GOB Chemistry Study Guide: Chapters 5–9 Key Concepts and Applications

Chapter 5: Chemical Reactions

5.4 Oxidation and Reduction

Oxidation-reduction (redox) reactions are fundamental to both inorganic and organic chemistry. These reactions involve the transfer of electrons, resulting in changes to the oxidation states of atoms.

• Oxidation: Loss of electrons or increase in oxidation state.

• Reduction: Gain of electrons or decrease in oxidation state.

• Inorganic Redox Example: $2Na + Cl_2 \rightarrow 2NaCl$ (Na is oxidized, Cl is reduced)

• Organic Redox Example: Oxidation of alcohols to aldehydes/ketones.

• Predicting Products: Identify which species is oxidized/reduced and write the products accordingly.

Additional info: In organic chemistry, oxidation often involves the addition of oxygen or removal of hydrogen, while reduction is the opposite.

5.5 Organic Reactions: Condensation and Hydrolysis

Organic reactions can be classified by the changes they produce in molecules. Condensation and hydrolysis are two important types.

• Condensation Reaction: Two molecules combine to form a larger molecule, releasing a small molecule (often water). Example: Formation of esters from acids and alcohols.

• Hydrolysis Reaction: A molecule is split into two smaller molecules by the addition of water. Example: Hydrolysis of esters to acids and alcohols.

• Classification: Reactions may be identified as oxidation, reduction, condensation, or hydrolysis based on reactants and products.

5.6 Organic Addition Reactions to Alkenes

Addition reactions involve the addition of small molecules to double bonds in alkenes, converting them to single bonds.

• Hydrogenation: Addition of H2 to an alkene to form an alkane. $CH_2=CH_2 + H_2 \rightarrow CH_3-CH_3$

• Hydration: Addition of H2O to an alkene to form an alcohol. $CH_2=CH_2 + H_2O \rightarrow CH_3CH_2OH$

Chapter 6: Carbohydrates

6.1 Classes of Carbohydrates

Carbohydrates are classified based on the number of sugar units and their solubility.

• Monosaccharides: Single sugar units (e.g., glucose).

• Disaccharides: Two sugar units (e.g., sucrose).

• Oligosaccharides: 3–10 sugar units.

• Polysaccharides: Many sugar units (e.g., starch, cellulose).

• Soluble Fiber: Dissolves in water; found in oats, fruits.

• Insoluble Fiber: Does not dissolve; found in whole grains.

6.2 Functional Groups in Monosaccharides

Monosaccharides contain various functional groups that determine their chemical properties.

• Alcohols: Primary, secondary, tertiary alcohols based on the carbon's connectivity.

• Aldehyde Group: Present in aldoses (e.g., glucose).

• Ketone Group: Present in ketoses (e.g., fructose).

6.3 Stereochemistry in Monosaccharides

Stereochemistry describes the spatial arrangement of atoms in monosaccharides, leading to isomerism.

• D- and L-Isomers: Based on the orientation of the hydroxyl group on the chiral carbon furthest from the carbonyl.

• Enantiomers: Non-superimposable mirror images.

• Epimers: Differ at one chiral center.

• Diastereomers: Stereoisomers that are not mirror images.

• Common Monosaccharides: Glucose, galactose, fructose, ribose.

6.4 Reactions of Monosaccharides

Monosaccharides undergo reactions at the anomeric carbon, forming cyclic structures and redox products.

• Cyclic α and β Anomers: Formed by intramolecular reaction between carbonyl and hydroxyl groups.

• Oxidation of Aldoses: Produces acids.

• Reduction of Aldoses: Produces sugar alcohols.

6.5 Disaccharides

Disaccharides are formed by glycosidic bonds between two monosaccharides.

• Condensation and Hydrolysis: Formation and breakdown of disaccharides.

• Glycosidic Bond: Covalent bond joining two sugars.

• Common Disaccharides: Sucrose, lactose, maltose.

• Sweetness Scale: Sucrose is standard; fructose is sweeter, lactose is less sweet.

6.6 Polysaccharides

Polysaccharides are large carbohydrates composed of repeating sugar units linked by glycosidic bonds.

• Identification: By type of glycosidic bond and sugar subunit.

• Examples: Starch (α-1,4 bonds), cellulose (β-1,4 bonds), glycogen (branched α-1,6 bonds).

6.7 Carbohydrates and Blood

Carbohydrates play a role in blood group compatibility and anticoagulation.

• ABO Compatibility: Determined by specific oligosaccharides on red blood cells.

• Heparin: A polysaccharide with anticoagulant properties.

Chapter 7: States of Matter and Attractive Forces

7.2 Liquids and Solids: Predicting Properties Through Attractive Forces

Attractive forces between molecules determine the physical properties of liquids and solids.

• Types of Attractive Forces:

  • London Dispersion Forces

  • Dipole-Dipole Interactions

  • Hydrogen Bonding

  • Ionic Bonds

  • Covalent Bonds

• Hydrogen Bonds: Occur between molecules with N-H, O-H, or F-H bonds.

• States of Matter: Solid, liquid, gas; changes depend on attractive forces.

7.3 Attractive Forces and Solubility

Solubility is influenced by the types of attractive forces present in molecules.

• Golden Rule of Solubility: "Like dissolves like"—polar substances dissolve in polar solvents, nonpolar in nonpolar.

• Solubility in Water: Polar and ionic compounds are generally soluble; nonpolar compounds are not.

Chapter 8: Solution Chemistry

8.1 Solutions Are Mixtures

Solutions are homogeneous mixtures of two or more substances.

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving.

• Colloids: Intermediate particle size; do not settle.

• Suspensions: Large particles; settle out over time.

8.2 Formation of Solutions

Temperature and pressure affect the formation and solubility of solutions.

• Saturated Solution: Contains maximum solute at given conditions.

• Dilute Solution: Contains small amount of solute.

• Temperature Effect: Solubility of solids increases with temperature; gases decrease.

8.3 Chemical Equations for Solution Formation

Electrolytes dissociate in water to form ions, while nonelectrolytes do not.

• Strong Electrolytes: Completely dissociate.

• Weak Electrolytes: Partially dissociate.

• Nonelectrolytes: Do not dissociate.

8.4 Concentration

Concentration expresses the amount of solute in a given amount of solution.

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Percent Units: $\% = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100$

• Parts per Million/Billion: $ppm = \frac{\text{mg solute}}{\text{L solution}}$

8.5 Dilution

Dilution involves adding solvent to decrease the concentration of a solution.

• Dilution Equation: $M_1V_1 = M_2V_2$

8.6 Osmosis and Diffusion

Osmosis and diffusion are processes that move substances across membranes.

• Osmosis: Movement of water across a semipermeable membrane from low to high solute concentration.

• Diffusion: Movement of particles from high to low concentration.

8.7 Transport Across Cell Membranes

Transport across cell membranes occurs via different mechanisms.

• Passive Transport: No energy required (diffusion, osmosis).

• Facilitated Transport: Uses proteins to help substances cross.

• Active Transport: Requires energy to move substances against concentration gradient.

Chapter 9: Acids, Bases, and Buffers in the Body

9.1 Acids and Bases—Definitions

Acids and bases are defined by their ability to donate or accept protons (H+).

• Arrhenius Definition: Acids produce H+ in water; bases produce OH-.

• Brønsted-Lowry Definition: Acids donate H+; bases accept H+.

• Physical Properties: Acids taste sour, bases taste bitter and feel slippery.

9.2 Strong Acids and Bases

Strong acids and bases dissociate completely in water.

• Strong Acids: HCl, HNO3, H2SO4

• Strong Bases: NaOH, KOH

• Neutralization Reaction: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

9.3 Chemical Equilibrium

Chemical equilibrium occurs when the rates of forward and reverse reactions are equal.

• Equilibrium Expression: $K = \frac{[\text{products}]}{[\text{reactants}]}$

• Le Châtelier's Principle: System shifts to counteract changes in concentration, pressure, or temperature.

9.4 Weak Acids and Bases

Weak acids and bases only partially dissociate in water, establishing equilibrium.

• Conjugate Acid-Base Pairs: Acid and base differ by one proton.

• Equilibrium Equation Example: $HA + H_2O \rightleftharpoons H_3O^+ + A^-$

9.5 pH and the pH Scale

pH measures the acidity or basicity of a solution.

• pH Equation: $pH = -\log[H^+]$

• Relationship: $[H^+][OH^-] = 1 \times 10^{-14}$ at 25°C

• Acidic: pH < 7; Neutral: pH = 7; Basic: pH > 7

9.6 pKa

pKa is the negative logarithm of the acid dissociation constant, indicating acid strength.

• Lower pKa: Stronger acid.

• Prediction: Compare pH and pKa to determine predominant species.

9.7 The Relationship Between pH, pKa, Drug Solubility, and Diffusion

The ionization state of pharmaceuticals affects their solubility and ability to diffuse across membranes.

• Henderson-Hasselbalch Equation: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$

• Charged vs. Uncharged: Uncharged forms diffuse more easily; charged forms are more soluble in water.

9.8 Buffers and Blood—The Bicarbonate Buffer System

Buffers resist changes in pH. The bicarbonate buffer system maintains blood pH.

• Bicarbonate Buffer Equation: $CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons HCO_3^- + H^+$

• Physiological Importance: Regulates blood pH during changes in ventilation.