manner, it is important that we consider the bonding in all reactants and products. Because D values are typically averages for one type of bond in many different molecules, this calculation provides a rough estimate, not an exact value, for the enthalpy of reaction.

Consider the following reaction:

H2(g)+Cl2(g)⟶2HCl(g)or

H–H(g)+Cl–Cl(g)⟶2H–Cl(g)To form two moles of HCl, one mole of H–H bonds and one mole of Cl–Cl bonds must be broken. The energy required to break these bonds is the sum of the bond energy of the H–H bond (436 kJ/mol) and the Cl–Cl bond (243 kJ/mol). During the reaction, two moles of H–Cl bonds are formed (bond energy = 432 kJ/mol), releasing 2 × 432 kJ; or 864 kJ. Because the bonds in the products are stronger than those in the reactants, the reaction releases more energy than it consumes:

ΔH=ƩDbonds broken−ƩDbonds formedΔH=[DH−H+DCl−Cl]−2DH−Cl=[436+243]−2(432)=−185kJ

This excess energy is released as heat, so the reaction is exothermic.

EXAMPLE 1.4.6.1

Using Bond Energies to Calculate Approximate Enthalpy Changes

Methanol, CH3OH, may be an excellent alternative fuel. The high-temperature reaction of steam and carbon produces a mixture of the gases carbon monoxide, CO, and hydrogen, H2, from which methanol can be produced. Using the bond energies in Table 1.4.6.2, calculate the approximate enthalpy change, ΔH, for the reaction here:

CO(g)+2H2(g)⟶CH3OH(g)Solution

First, we need to write the Lewis structures of the reactants and the products:A set of Lewis diagrams show a chemical reaction. The first structure shows a carbon atom with a lone pair of electrons triple bonded to an oxygen with a lone pair of electrons. To the right of this structure is a plus sign, then the number 2 followed by a hydrogen atom single bonded to a hydrogen atom. To the right of this structure is a right-facing arrow followed by a hydrogen atom single bonded to a carbon atom that is single bonded to two hydrogen atoms and an oxygen atom with two lone pairs of electrons. The oxygen atom is also single bonded to a hydrogen atom.

From this, we see that ΔH for this reaction involves the energy required to break a C–O triple bond and two H–H single bonds, as well as the energy produced by the formation of three C–H single bonds, a C–O single bond, and an O–H single bond. We can express this as follows:

ΔH=ƩDbonds broken−ƩDbonds formedΔH=[DC≡O+2(DH−H)]−[3(DC−H)+DC−O+DO−H]ΔH=ƩDbonds broken−ƩDbonds formedΔH=[DC≡O+2(DH−H)]−[3(DC−H)+DC−O+DO−H]Using the bond energy values in Table 1.4.6.1, we obtain:

 

ΔH=[1080+2(436)]−[3(415)+350+464]=−107kJ 

Check Your Learning

Ethyl alcohol, CH3CH2OH, was one of the first organic chemicals deliberately synthesized by humans. It has many uses in industry, and it is the alcohol contained in alcoholic beverages. It can be obtained by the fermentation of sugar or synthesized by the hydration of ethylene in the following reaction:A set of Lewis structures show a chemical reaction. The first structure shows two carbon atoms that are double bonded together and are each single bonded to two hydrogen atoms. This structure is followed by a plus sign, then an oxygen atom with two lone pairs of electrons single bonded to two hydrogen atoms. A right-facing arrow leads to a carbon atom single bonded to three hydrogen atoms and a second carbon atom. The second carbon atom is single bonded to two hydrogen atoms and an oxygen atom with two lone pairs of electrons. The oxygen atom is single bonded to a hydrogen atom as well.

Using the bond energies in Table 1.4.6.2, calculate an approximate enthalpy change, ΔH, for this reaction.

ANSWER:

–35 kJ

Gradient

Sometimes solutes are found at uneven concentrations within a solution. The concentration of a solute in a solution is the amount of solute per volume, usually measured in moles per litre or molarity (M) or as the percent mass of solute per volume (%). If the concentration of a solute is non-uniform within a solution, this results in a concentration gradient within that solution. A gradient refers to a difference in the distribution of matter within a system.

Beaker containing a solution with a gradient of solute.
Figure 1.4.6.6 Beaker containing a solution with a gradient of solute.

Concentration gradients are a form of chemical potential energy. This is because solutes tend to move from high concentration to low concentration. This phenomenon is known as diffusion and you will discuss it further in lab and lecture. Solutions move toward even mixing or equilibrium. Therefore, a concentration gradient within a solute represents a chemical potential where the molecules have the potential to move to an area of lower concentration within the solution. A potential refers to a difference in the distribution of energy within a system.

Ions form a chemical concentration gradient within a solution but they also bear charges. If the charge distribution within a solution is non-uniform, this results in an electrical gradient. Electrical gradients result in one part of the solution being more negatively charged and another part of the solution being more positively-charged. The resulting non-uniform distribution of charge is an electrical potential because negative charges tend to move toward areas of positive charge and vice versa. Ion gradients represent both a chemical concentration gradient AND an electrical gradient. The sum of these two gradients is collectively referred to as an electrochemical gradient and the difference in chemical and electrical energy within that system is called an electrochemical potential.

Electrochemical gradients and the resulting electrochemical potentials are biologically important. Cells actively maintain electrochemical gradients across their membranes to ensure that electrochemical signals can be sent from the brain, along nerves, to muscle and gland cells. Electrochemical gradients across cell membranes ensure that cells can generate energy. You will have a fuller understanding of these processes by the completion of this course.

In a solution, equilibrium is the even mixing of solutes within a solvent. Solutes tend to move from areas of high concentration to areas of low concentration, leading to even dispersal of the solute in solution.

Chemical reactions may also be in equilibrium. What does this mean?

Chemical reactions can occur in two directions: forward and reverse. For example, the formation of carbonic acid is endergonic in the forward reaction and exergonic in the reverse reaction. Chemical equilibrium occurs when the forward reaction and the reverse reaction occur at the same rate. Special conditions are required to achieve chemical equilibrium for each reaction. You will learn more about chemical equilibrium in the future chapters.

Inside of cells, most chemical reactions do not occur in conditions that favour chemical equilibrium. The special environment maintained within the cell ensures that some chemical reactions are irreversible, meaning that the reaction performed cannot convert the products back into the reactants. Most metabolic reactions are irreversible reactions. However, cells may change the intracellular environment to increase the rate of a reversible reaction in a particular direction. This is also true at the tissue, organ, and organ system level. For example, you will learn about how red blood cells, blood, and the respiratory system cooperate to ensure that carbon dioxide is effectively removed from the blood in a timely manner during our discussion of the respiratory system.

Cells change the conditions within the cell to affect whether certain chemical reactions will occur and how fast they will occur. This phenomenon is called regulation, and you will examine this topic in nearly every lecture of this course.

Section Summary

  • Energy is the capacity to do work. It can be neither created nor destroyed and every energy transformation dissipates some energy as heat energy.
  • A gradient is an uneven distribution of matter within a system.
  • A potential is an uneven distribution of energy within a system.
  • Cells use both chemical and electrical gradients, or electrochemical gradients, to store energy and prepare for the many energy transformations that sustain life.
  • Chemical reactions rearrange matter:
    • Endergonic chemical reactions require energy input to make/break bonds.
    • Exergonic chemical reactions release energy as bonds are made/broken.
  • Chemical reactions achieve equilibrium when the rate of the forward and reverse reactions is approximately the same:
    • The intracellular environment may favour one direction, but the intracellular environment can change.
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