OCT 15- Chemical reactions & Energy

Overview

• Weekly schedule information provided regarding impending exams, emphasizing the critical importance of timely preparation due to the rapid approach of test dates.

Amino Acids

• Definition of Amino Acids: The fundamental building blocks of proteins, each characterized by a central alpha-carbon atom bonded to an amino group ($-NH_2$), a carboxyl group ($-COOH$), a hydrogen atom, and a unique R group (side chain).

  • Key Component: The R group (side chain) is the variable component that dictates the specific chemical properties and functions of each amino acid.

  • Importance of R group: The R group primarily determines whether an amino acid is polar, nonpolar, acidic, or basic. This characteristic is crucial for protein folding and function.

  • Method for determining the nature of R group: While a chart is the definitive reference, general guidelines include:

    • The presence of highly electronegative atoms like oxygen (e.g., in hydroxyl $-OH$ or carboxyl $-COOH$ groups) or nitrogen (e.g., in amine groups) within the R group typically indicates polarity, allowing for hydrogen bonding.

    • Hydrocarbon R groups, consisting primarily of carbon and hydrogen atoms, are generally nonpolar and hydrophobic.

    • The presence of sulfur can be tricky; for instance, cysteine contains a polar sulfhydryl group ($-SH$), while methionine contains a nonpolar thioether linkage.

Discussion of Enzymes

• Definition of Enzymes: Biological catalysts, primarily proteins, that significantly accelerate the rate of specific biochemical reactions without being consumed in the process. They achieve this by lowering the activation energy required for a reaction to proceed.

• Transitioning to a discussion on energy and its critical role in driving these enzyme-catalyzed reactions.

Types of Energy

• Potential Energy vs. Kinetic Energy:

  • Potential Energy: Stored energy that possesses the capacity to do work but is not actively doing so. This energy can be stored in various forms:

  • Gravitational potential energy: E.g., a rock at the top of a hill, water behind a dam.

  • Chemical potential energy: Stored in the bonds between atoms within molecules (e.g., glucose, ATP).

  • Elastic potential energy: Stored in a stretched spring or an inflated balloon.

  • Analogy: Individuals not engaging with their full capabilities; comparison of potential energy stored in the chemical bonds of a fuel versus its release when oxidized.

  • Example: An inflated balloon holds elastic potential energy, which can be rapidly converted into kinetic energy (motion and sound) when the valve is opened and air escapes.

  • Kinetic Energy: The energy of motion, actively doing work. It is the energy released when potential energy is converted.

  • Examples: A student walking (mechanical kinetic energy), light (electromagnetic kinetic energy), heat (thermal kinetic energy), electricity (electrical kinetic energy).

Hydroelectric Dams

• Process of Hydroelectric Energy Generation:

  • Mechanism: Water buildup on one side of a dam stores gravitational potential energy. When released, this water flows downwards.

  • The flowing water possesses kinetic energy, which then turns large turbines (mechanical energy).

  • These turbines are connected to generators, which are typically surrounded by copper wire coils. The rotation of the turbines within a magnetic field induces an electric current in the copper wires through electromagnetic induction, converting mechanical energy into electrical energy.

  • Thus, water stored in the dam (potential energy) transforms into kinetic energy as it flows, then into mechanical energy in the turbines, and finally into electrical energy.

Energy in Food

• Calories: A common unit used to measure the energy content in food. Specifically, 1 calorie (cal) is the amount of energy required to raise the temperature of 1 gram of water by 1 degree Celsius. Food labels typically use kilocalories (kcal or Cal), where $1 \, \text{Cal} = 1000 \, \text{cal}$.

  • Example: An orange (approximately 60 Calories) provides significantly less energy than a candy bar (around 200 Calories), indicating differences in the amount of stored chemical potential energy within their molecular bonds.

  • Explanation of calories: The energy stored within the chemical bonds of macronutrients (carbohydrates, fats, proteins) is released through metabolic processes.

  • Energy Conversion: The conversion of chemical energy from food into usable forms, like ATP, is not perfectly efficient. A significant portion of this energy (typically 10-15% in biological systems) is inevitably lost as heat, primarily due to the Second Law of Thermodynamics and the inherent inefficiencies of biochemical pathways.

Types of Energy in Biological Systems

• Categorization: Biological systems primarily utilize light energy, heat energy, and chemical energy.

  • Light energy: Crucial for photosynthesis, where radiant energy from the sun is captured by plants and converted into chemical energy (glucose).

  • Heat energy: Plays a vital role in biological processes, such as maintaining optimal body temperature (thermogenesis) in warm-blooded organisms and influencing reaction rates.

  • Chemical energy: Stored in the chemical bonds of molecules like glucose and adenosine triphosphate (ATP) from food and drink. This energy is key in driving metabolic processes, muscle contraction, and active transport.

Chemical Reactions in Biology

• Dehydration Synthesis and Hydrolysis:

  • Dehydration Synthesis (Condensation Reaction): A process where two smaller molecules (monomers) are joined together to form a larger molecule (polymer) with the removal of a water molecule. This reaction typically requires an input of energy and forms new covalent bonds.

  • Hydrolysis: The reverse process of dehydration synthesis, where a larger molecule is broken down into smaller constituent molecules by the addition of a water molecule. This reaction typically releases energy and breaks covalent bonds.

Energy Reactions: Endergonic vs. Exergonic

• Endergonic Reaction: A non-spontaneous chemical reaction that requires a net input of energy from its surroundings to proceed. The products of an endergonic reaction have higher free energy than the reactants, meaning "Delta G" > 0.

  • Associated with anabolic reactions (biosynthesis), which build complex molecules from simpler ones.

  • Example: Photosynthesis, where light energy is absorbed to convert carbon dioxide and water into glucose.

• Exergonic Reaction: A spontaneous chemical reaction that releases a net amount of free energy into its surroundings. The products have lower free energy than the reactants, meaning "Delta G" < 0.

  • Associated with catabolic reactions, which break down complex molecules into simpler ones.

  • Example: Cellular respiration, where glucose is broken down to produce carbon dioxide, water, and ATP.

Thermodynamics

• First Law of Thermodynamics (Law of Conservation of Energy): States that energy cannot be created or destroyed in an isolated system; it can only be transformed from one form to another. In biological systems, this means the total energy before and after any biological process remains constant.

• Second Law of Thermodynamics: States that during every energy transfer or transformation, some energy becomes unusable, and the entropy (disorder or randomness) of the universe or an isolated system increases. Biological systems must constantly input energy to maintain their highly ordered state.

  • Entropy: A measure of disorder or randomness within a system. Examples like ice melting demonstrate increasing entropy as water molecules gain more freedom of movement, leading to a less ordered liquid state.

Gibbs Free Energy

• Definition: Gibbs free energy ($G$) is a thermodynamic potential that measures the useful or process-initiating work obtainable from an isothermal, isobaric thermodynamic system. It represents the amount of energy available to do work in a system.

• Formula: $G = H - TS$

  • Components:

  • $G$: Gibbs free energy, representing the maximum reversible work that may be performed by a thermodynamic system at a constant temperature and pressure.

  • $H$: Enthalpy, which is the total heat content of a system at constant pressure. It includes the internal energy of the system plus the product of pressure and volume.

  • $T$: Absolute temperature in Kelvin ($K$). This factor scales the importance of entropy to the overall free energy.

  • $S$: Entropy, the measure of the system's disorder or randomness.

  • Interpretation of Delta G ($\Delta G$), the change in Gibbs free energy during a reaction:

  • Positive $\Delta G$: The products have more free energy than the reactants. The reaction is endergonic and non-spontaneous, requiring an input of energy to proceed.

  • Negative $\Delta G$: The products have less free energy than the reactants. The reaction is exergonic and spontaneous, releasing energy and capable of doing work.

  • $\Delta G = 0$: The system is at equilibrium.

Example Reactions

• Photosynthesis: A classic example of an endergonic (anabolic) reaction.

  • Chemical Equation: $6CO<em>2 + 6H</em>2O + \text{light energy} \rightarrow C<em>6H</em>{12}O<em>6 + 6O</em>2$

  • Low-energy reactants (carbon dioxide and water) are converted into a high-energy product (glucose) and oxygen, primarily due to the input of light energy. This process increases the order of the system and thus consumes free energy.

• Cellular Respiration: A classic example of an exergonic (catabolic) reaction.

  • Chemical Equation: $C<em>6H</em>{12}O<em>6 + 6O</em>2 \rightarrow 6CO<em>2 + 6H</em>2O + \text{ATP (energy)}$

  • High-energy reactants (glucose and oxygen) are broken down to produce lower-energy products (carbon dioxide and water), releasing a significant amount of free energy primarily in the form of ATP. This process increases the entropy of the system and releases free energy usable for cellular work.

Enzyme Activity and Speeding Reactions

• Discussion of Hydrogen Peroxide:

  • Hydrogen peroxide ($H<em>2O</em>2$) naturally decomposes into water ($H<em>2O$) and oxygen ($O</em>2$): $2H<em>2O</em>2 \rightarrow 2H<em>2O + O</em>2$.

  • When applied to a cut, the enzyme catalase (found in blood and tissues) rapidly breaks it down, demonstrating a spontaneous reaction that requires a catalyst to occur at an observable rate.

  • It's important to understand that a reaction being spontaneous (thermodynamically favorable, a negative $"Delta G"$) does not mean it is instantaneous (occurs rapidly). Many spontaneous reactions have high activation energies and thus proceed very slowly without a catalyst.

• Factors Influencing Reaction Rates: Besides the presence of catalysts (enzymes), reaction rates can be influenced by:

  • Temperature: Generally, higher temperatures increase reaction rates (up to an optimum).

  • Pressure: Particularly for gaseous reactants, higher pressure increases reaction rates.

  • Concentration of reactants: Higher concentrations lead to more frequent collisions.

  • Enzymes: Enzymes dramatically increase the rate of spontaneous reactions by lowering their activation energy. They are highly specific, often following a lock-and-key or induced-fit model, and their activity can be regulated by factors like pH, temperature, and substrate concentration.

Misconceptions Clarified

• Spontaneous vs. Instantaneous: It is crucial to distinguish between these terms. A spontaneous reaction is one that is thermodynamically favored (i.e., it can occur without a continuous input of energy, having a negative $"Delta G"$). An instantaneous reaction occurs immediately or very rapidly. Many spontaneous reactions, like the rusting of iron or the decomposition of hydrogen peroxide without catalase, can be extremely slow because they possess a high activation energy. Enzymes, therefore, primarily affect the rate of a reaction by lowering the activation energy, not its spontaneity.