Chemical Context of Life and Molecular Diversity Flashcards (copy)

Overview: The Importance of Chemistry to Life

Global Environmental Context: As Earth warms due to climate change, Arctic sea ice and glaciers are melting, fundamentally altering life support systems.
Ecological Shifts: Warmer Arctic waters and reduced ice packs lead to phytoplankton blooms (microscopic aquatic photosynthetic organisms) visible from space. While these thrive, ice-dependent organisms like black guillemots are suffering and threatened.
Water's Dominance: Three-quarters of Earth's surface is covered by water. It is the only common substance to exist in nature in three physical states: solid (ice), liquid, and gas (water vapor).
Rare Physical Property: Solid water (ice) floats on liquid water, a life-supporting property emerging from the chemical structure of the water molecule.
Biological Foundations: Biology is interdisciplinary, applying concepts of chemistry and physics to natural systems and organisms.

CONCEPT 2.1: Matter consists of chemical elements in pure form and in combinations called compounds

Definition of Matter: Anything that takes up space and has mass; it exists in forms such as rocks, metals, oils, gases, and organisms.
Elements: Substances that cannot be broken down into other substances by chemical reactions. There are 92 naturally occurring elements recognized today (e.g., gold, copper, carbon, oxygen).
Symbols: Each element has a one- or two-letter symbol, sometimes derived from Latin (e.g., Sodium is Na from natrium).
Compounds: Substances consisting of two or more different elements combined in a fixed ratio (e.g., $NaCl$ at a 1:1 ratio).
Emergent Properties: Compounds have chemical and physical characteristics different from those of their constituent elements. Sodium is a metal and chlorine is a poisonous gas, but together they form edible table salt.
The Elements of Life:
- Essential Elements: About $20-25\%$ of the 92 natural elements are needed for a healthy life and reproduction.
- Humans require 25 elements; plants require 17.
- Big Four Elements: Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N) make up approximately $96\%$ of living matter.
- Secondary Elements: Calcium (Ca), Phosphorus (P), Potassium (K), and Sulfur (S) account for most of the remaining $4\%$ .
- Trace Elements: Required in minute quantities. Iron (Fe) is needed by all life. Iodine (I) is essential for vertebrate thyroid hormones; a deficiency causes goiter. Daily intake for humans is $0.15\,mg$ .
Evolution of Tolerance: Some organisms adapt to toxic elements. Sunflowers can take up lead and zinc in concentrations lethal to most species, and were used to detoxify soil after Hurricane Katrina.

CONCEPT 2.2: An element's properties depend on the structure of its atoms

Atom: The smallest unit of matter that retains the properties of an element.
Subatomic Particles:
- Protons: Positively charged ( $+1$ unit), located in the atomic nucleus. Mass is approximately $1.7 \times 10^{-24}\,g$ or $1\,\text{dalton}$ .
- Neutrons: Electrically neutral, located in the nucleus. Mass is approximately $1\,\text{dalton}$ .
- Electrons: Negatively charged ( $-1$ unit), form a rapidly moving "cloud" around the nucleus. Mass is negligible (approx. $1/2000$ of a proton).
Atomic Measurements:
- Atomic Number: The number of protons unique to an element, written as a subscript to the left (e.g., $_{2}He$ ).
- Mass Number: Total number of protons and neutrons, written as a superscript (e.g., $^4He$ ).
- Determining Neutrons: Mass number minus the atomic number.
- Atomic Mass: The total mass of an atom (often close to the mass number).
Isotopes: Different atomic forms of the same element with the same number of protons but different numbers of neutrons.
Adding on..they are: Different structural forms of the same chemical element. They contain the same number of protons but differ in the number of neutrons in their nucleus. Because they have the same number of protons and electrons, they behave identically in chemical reactions, but their different weights give them unique nuclear and physical properties.
- Carbon Isotopes: $^{12}C$ ( $99\%$ ), $^{13}C$ ( $1\%$ ), and the rare radioactive $^{14}C$ .
- Radioactive Isotopes: Decay spontaneously, giving off particles and energy. Decay can change the element (e.g., $^{14}C$ decays into $^{14}N$ ).
- Applications: Fossil dating, metabolic tracers (e.g., kidney diagnostic tools, PET scans to find cancerous tissue with high metabolic activity).
- Hazards: Radiation can damage cellular molecules; radioactive fallout is a serious environmental threat.
Electron Energy Levels:
- Energy: Capacity to cause change or do work.
- Potential Energy: Energy matter possesses due to location or structure. Electrons have potential energy because of their distance from the nucleus.
- Electron Shells: Represent different energy levels. The first shell is closest to the nucleus (lowest energy). Successive shells have higher energy.
- Transitions: Electrons move between shells by absorbing or losing a discrete amount of energy equal to the difference between energy levels.
Electron Distribution:
- Valence Electrons: Electrons in the outermost shell (valence shell).
- Chemical Behavior: Determined by the distribution of electrons. Atoms with full valence shells are inert (e.g., Helium, Neon, Argon).
- Reactivity: Arises from the presence of one or more unpaired electrons in the valence shell.

CONCEPT 2.3: The formation and function of molecules depend on chemical bonding between atoms

Chemical Bonds: Attractions that keep atoms close together to complete their valence shells.
Covalent Bonds: The sharing of a pair of valence electrons by two atoms.
When two nonmetal atoms share pairs of valence electrons.
- Molecule: Two or more atoms held by covalent bonds.
- Single Bond: One pair of shared electrons ( $H-H$ ).
- Double Bond: Two pairs of shared electrons ( $O=O$ ).
- Valence: The bonding capacity of an atom, usually equal to the number of unpaired electrons (H=1, O=2, N=3, C=4).
Electronegativity: The attraction of a particular atom for the electrons of a covalent bond.
- Nonpolar Covalent Bond: Electrons shared equally (e.g., $H_2$ , $O_2$ ).
- Polar Covalent Bond: Electrons shared unequally (e.g., $H_2O$ where Oxygen is more electronegative, creating partial charges, $\delta-$ and $\delta+$ ).
Ionic Bonds: The complete transfer of electrons from one atom to a much more electronegative partner.
Nonmetal and metal bond-Occurs when one atom completely transfers one or more valence electrons to another, creating a positive ion (cation) and a negative ion (anion).
- Ions: Oppositely charged atoms or molecules. Cation (+); Anion (-).
- Bond: The attraction between oppositely charged ions ( $Na^+$ and $Cl^-$ ).
- Ionic Compounds: Also called salts, often forming crystalline lattices in nature. They are not discrete molecules but ratios (e.g., $NaCl$ is 1:1, $MgCl_2$ is 1:2).
Weak Chemical Interaction: Crucial for emergent properties of life and molecular recognition.
- Hydrogen Bonds: Noncovalent attraction between a hydrogen atom (covalently bonded to an electronegative atom) and another electronegative atom (usually O or N).
- Van der Waals Interactions: Brief, weak attractions between transiently charged regions of nonpolar molecules.
- Example: Geckos walking up walls via millions of tiny hairs; inspired Geckskin adhesive that can hold $700\,\text{pounds}$ .
Molecular Shape and Function: A molecule's shape (e.g., water is V-shaped at a $104.5^\circ$ angle; methane is a tetrahedron) determines how it interacts with other molecules.
- Molecular Mimicry: Opiates (morphine/heroin) relieve pain because they share a similar shape to natural endorphins and can bind to the same brain receptors.

CONCEPT 2.4: Chemical reactions make and break chemical bonds

Definition: The making and breaking of chemical bonds leading to changes in the composition of matter.
Components: Reactants (starting materials) and Products (resulting materials).
Conservation: Matter is not created or destroyed, only rearranged.
Photosynthesis: $6CO_2 + 6H_2O \rightarrow C_6H_{12}O_6 + 6O_2$ . Sunlight powers the conversion of carbon dioxide and water into glucose and oxygen.
Chemical Equilibrium: Reached when the forward and reverse reaction rates are equal. It is a dynamic equilibrium where concentrations stabilize at a specific ratio.

CONCEPT 2.5: Hydrogen bonding gives water properties that help make life possible on Earth

Polar Molecule: Water's V-shape and oxygen's electronegativity cause uneven charge distribution.
Emergent Properties of Water:
1. Cohesion: Hydrogen bonds hold water molecules together. Facilitates water transport in plants against gravity. Adhesion (water to cell walls) also helps.
2. Surface Tension: Measure of how difficult it is to stretch/break the surface. High in water, allowing some animals (e.g., raft spiders) to walk on it.
3. Moderation of Temperature:
  - Kinetic Energy vs. Thermal Energy: Thermal energy is total kinetic energy (depends on volume); Temperature is average kinetic energy.
  - Calorie (cal): Heat to raise $1\,g$ of water by $1^\circ C$ . $1\,cal = 4.184\,J$ .
  - High Specific Heat: Water resists temperature changes. Much heat goes into breaking H-bonds before speed increases. Stabilizes ocean and coastal temperatures.
  - Evaporative Cooling: The "hottest" molecules leave as gas, cooling the remaining liquid surface (e.g., sweating, plant leaf cooling).
  - Heat of Vaporization: $1\,g$ of water at $25^\circ C$ needs $580\,cal$ to evaporate.
4. Expansion Upon Freezing: Water is most dense at $4^\circ C$ . Below this, it expands into a crystalline lattice. Ice is $10\%$ less dense than liquid water. Provides an insulating habitat for marine life.
5. Versatility as a Solvent:
  - Definitions: Solution (homogeneous mixture), Solvent (dissolving agent), Solute (dissolved substance).
  - Hydration Shell: The sphere of water molecules around each dissolved ion.
  - Hydrophilic vs. Hydrophobic: Hydrophilic substances have an affinity for water (even if they don't dissolve, like cotton cellulose). Hydrophobic substances (nonpolar/nonionic) repel water (e.g., vegetable oil).
Solute Concentration:
- Molecular Mass: Sum of masses of all atoms in a molecule (Sucrose $C_{12}H_{22}O_{11} \approx 342\,\text{daltons}$ ).
- Mole (mol): $6.02 \times 10^{23}$ objects (Avogadro's number).
- Molarity (M): Number of moles of solute per liter of solution.
Acids and Bases:
- Dissociation: Water can dissociate into a Hydrogen ion ( $H^+$ /hydronium $H_3O^+$ ) and a Hydroxide ion ( $OH^-$ ).
- Acid: Increases $H^+$ concentration ( $HCl \rightarrow H^+ + Cl^-$ ).
- Base: Reduces $H^+$ concentration (by accepting $H^+$ or donating $OH^-$ ).
- pH Scale: $pH = -\log[H^+]$ . In neutral water ( $25^\circ C$ ), $[H^+] = 10^{-7}$ , so $pH = 7$ .
- Scale Range: 0 (acidic) to 14 (basic). Each unit is a tenfold difference.
Buffers: Substances that minimize pH changes by accepting/donating $H^+$ (e.g., Carbonic acid $H_2CO_3 \rightleftharpoons HCO_{3}^- + H^+$ in human blood).
Ocean Acidification: $CO_2$ absorption by oceans forms carbonic acid, lowering pH and reducing carbonate ions ( $CO_{3}^{2-}$ ). This inhibits calcification (formation of $CaCO_3$ ) in corals and shell-building animals.

CONCEPT 3.1: Carbon atoms can form diverse molecules by bonding to four other atoms

Organic Compounds: Compounds containing carbon.
Carbon Versatility: Carbon has 4 valence electrons, allowing it to form four covalent bonds in a tetrahedral shape (angles of $109.5^\circ$ ).
Carbon Skeletons: Vary in length, branching, double bond position, and presence of rings.
Hydrocarbons: Organic molecules consisting only of carbon and hydrogen (e.g., petroleum, fat tails). Highly hydrophobic and energy-rich.
Isomers: Same molecular formula, different structures.
1. Structural Isomers: Differ in covalent arrangement of atoms.
2. Cis-trans Isomers: Differ in spatial arrangement around a double bond. (Cis = same side; Trans = opposite sides).
3. Enantiomers: Mirror images due to an asymmetric carbon. Important in pharmaceuticals (e.g., ibuprofen, meth).
Chemical Groups:
- Hydroxyl ( $-OH$ ): Alcohols; polar.
- Carbonyl ( $C=O$ ): Ketones (internal) or Aldehydes (end of skeleton).
- Carboxyl ( $-COOH$ ): Acts as an acid.
- Amino ( $-NH_2$ ): Acts as a base.
- Sulfhydryl ( $-SH$ ): Thiols; can form cross-links (disulfide bridges) in proteins.
- Phosphate ( $-OPO_{3}^{2-}$ ): Contributes negative charge; involved in energy transfer.
- Methyl ( $-CH_3$ ): Affects gene expression and hormone shape.
ATP (Adenosine Triphosphate): Contains three phosphate groups; reacts with water to release energy, becoming ADP.

CONCEPT 3.2: Macromolecules are polymers, built from monomers

Definitions: Polymer (long chain of similar building blocks); Monomer (smaller building blocks).
Dehydration Reaction: Synthesizing a polymer by removing a water molecule to form a covalent bond.
Hydrolysis: Breaking down a polymer by adding a water molecule.
Diversity: Immense variety from only $40-50$ common monomers, analogous to 26 letters forming infinite words.

CONCEPT 3.3: Carbohydrates serve as fuel and building material

Monosaccharides: Simple sugars ( $CH_2O$ ratio). Glucose ( $C_6H_{12}O_6$ ) is the most common. Form rings in water.
Disaccharides: Two monosaccharides joined by a glycosidic linkage (e.g., Sucrose = glucose + fructose).
Polysaccharides:
- Storage: Starch (plants; polymers of $\alpha$ -glucose) and Glycogen (animals; highly branched polymer of $\alpha$ -glucose stored in liver/muscles).
- Structural: Cellulose (plant cell walls; polymers of $\beta$ -glucose forming straight microfibrils) and Chitin (arthropod exoskeletons and fungal cell walls).
Digestion: Humans cannot digest cellulose (insoluble fiber) because we lack enzymes to break $\beta$ linkages.

CONCEPT 3.4: Lipids are a diverse group of hydrophobic molecules

Characteristics: Not true polymers; hydrophobic due to hydrocarbon regions.
Fats (Triacylglycerols): One glycerol + 3 fatty acids joined by ester linkages.
- Saturated Fats: No double bonds; solid at room temperature (animal fats).
- Unsaturated Fats: One or more cis double bonds creating kinks; liquid at room temperature (plant/fish oils).
- Trans Fats: Unsaturated fats with trans double bonds; linked to heart disease.
Phospholipids: Glycerol + 2 fatty acids + phosphate group. Have a hydrophilic head and hydrophobic tails. Form bilayers in cell membranes.
Steroids: Carbon skeleton with four fused rings (e.g., Cholesterol, signaling hormones).

CONCEPT 3.5: Proteins include a diversity of structures and functions

Functions: Enzymes (catalysts), defense, storage, transport, cellular communication, movement, structural support.
Amino Acids: 20 types, each with an $\alpha$ carbon, amino group, carboxyl group, and a variable R group (side chain).
Polypeptides: Polymers of amino acids linked by peptide bonds.
Four Levels of Protein Structure:
1. Primary: Unique linear sequence of amino acids.
2. Secondary: Coils ( $\alpha$ helix) and folds ( $\beta$ pleated sheet) from hydrogen bonds in the polypeptide backbone.
3. Tertiary: 3D shape from R-group interactions (hydrophobic interactions, H-bonds, ionic bonds, and disulfide bridges).
4. Quaternary: Association of two or more polypeptide subunits (e.g., Hemoglobin with 4 subunits, Collagen with 3).
Sickle-cell Disease: Result of a single amino acid substitution (Glutamic acid to Valine) in hemoglobin.
Denaturation: Loss of native shape and function due to pH, salt, or temperature changes.
Protein Folding: Controlled by amino acid sequence. Assisted by other proteins. Diseases like Alzheimer's are linked to misfolded proteins.
Determination: X-ray crystallography, NMR spectroscopy, and cryo-electron microscopy are used to find 3D shapes.

CONCEPT 3.6: Nucleic acids store, transmit, and help express hereditary information

Gene Expression: DNA $\rightarrow$ RNA $\rightarrow$ Protein.
Nucleic Acids: Polymers called polynucleotides.
Nucleotide Components: Five-carbon sugar (Pentose) + Nitrogenous base + Phosphate group.
Nitrogenous Bases:
- Pyrimidines: Single ring (Cytosine [C], Thymine [T] in DNA, Uracil [U] in RNA).
- Purines: Double ring (Adenine [A], Guanine [G]).
Sugars: Deoxyribose (DNA) and Ribose (RNA).
DNA Structure: Double helix, antiparallel strands, sugar-phosphate backbones on the outside. A pairs with T; G pairs with C.
RNA Structure: Usually single-stranded; variable shapes. A pairs with U.

CONCEPT 3.7: Genomics and proteomics have transformed biological inquiry

Genomics: Analyzing large sets of genes or whole genomes.
Proteomics: Analyzing large sets of proteins.
Bioinformatics: Computational tools used to handle large biological data sets.
Cost Efficiency: Sequencing 1 million bases cost >\$5000 in 2001; cost dropped to <\$0.02 by 2017.
Evolution: Molecular genealogy confirms kinship. Humans and gorillas differ by 1 amino acid in hemoglobin; humans and frogs differ by 67. Human genome is $95-98\%$ identical to chimpanzees.
Applications: Medical science (personalized medicine), Paleontology (Neanderthal genome), Conservation (tracking illegal ivory), and understanding plant-microbe interactions.

Questions & Discussion

Q: Is a trace element an essential element?
A: Yes, they are required for an organism to live a healthy life and reproduce, even if only in minute quantities.
Q: What are the effects of iron deficiency?
A: Iron is needed for hemoglobin; deficiency could result in reduced oxygen transport in the blood, leading to fatigue and other health issues.
Q: What determines the atomic number of magnesium?
A: The number of protons (12). Magnesium has 12 electrons, 3 electron shells, and 2 valence electrons.
Q: Why are neon and argon unreactive?
A: They have full valence shells (8 electrons in the outer shell).
Q: How many water molecules are needed to hydrolyze a 10-monomer polymer?
A: 9 water molecules (one for each bond between monomers).
Q: Why does the structure $H-C=C-H$ fail to make sense chemically?
A: Each carbon is only forming 3 bonds, but carbon's valence is 4.
Q: What holds magnesium chloride ( $MgCl_2$ ) together?
A: Ionic bonds resulting from the attraction between $Mg^{2+}$ cations and $Cl^-$ anions.
Q: How would properties of water change if O and H had the same electronegativity?
A: Water would be nonpolar, meaning it would not form hydrogen bonds, losing its cohesive properties and high specific heat, making life as we know it impossible.