BIO 111 Lecture 3: Energy & Chemical Reactions

Energy Fundamentals

Defining Energy

• Energy: The capacity to do work or supply heat.

• Potential Energy: Sto red energy; energy a substance possesses due to its location or structure.

  • Examples: Water behind a dam, chemical bonds, an electron's position relative to the nucleus.

• Kinetic Energy: Energy of motion.

  • Examples: Moving water, a running animal, vibra ting molecules.

Potential Energy and Stability

• Matter tends to move towards the lowest possible state of potential energy.

• The potential energy of an electron (e-) depends on its position relative to the nucleus:

  • Moving electrons to outer shells requires energy, so electrons in outer shells have higher potential energy.

  • Changes in an electron's potential energy occur only in fixed amounts (quantized steps).

Chemical Energy

• Definition (Freeman et al., 2014): The potential energy "stored in bonds" between atoms.

• More Appropriate Definition (Campbell et al., 2011): The potential energy available for release in chemical reactions.

• Elaboration on Chemical Energ

  • Electrons possess energy that varies with their distance from the nucleus:

    • Closer to the atomic nucleus $\rightarrow$ lower potential energy.

    • Further from the atomic nucleus $\rightarrow$ higher potential energy.

  • Atoms achieve stability by filling their valence shells, often by releasing, gaining, or sharing electrons (e.g., forming a covalent bond).

  • Energy is associated with electrons within chemical bonds.

    • This is the chemically available energy found in substances like food.

    • When food molecules are broken down in chemical reactions, bonds are broken and new ones formed, releasing energy that powers life processes.

  • In nonpolar covalent bonds, electrons are shared evenly. In polar covalent bonds, more electronegative atoms pull electrons closer to their nucleus.

    • Electrons will have lower potential energy in bonds where they are held closer to a nucleus, such as in polar covalent bonds, due to stronger attraction.

Energy and Bond Formation/Breakage

• Bond Formation: Energy is released during bond formation because the new, more stable arrangement of atoms stores less energy than the separate atoms did. The energy difference is released (e.g., as heat).

• Bond Breakage: Energy is consumed (or absorbed) when chemical bonds are broken, as energy must be supplied to overcome the attractive forces holding atoms together.

• "Energy stored in bonds": This is shorthand for the potential energy that can be released when new, more stable bonds are formed after original bonds are broken.

Chemical Reactions

• Chemical Reaction: A process in which one or more substances are changed into other substances.

  • Involves the making and/or breaking of chemical bonds.

  • Substances are combined, broken down, or molecules are rearranged.

• Metabolism: The totality of all chemical reactions occurring in a living cell or organism.

Writing a Chemical Reaction

• Format: Reactants on the left, products on the right, separated by a reaction arrow ($\rightarrow$).

  • Example: $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(l)$

• Components:

  • Chemical Formula: Represents the specific molecule (e.g., $H_2$).

  • Subscript: Indicates the number of atoms of an element within a molecule (e.g., $O_2$ has two Oxygen atoms).

  • Coefficients: Numbers preceding chemical formulas, indicating the number of molecules involved in the reaction (e.g., $2H_2$ means two molecules of hydrogen).

  • State of Matter: (g) for gas, (l) for liquid, (s) for solid, (aq) for aqueous solution.

  • Reaction Arrow ($\rightarrow$): Indicates the direction of the reaction.

  • Reversible Reaction ($\rightleftharpoons$): Indicates that the reaction can proceed in both forward and reverse directions.

Chemical Equilibrium

• Definition: Occurs when the forward and reverse reaction rates are equal.

  • Quantities of reactants and products remain constant, though not necessarily equal.

  • It is a dynamic but stable state.

• Disturbance: Chemical equilibrium can be disturbed by changing the concentration of reactants or products.

• Collision Theory: For most reactions, molecules must collide in a specific orientation that brings the involved electrons near each other.

  • Concentration's Role: Higher concentration leads to a higher number of collisions, which in turn leads to a higher reaction rate.

    • Example: Adding more $CO<em>2$ to the reaction $H</em>2O + CO<em>2 \rightleftharpoons H</em>2CO<em>3$ would shift the equilibrium to the right, producing more $H</em>2CO<em>3$. Removing $H</em>2CO_3$ would also shift the equilibrium to the right.

Thermal Energy

• Definition: A form of kinetic energy associated with the random movement of atoms or molecules.

  • Low Temperature: Molecules move slowly, resulting in low average kinetic energy.

  • High Temperature: Molecules move quickly, resulting in high average kinetic energy.

Energy Transformations and Thermodynamics

• Energy Transformation: Energy can change from one form to another.

  • Example: Sunlight can excite electrons to higher energy levels. As these electrons return to lower energy levels, the potential energy difference is released as heat or light.

• Thermodynamics: The study of energy transformations that occur in collections of matter.

• System: The specific matter undergoing study.

  • Isolated System: Cannot exchange either matter or energy with its surroundings.

  • Closed System: Can exchange energy but not matter with its surroundings.

  • Open System: Can exchange both matter and energy with its surroundings.

  • Organisms are open systems. They take in energy (food, sunlight) and matter, and release energy (heat) and matter (waste products).

Laws of Thermodynamics

First Law of Thermodynamics

• Principle: The energy of the universe is constant; energy cannot be created or destroyed, only transferred or transformed from one form to another.

  • Implication: Plants are better described as "transformers" of energy (converting light energy to chemical energy) rather than "producers" (as they don't create energy).

• Enthalpy ($H$): The internal (potential) energy of a system.

  • Change in Enthalpy ($\Delta H_{reaction}$): Calculated as the sum of product enthalpies minus the sum of reactant enthalpies.

    • $\Delta H<em>{reaction} = \Sigma H</em>{products} - \Sigma H_{reactants}$ (using capital sigma for summation)~

  • Exothermic Reaction:

    • Occurs when re actants have higher enthalpy than products.

    • $\Delta H$ is negative (\Delta H < 0).

    • Energy is released, typically as heat.

    • Examples: Combustion of methane, a hot pack (releasing heat).

  • Endothermic Reaction:

    • Occurs when reactants have lower enthalpy than products.

    • $\Delta H$ is positive (\Delta H > 0).

    • Energy is required or absorbed for the reaction to happen, typically from the surroundings as heat.

    • Example: A cold pack (absorbing heat).

Second Law of Thermodynamics

• Principle: Every energy transfer or transformation increases the entropy of the universe.

• Entropy ($S$): A measure of molecular disorder or randomness.

  • More disorder or a more randomly arranged system corresponds to greater entropy.

  • The second law suggests that the change in entropy for the universe ($\Delta S<em>{universe}$) tends to increase, meaning it's positive (\Delta S{universe} > 0).

  • Change in Entropy ($\Delta S_{reaction}$): Calculated as the sum of product entropies minus the sum of reactant entropies.

    • $\Delta S<em>{reaction} = \Sigma S</em>{products} - \Sigma S_{reactants}$ (using capital sigma for summation)

• Example: A brown bear converting chemical energy from fish into kinetic energy for running also increases disorder by releasing heat and small waste molecules (CO₂, H₂O) into the surroundings.

Spontaneous & Non-Spontaneous Reactions

• Spontaneous Reactions: Reactions that proceed on their own without any continuous external influences (e.g., added energy).

  • Important Note: "Spontaneous" refers to being "energetically favorable," not "instantaneous." Spontaneous reactions can be very quick or very slow.

• Reactions tend to be spontaneous if:

  1. The products have lower potential energy than the reactants (i.e., energy is released, \Delta H < 0).

  2. The product molecules are less ordered than the reactant molecules (i.e., entropy increases, \Delta S > 0).

Exothermic vs. Endothermic (Revisited in context of Spontaneity)

• If products have lower potential energy (enthalpy):

  • $\Delta H$ is negative.

  • The reaction is exothermic; the difference in potential energy is given off as heat.

  • Example: Combustion of methane ($CH<em>4(g) + 2O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(g)$). Electrons are held more tightly in the products ($CO<em>2$ and $H</em>2O$) than in the reactants ($CH<em>4$ and $O</em>2$), resulting in lower potential energy for products and a negative $\Delta H$.

• If products have higher potential energy (enthalpy):

  • $\Delta H$ is positive.

  • The reaction is endothermic; heat energy is absorbed.

Entropy and Spontaneity

• If product molecules are less ordered (higher entropy):

  • $\Delta S$ is positive.

  • Example: If a reaction results in an increase in the number of gaseous molecules or broken bonds, that typically means higher entropy.

Gibbs Free Energy Change ($\Delta G$)

• Free Energy ($G$): The portion of a system's energy that can be harnessed to do work.

• Gibbs Free Energy Change ($\Delta G$): A measure of the change in potential energy (enthalpy, $\Delta H$) and entropy ($\Delta S$) that occurs for a given chemical reaction.

  • Purpose: Used to predict whether a reaction will be spontaneous at constant temperature and pressure.

• Equation: $\Delta G = \Delta H - T\Delta S$

  • $\Delta G$: Gibbs free energy change.

  • $\Delta H$: Change in enthalpy (potential energy of products minus potential energy of reactants).

  • $T$: Absolute temperature in Kelvin (always positive).

  • $\Delta S$: Change in entropy (entropy of products minus entropy of reactants).

Predicting Spontaneity with Gibbs Free Energy

• Spontaneous Reaction (Exergonic):

  • $\Delta G$ is negative (\Delta G < 0).

  • Net release of free energy.

  • Products are more stable than reactants (lower potential energy) and/or products have higher entropy (are more disordered).

  • Example 1: (\Delta H < 0, \Delta S > 0 at any $T$): Decreasing potential energy and increasing entropy strongly favor spontaneity, always resulting in \Delta G < 0.

  • Example 3: (\Delta H > 0, but $\Delta S \gg 0$): Potential energy increases, but entropy increases significantly. The reaction can be spontaneous at higher temperatures because the $-T\Delta S$ term becomes more negative, making \Delta G < 0.

• Non-Spontaneous Reaction (Endergonic):

  • $\Delta G$ is positive (\Delta G > 0).

  • Absorbs free energy from the surroundings.

  • Products have higher potential energy and/or products have lower entropy (are more ordered).

  • Example 2: ($\Delta H \gg 0$, \Delta S > 0): A large increase in potential energy ($\Delta H$ is largely positive) makes $\Delta G$ positive, even if entropy increases. This reaction is non-spontaneous.

    • Example: Electrolysis of water ($2H2O \rightarrow 2H2 + O_2$): Requires continuous energy input to proceed.

Factors Influencing Spontaneity

• Effect of Temperature (T): An increase in temperature ($T$) inflates the value of the $-T\Delta S$ term. If $\Delta S$ is positive, increasing $T$ makes $-T\Delta S$ more negative, thus decreasing $\Delta G$ and increasing spontaneity. If $\Delta S$ is negative, increasing $T$ makes $-T\Delta S$ more positive, thus increasing $\Delta G$ and decreasing spontaneity.

Exergonic and Endergonic Reactions (Summary)

• Exergonic Reaction:

  • $\Delta G$ is negative (\Delta G < 0).

  • Net release of free energy.

  • Products are more stable than reactants (lower potential energy) and/or products have higher entropy (are more disordered).

  • Spontaneous.

• Endergonic Reaction:

  • $\Delta G$ is positive (\Delta G > 0).

  • Absorbs free energy from the surroundings.

  • Products have higher potential energy and/or products have lower entropy (are more ordered).

  • Non-spontaneous.

Free Energy of Activation

• Definition: The initial investment of energy required to overcome an activation barrier (to reach the transition state) for a reaction to begin.

  • Even spontaneous reactions may require this energy.

• Source: Usually provided as heat absorbed from the surroundings.

• Once the activation barrier is overcome, spontaneous reactions will proceed with a net release of energy.