Quiz Review

Chapter 2: Atoms & Chemical Bonds

1. Elements, Atoms, Molecules, and Compounds

• Element: A substance that cannot be broken down into simpler substances by chemical reactions.

• Atom: The smallest unit of matter that retains all properties of an element.

• Molecule: Two or more atoms held together by covalent bonds.

• Compound: Two or more different elements combined in a fixed ratio.

  • A molecule containing different elements is also considered a compound.

2. Atomic Structure & Subatomic Particles

• Subatomic Particles:

  • Protons: Positive charge ($+1$).

  • Neutrons: Neutral / no charge ($0$).

  • Electrons: Negative charge ($-1$); determine an atom's overall charge and chemical reactivity.

• Atomic Calculations:

  • $\text{Atomic Number} = \text{number of protons}$

  • $\text{Mass Number} = \text{number of protons} + \text{number of neutrons}$

3. Isotopes

• Isotopes: Atoms of the same element that have the same number of protons but different numbers of neutrons.

• Radioactive Isotopes: Unstable isotopes whose nuclei decay spontaneously, releasing particles and energy. Used as biological tracers to track metabolic pathways and biological processes.

4. Electrons & Valence Shells

• Electrons reside in energy levels or electron shells surrounding the nucleus.

• Electrons farther from the nucleus possess higher potential energy.

• Valence Shell: The outermost electron shell.

• Valence Electrons: Electrons in the valence shell available for chemical bonding.

• Atom reactivity is driven by filling or completing its valence shell.

5. Chemical Bonds & Interactions

• Covalent Bonds: Sharing of a pair of valence electrons by two atoms.

  • Single Bond: Sharing $1$ pair of electrons.

  • Double Bond: Sharing $2$ pairs of electrons.

  • Bonding capacity depends on the number of unpaired valence electrons.

• Electronegativity: An atom's attraction for shared electrons in a covalent bond.

  • Nonpolar Covalent Bond: Electrons are shared equally (or nearly equally) due to similar electronegativity.

  • Polar Covalent Bond: Electrons are shared unequally due to differing electronegativity.

  • Example in $H_2O$: Oxygen is more electronegative than hydrogen, pulling shared electrons closer to form a partial negative charge ($\delta^-$) on oxygen and partial positive charges ($\delta^+$) on hydrogen.

Electronegativity Difference & Bond Types

• $\text{Difference } 0.0 \text{ to } 0.4$: Nonpolar Covalent (shared equally, no charges)

• $\text{Difference } 0.4 \text{ to } 1.8$: Polar Covalent (shared unequally, partial charges $\delta^-$, $\delta^+$)

• $\text{Difference } > 1.8$: Ionic (electrons transferred completely, full charges $+$, $-$)

• Ionic Bonds: Complete transfer of electrons from one atom to another, resulting in oppositely charged ions that attract each other.

  • Cation: A positively charged ion formed by losing electrons (e.g., $Na^+$).

  • Anion: A negatively charged ion formed by gaining electrons (e.g., $Cl^-$).

  • Example: $NaCl$ formation where $Na$ loses $1$ electron ($Na^+$) and $Cl$ gains $1$ electron ($Cl^-$).

• Hydrogen Bonds: Noncovalent attraction between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom.

• Van der Waals Interactions: Weak, transient attractions between adjacent molecules caused by temporary local partial charges.

  • Example: Enables geckos to climb vertical surfaces.

• Shape & Function: Molecular structure dictates biological activity, essential for signal recognition, enzymes, and cell receptors.

Chapter 3: Water & the Fitness of the Environment

1. Water & Polarity

• Water ($H_2O$) is a polar molecule held together internally by polar covalent bonds between oxygen and hydrogen.

• Hydrogen bonds continuously form and break between adjacent water molecules.

2. Four Emergent Properties of Water

1. Cohesion & Adhesion:

   • Cohesion: Attraction between identical water molecules via hydrogen bonding.

   • Adhesion: Attraction between water molecules and other polar or charged substances.

   • Surface Tension: Measure of how difficult it is to stretch or break the surface of a liquid.

   • Enables water transport against gravity in plants.

2. Expansion Upon Freezing:

   • Water expands as it freezes because hydrogen bonds stabilize into a crystalline lattice structure, keeping ice molecules farther apart than in liquid water.

   • Ice is less dense than liquid water and floats, insulating aquatic environments below.

3. Versatility as a Solvent:

   • Water is an effective solvent due to its polarity.

   • Hydration Shell: A sphere of water molecules surrounding dissolved ions ($\delta^-$ oxygen attracts $Na^+$; $\delta^+$ hydrogen attracts $Cl^-$).

   • Hydrophilic: Substances with an affinity for water (polar or ionic).

   • Hydrophobic: Substances that repel water (nonpolar or nonionic).

4. Temperature Moderation:

   • Kinetic Energy: Energy of motion; average kinetic energy is measured as temperature.

   • Thermal Energy: Total kinetic energy of molecules in a body of matter.

   • High Specific Heat: Water absorbs or releases large amounts of heat with minimal temperature change.

   • Evaporative Cooling: As liquid evaporates, high-energy molecules escape, cooling the remaining surface.

Chapter 4: Carbon & Molecular Diversity

1. Importance of Carbon

• Carbon forms organic compounds by forming up to $4$ covalent bonds with elements such as hydrogen, oxygen, nitrogen, and other carbon atoms.

2. Carbon Skeletons & Hydrocarbons

• Carbon chains form skeletons that vary in length, branching, double bond position, and rings.

• Hydrocarbons: Organic molecules consisting solely of carbon and hydrogen (e.g., nonpolar regions of lipids) that store large amounts of energy.

3. Isomers

• Isomer: Compounds with identical molecular formulas but distinct structural arrangements and properties.

  • Structural Isomers: Differ in the covalent arrangement of their atoms.

  • Cis-Trans Isomers: Differ in spatial arrangements around a rigid double bond.

4. Functional Groups

• Specific chemical groups attached to carbon skeletons that participate in chemical reactions or modify molecular function.

• Example Steroids (Estradiol vs. Testosterone):

  • Both share the same four-ring carbon skeleton structure.

  • Estradiol: Features a hydroxyl group ($-OH$).

  • Testosterone: Features a carbonyl group ($C=O$) and a methyl group ($-CH_3$).

  • Differences in functional groups lead to binding with different cellular receptors.

5. ATP (Adenosine Triphosphate)

• Primary energy-transferring molecule in cells consisting of adenosine attached to three phosphate groups.

• Reaction for energy release:   $\text{ATP} \rightarrow \text{ADP} + \text{P}_i + \text{energy}$

Chapter 5: Macromolecules

1. Proteins

• Building Blocks: Amino acids linked by peptide bonds to form polypeptide chains.

• Amino Acid Structure: Central carbon (alpha-carbon), amino group ($-NH_2$), carboxyl group ($-COOH$), hydrogen atom, and a variable R group (side chain).

  • Nonpolar R groups = Hydrophobic (positioned in protein interior).

  • Polar / Charged R groups = Hydrophilic (positioned on protein exterior).

• Four Levels of Protein Structure:

  1. Primary Structure: Linear sequence of amino acids.

  2. Secondary Structure: Coils (alpha-helices) and folds (beta-pleated sheets) held by hydrogen bonds along the polypeptide backbone.

  3. Tertiary Structure: Overall 3D shape resulting from interactions among R group side chains.

  4. Quaternary Structure: Association of two or more individual polypeptide chains working together.

• Denaturation: Loss of a protein's native 3D structure and function due to changes in temperature, pH, or salt concentration.

• Structural variation example: A single amino acid change in hemoglobin causes sickle-cell anemia.

2. Nucleic Acids

• Function: Store, transmit, and express hereditary information (DNA and RNA).

• Building Blocks: Nucleotides linked in a $5' \rightarrow 3'$ direction.

• Nucleotide Components:

  • $5$-carbon sugar (deoxyribose or ribose)

  • Phosphate group

  • Nitrogenous base

• Nitrogenous Bases:

  • Purines: Adenine ($A$) and Guanine ($G$) (two rings).

  • Pyrimidines: Cytosine ($C$), Thymine ($T$), and Uracil ($U$) (single ring).

• Base Pairing Rules:

  • $A \leftrightarrow T$ (DNA) or $A \leftrightarrow U$ (RNA) via $2$ hydrogen bonds.

  • $G \leftrightarrow C$ via $3$ hydrogen bonds.

3. DNA vs. RNA

• Sugar: Deoxyribose in DNA vs. Ribose in RNA.

• Bases: Thymine ($T$) in DNA vs. Uracil ($U$) in RNA.

• Structure: DNA is double-stranded (forming a double helix); RNA is single-stranded (can fold into diverse 3D structures).