Test 3 Study Guide: Chemical Reactions, Carbohydrates, and Related Concepts

Factors Affecting Reaction Rate

Overview

The rate of a chemical reaction describes how quickly reactants are converted into products. Several factors influence this rate, which are crucial for understanding both laboratory and biological reactions.

• Concentration: Higher concentration of reactants increases the likelihood of collisions, speeding up the reaction.

• Temperature: Raising temperature increases kinetic energy, leading to more frequent and energetic collisions.

• Catalysts: Catalysts lower the activation energy, allowing reactions to proceed faster without being consumed.

• Surface Area: Greater surface area (e.g., powdered solids) exposes more reactant particles, increasing reaction rate.

• Nature of Reactants: Ionic compounds generally react faster than covalent compounds due to easier bond rearrangement.

Example: Enzymes in the body act as biological catalysts, speeding up metabolic reactions.

Reaction Types

Synthesis, Decomposition, Single/Double Exchange

Chemical reactions can be classified based on how atoms and molecules are rearranged.

• Synthesis (Combination): Two or more substances combine to form a single product. $A + B \rightarrow AB$

• Decomposition: A single compound breaks down into two or more simpler substances. $AB \rightarrow A + B$

• Single Exchange (Replacement): One element replaces another in a compound. $A + BC \rightarrow AC + B$

• Double Exchange (Replacement): Two compounds exchange ions to form two new compounds. $AB + CD \rightarrow AD + CB$

Example: The reaction of sodium chloride with silver nitrate is a double exchange reaction.

Blood Types

Overview

Blood types are determined by the presence of specific antigens on the surface of red blood cells.

• ABO System: Types A, B, AB, and O are based on the presence or absence of A and B antigens.

• Rh Factor: Positive or negative, depending on the presence of the Rh antigen.

Example: Type O negative is the universal donor; type AB positive is the universal recipient.

Fischer Projections

Overview

Fischer projections are a two-dimensional representation of three-dimensional molecules, commonly used for carbohydrates and amino acids.

• Vertical lines: Represent bonds going away from the viewer.

• Horizontal lines: Represent bonds coming toward the viewer.

Example: Glucose can be drawn in a Fischer projection to show its stereochemistry.

Diastereomers vs. Enantiomers

Overview

Stereoisomers are molecules with the same connectivity but different spatial arrangements.

• Enantiomers: Non-superimposable mirror images; have identical physical properties except for optical activity.

• Diastereomers: Not mirror images; differ in physical properties and reactivity.

Example: D-glucose and L-glucose are enantiomers; D-glucose and D-mannose are diastereomers.

Drawing Pyranose from Fischer Projection

Overview

Pyranose forms are six-membered ring structures of sugars, derived from the cyclization of the Fischer projection.

• Step 1: Identify the carbonyl group and the hydroxyl group that will react.

• Step 2: Draw the ring, placing substituents according to the Fischer projection.

Example: D-glucose forms a pyranose ring by cyclizing between C1 and C5.

Structures of Disaccharides and Polysaccharides

Overview

Disaccharides and polysaccharides are carbohydrates formed by linking monosaccharides via glycosidic bonds.

• Disaccharides: Two monosaccharides joined (e.g., sucrose, lactose).

• Polysaccharides: Long chains of monosaccharides (e.g., starch, cellulose, glycogen).

Example: Sucrose is composed of glucose and fructose linked by an α-1,2-glycosidic bond.

Carbohydrate Classification

Overview

Carbohydrates are classified based on their structure and complexity.

• Monosaccharides: Simple sugars (e.g., glucose, fructose).

• Disaccharides: Two monosaccharides (e.g., sucrose).

• Oligosaccharides: 3-10 monosaccharides.

• Polysaccharides: More than 10 monosaccharides (e.g., starch).

Example: Cellulose is a polysaccharide composed of β-glucose units.

Glycosidic Bond Types

Overview

Glycosidic bonds are covalent bonds that link carbohydrate molecules.

• α-Glycosidic bond: Formed when the OH group on the anomeric carbon is in the alpha position.

• β-Glycosidic bond: Formed when the OH group is in the beta position.

Example: Lactose contains a β-1,4-glycosidic bond between galactose and glucose.

Anomers

Overview

Anomers are isomers of monosaccharides that differ in configuration at the anomeric carbon.

• α-Anomer: OH group on the anomeric carbon is trans to the CH2OH group.

• β-Anomer: OH group is cis to the CH2OH group.

Example: α-D-glucose and β-D-glucose are anomers.

Alcohol Classification

Overview

Alcohols are classified based on the number of alkyl groups attached to the carbon bearing the hydroxyl group.

• Primary (1°): One alkyl group attached.

• Secondary (2°): Two alkyl groups attached.

• Tertiary (3°): Three alkyl groups attached.

Example: Ethanol is a primary alcohol; isopropanol is a secondary alcohol.

Reactions: Hydrogenation, Hydration, Hydrolysis, Combustion, Redox

Overview

These reactions are fundamental in both organic and biological chemistry.

• Hydrogenation: Addition of hydrogen to unsaturated bonds.

• Hydration: Addition of water to a molecule.

• Hydrolysis: Splitting a molecule by adding water.

• Combustion: Reaction with oxygen producing energy, CO2, and H2O.

• Redox: Involves transfer of electrons; oxidation is loss, reduction is gain.

Example: Hydrolysis of sucrose yields glucose and fructose.

Calories in Foods

Overview

The caloric content of food is determined by the energy released during combustion.

• Carbohydrates: 4 kcal/g

• Proteins: 4 kcal/g

• Fats: 9 kcal/g

Example: A food containing 10g fat, 20g carbohydrate, and 5g protein has $10 \times 9 + 20 \times 4 + 5 \times 4 = 186$ kcal.

Reaction Energy Diagram

Overview

Reaction energy diagrams illustrate the energy changes during a chemical reaction.

• Activation Energy: The energy required to initiate the reaction.

• Exothermic: Products have lower energy than reactants; energy is released.

• Endothermic: Products have higher energy than reactants; energy is absorbed.

Example: Combustion reactions are exothermic, releasing heat.

Endo- and Exothermic/Exer- and Endergonic Reactions

Overview

These terms describe whether energy is absorbed or released in a reaction.

• Exothermic: Releases heat; $ext{ΔH} < 0$.

• Endothermic: Absorbs heat; $ext{ΔH} > 0$.

• Exergonic: Releases free energy; $ext{ΔG} < 0$.

• Endergonic: Absorbs free energy; $ext{ΔG} > 0$.

Example: Cellular respiration is exergonic; photosynthesis is endergonic.