Enzymes, Water, Acids & Bases: Chemical Energetics and Biological Importance

Enzymes, Water, Acids & Bases

Learning Outcomes

• Understand the notation and basic energetics of chemical bond formation

• Describe the factors that influence reaction rates

• Explain the properties, actions, and importance of enzymes

• Discuss the physiologically important properties of water

• Define and describe the properties of acids and bases with respect to hydrogen ions

• Explain what the pH scale measures

• Explain how a pH buffer works

Chemical Bond Formation and Energetics

Types of Chemical Energy

Chemical reactions in biological systems involve the making and breaking of chemical bonds, which is associated with energy changes.

• Chemical energy: Potential energy stored in chemical bonds.

• Electrical energy: Energy from the flow of ions (charged particles).

• Mechanical energy: Energy directly transferred from one object to another, often as movement.

Endergonic and Exergonic Reactions

Chemical reactions can be classified based on their energy requirements or release:

• Endergonic reactions: Products have higher energy than reactants; these reactions require an input of energy to proceed.

• Exergonic reactions: Reactants have higher energy than products; these reactions release energy.

Endergonic Reaction Example: Photosynthesis

• Photosynthesis is the original source of energy for most life on Earth.

• Overall reaction: $6CO_2 + 6H_2O \rightarrow C_6H_{12}O_6 + 6O_2$

• This process is highly endergonic, requiring energy input from sunlight.

Exergonic Reaction Example: Cellular Respiration

• Cellular respiration breaks down glucose, releasing energy used to produce ATP.

• Overall reaction: $C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O$

• This process is highly exergonic, releasing energy for cellular work.

Chemical Reactions in Cells

Key Terms

• Reactants: Substances that enter into a reaction.

• Intermediates: Substances formed in the middle of a reaction.

• Products: End results of a reaction.

Chemical reactions may be reversible or irreversible. For example:

• Reversible reaction: $AB \rightleftharpoons A + B$

• Exchange reaction: $HCl + NaOH \rightarrow H_2O + NaCl$

Factors Influencing Reaction Rates

Several factors affect how quickly chemical reactions occur in biological systems:

• Physical contact: Reactants must collide to react.

• Overcoming repulsive forces: Reactants must overcome electron repulsion.

• Concentration: Higher concentrations of reactants increase reaction rates.

• Temperature: Higher temperatures generally increase reaction rates, but extreme heat can denature proteins (including enzymes).

• Catalysts: Substances that speed up reactions without being consumed.

Enzymes: Properties, Actions, and Importance

Enzymes are biological catalysts that play a crucial role in facilitating chemical reactions in the body.

• Lower activation energy: Enzymes reduce the energy required to start a reaction.

• Specificity: Each enzyme has a unique active site that binds specific substrates.

• Reusable: Enzymes are not consumed in the reaction and can be used repeatedly.

Activation Energy

• Activation energy (EA): The minimum energy required to initiate a chemical reaction.

• Enzymes lower EA, making reactions proceed faster at body temperature.

Factors Affecting Enzyme Activity

• Temperature: Increases reaction rate up to an optimal point; high temperatures denature enzymes.

• Substrate concentration: Higher substrate increases reaction rate until enzymes are saturated.

• Product concentration: High product levels can slow reactions (negative feedback).

• Enzyme concentration: More enzyme increases reaction rate up to a limit.

• pH: Each enzyme has an optimal pH; deviations can reduce activity or denature the enzyme.

Example: Pepsin (a digestive enzyme in the stomach) has an optimal pH of ~2.

Enzyme Deficiencies

• Tay-Sachs Disease: Deficiency of hexosaminidase; leads to accumulation of gangliosides in neurons, causing early death.

• Severe Combined Immunodeficiency Syndrome (SCIDs): May be due to adenosine deaminase deficiency; results in a nearly absent immune system.

• Phenylketonuria (PKU): Deficiency of phenylalanine hydroxylase; leads to buildup of phenylalanine, causing seizures and intellectual disability if untreated.

Physiologically Important Properties of Water

Water is essential for life and has several unique properties that support physiological processes:

• Heat absorption: Water absorbs heat without significant temperature change.

• Evaporative cooling: Water carries heat away as it evaporates.

• Cushioning and protection: Water protects body structures from trauma.

• Lubrication: Water reduces friction between adjacent surfaces.

Water as a Solvent

• Hydrophilic substances: Water dissolves ionic and polar covalent solutes ("like dissolves like").

• Hydrophobic substances: Nonpolar molecules do not dissolve well in water.

• Hydration spheres: Water molecules surround ions, keeping them dissociated and mobile.

Acids, Bases, and the pH Scale

Acids

• Electrolytes: Acids ionize and dissociate in water.

• Proton donors: Acids release hydrogen ions (H+) into solution.

• Acidity: Determined by the concentration of H+ in solution.

• Example: $HCl \rightarrow H^+ + Cl^-$

Bases

• Electrolytes: Bases ionize and dissociate in water.

• Proton acceptors: Bases accept H+ ions or release hydroxide ions (OH-).

• Alkalinity: Determined by the concentration of OH- in solution.

• Example: $NaOH \rightarrow Na^+ + OH^-$

• $OH^- + H^+ \rightarrow H_2O$

pH Scale

• Definition: pH measures the concentration of H+ ions in a solution.

• Formula: $pH = -\log_{10}[H^+]$

• Lower pH = higher [H+] = more acidic.

• Higher pH = lower [H+] = more alkaline/basic.

• Pure water is neutral (pH ~7).

Buffers and pH Regulation

Buffers are crucial for maintaining stable pH in biological systems, such as blood.

• Definition: Buffers are weak acids and bases that prevent rapid shifts in pH.

• They release H+ when pH rises and bind H+ when pH falls.

• Example: Carbonic acid-bicarbonate buffer in blood:

  • When pH falls: $H_2CO_3 \rightarrow HCO_3^- + H^+$

  • When pH rises: $HCO_3^- + H^+ \rightarrow H_2CO_3$

pH Indicators

• pH indicators are substances that change color depending on the pH of the solution.

• Common indicators include litmus paper and natural indicators like red cabbage extract.

Summary Table: Properties of Acids and Bases

Property	Acids	Bases
Ionization in Water	Release H+	Release OH- or accept H+
pH Range	< 7	> 7
Example	HCl	NaOH
Effect on Litmus	Turns red	Turns blue

Additional info: Some explanations and examples have been expanded for clarity and completeness, including the summary table and more detailed descriptions of enzyme deficiencies and buffer systems.