Study Notes: Introduction to Biology and Chemistry for Biology

Introduction to Biology

• Concepts

  • Biological systems are structured at many levels that interrelate and interact.

  • Cells and organisms are made of organic molecules with specific properties and water.

  • Scientific methods validate predictions through experimentation by testing a hypothesis and finding substantial evidence.

• Outline

  • I. Why study biology

  • II. What is life

  • III. How to study life

  • IV. Classification system in biology

I. WHY STUDY BIOLOGY

• The study of biology helps us understand the nature of life and the mechanisms underlying life processes.

• Applications include finding cures for diseases, improving agriculture, and protecting the environment.

• Why biology is increasingly important today:

  • The 21st century will be shaped by information from biological sciences merging with other fields.

  • Large volume of biology publications (e.g., >500 thousand papers per year).

  • DNA as the secret code of life; genome sequencing progress; genetic modification of organisms.

• Applications span medical, agricultural, and veterinary sciences.

• Why study cellular and molecular biology:

  • The connecting basis of all life is at the cell and molecular level.

  • DNA → RNA → protein and cellular mechanisms form the fundamental basis of life.

  • Many biological phenomena are better understood at biochemical and molecular levels.

  • Gives ability to genetically alter DNA, develop cures, and diagnose precisely.

• To study biology, start with properties of living organisms, model systems, and broad approaches to study life.

II. WHAT IS LIFE

• Life is carbon-based, organic in nature, and built from major molecules: carbohydrates, proteins, lipids, nucleic acids.

• Living organisms require macro- and micronutrients and water (up to ~95% water).

• The basic unit of life is the cell.

• Emergent properties of living organisms:

  1. Reproduction: Life comes from life; genetic material is DNA.

  2. Growth and development: Organisms undergo growth and development.

  3. Order and Structure: Highly ordered; structure correlates with function.

  4. Metabolism: Energy intake and turnover; synthesis and breakdown of molecules.

  5. Respiration: Oxygen uptake and CO₂ release to generate energy.

  6. Response to environmental stimuli: Sensing surroundings and maintaining homeostasis.

  7. Adaptation and evolution: Short-term and long-term adaptations; billions of years of evolution.

  8. Autonomous movement: Bacteria, protists, animals move; fungi/plants respond to nutrition/light.

III. HOW TO STUDY LIFE

• Model systems:

  • Used because it’s impractical to study every organism in detail.

  • Definition: a representative organism or cell type used for experiments.

  • Criteria: easy to grow, manipulate, and study; abundant genetic information available.

• Examples of model systems:

  • Prokaryotes: E. coli, Salmonella (Salmonella typhimurium)

  • Eukaryotes: Arabidopsis (A. thaliana), Corn (Zea mays), Rice (Oryza sativa), Yeast (Saccharomyces cerevisiae)

  • Animals: Drosophila melanogaster, Caenorhabditis elegans, Mouse (Mus musculus), Zebra fish (Danio rerio)

  • Human cell lines: HeLa cells

• Broad approaches to study life:

  • Holism: Study whole organisms for behavioral, physiological, and nutritional aspects (e.g., aging in rats).

  • Reductionism: Study at cellular, tissue, or molecular levels; use cell lines or tissue suspensions.

  • In vivo: Living conditions; physiology/ecology; holistic or reductionist with cells/tissues.

  • In vitro: Non-living conditions; chemical and biochemical experiments; strictly reductionist.

  • In situ: Determine presence of DNA, RNA, or protein at a site (e.g., FISH for locating genes on chromosomes).

  • In silico: Computer analysis of genome/proteome data; bioinformatics/computational biology.

• Scientific reasoning:

  • Hypothesis as testable prediction; experiments with proper controls.

  • Inductive approach: Specific observations lead to generalizations (e.g., Darwin’s evolution).

  • Deductive approach: General concepts predict specific outcomes (e.g., birds have feathers → a peacock has feathers).

  • Hypotheses must be testable; experiments require proper treatments and controls (positive/negative).

  • Reproducibility across researchers and conditions can lead to a theory; universally proven results may become a law.

IV. CLASSIFICATION SYSTEM

• Taxonomic Classification:

  • Standard system to group and classify living organisms; updated as consensus grows.

  • Current system overview described with a chart/concept map (Figure references: 1.2, Concept Map 1.1).

• Concept Map 1.1: Domains and kingdoms of living organisms

  • 1. Bacteria (prokaryotes): diverse unicellular organisms before nucleus development.

  • 2. Archaea (archaebacteria): prokaryotic, extreme environments, some eukaryote-like features.

  • 3. Eukarya: true nucleus; membrane-bound.

  • Eukaryotic kingdoms:

  • Protista: predominantly unicellular eukaryotes (heterotrophic like paramecium/amoeba; phototrophic like algae); includes some multicellular forms (kelp).

  • Fungi: mostly multicellular; some unicellular; heterotrophic (e.g., yeast).

  • Plantae: higher plants (flowering and non-flowering); includes Monocots and Dicots; photoautotrophs.

  • Animalia: animals; multicellular; heterotrophic.

• V. BIOLOGICAL HIERARCHY

  • Living and non-living are built from atoms and molecules; hierarchy studied at different levels.

  • Levels (from simple to complex):

  • Atoms, Molecules, Macromolecules, Parts of cells (membranes, organelles), Cells, Tissues, Organs, Organ systems, Multicellular organisms, Population, Community, Ecosystem, Biomes, Biosphere.

  • Diagram references: Figure 1.3 (Biological Hierarchy).

Chemistry for Biology

Concepts

• Fundamental ideas:

  • Protons define elements; neutrons define isotopes; electrons determine chemical/physical properties.

  • Atoms form molecules via covalent, ionic, and hydrogen bonds.

  • Polarity differences (electronegativity) affect bonds and interactions.

  • Polarity of water and hydrogen bonding underlie many properties essential to life.

  • Water is a major component of life; participates in biochemical reactions.

  • pH and hydrogen ion concentration affect biomolecules and reactions.

  • Organic compounds are diverse; structure and functional groups define properties.

• Outline:

  • I. Atoms and Molecules

  • II. Water and Aqueous Solutions

  • III. Carbon Compounds

I. ATOMS AND MOLECULES

• Basic definitions:

  • Matter: substance with mass and volume.

  • Element: substance of a single type of atom; cannot be broken down by ordinary means; defined by proton number (atomic number).

  • Compound: two or more different elements combined.

  • Molecule: two or more atoms bonded together.

  • Atom: basic unit of matter; consists of protons, neutrons, electrons.

• Subatomic particles and masses:

  • Proton: positive, mass ~1 amu (Da).

  • Neutron: neutral, mass ~1 amu (Da).

  • Electron: negative, mass ~1/1000 amu (Da).

  • Mass number / atomic weight: total mass of protons + neutrons; reference to H or O for absolute weight; average isotopic weights.

• Key biological elements:

  • Living matter ~96% made of C, O, H, N; ~3–4% P, S, Ca, K; trace elements (e.g., Fe, Mg, I).

• Isotopes and radioisotopes:

  • Isotopes: same Z (protons), different neutrons; examples: 12C, 13C, 14C; 14C is radioactive with t1/2 = $5{,}730\ ext{ years}$.

  • Stable isotopes used in experiments (e.g., 14N, 15N) for labeling and mass spectrometry.

  • Radioisotopes emit radiation; applications include dating and molecular tracing.

• Electron configuration and bonding concepts:

  • Valence electrons: outermost electrons; determine bonding capacity.

  • Octet rule: atoms tend to have 2 electrons in first shell and 8 in second/third shells to be stable.

  • Electronegativity: atom’s tendency to attract electrons; higher values indicate stronger pull.

  • Typical electronegativity values (approximate): F ~ 4.0, O ~ 3.5, C ~ 3.5, H ~ 2.1, Na ~ 1.0.

  • Valence and bonding outcomes:

  • H, O, N, C valences: H = 1, O = 2, N = 3, C = 4 (illustrative).

• Chemical bonds and interactions (overview):

  • Covalent bonding: sharing electrons; strongest type (~50–170 kcal/mol); includes glycosidic, ester, peptide, and phosphodiester bonds.

  • Covalent bonds can be polar or non-polar:

  • Non-polar covalent: equal sharing; examples: H2, O2, CO2, CH4, C2H6, C3H8.

  • Polar covalent: unequal sharing; examples: H2O, NH3; partial charges due to electronegativity differences.

  • Ionic bonds: transfer of electrons; formation of cations and anions; weaker in aqueous solutions (~3–7 kcal/mol); example: NaCl.

  • Hydrogen bonds: dipole-dipole interaction involving partially positive H and electronegative partners (O, N, F); ~3–7 kcal/mol; important in water, DNA/RNA, proteins.

  • Hydrophilic and hydrophobic interactions:

  • Hydrophilic: polar/charged, water-soluble.

  • Hydrophobic: non-polar, water-repelling; drive membrane assembly and protein folding.

  • Van der Waals forces: weak attractions from transient dipoles in close-packed molecules; contribute to interactions in lipids and cellulose.

• Summary of bonds/interactions:

  • Covalent: sharing electrons; can be non-polar or polar.

  • Ionic: transfer of electrons; attraction of opposite charges; weaker in aqueous environments.

  • Hydrogen bonds: dipole attractions involving H; moderate strength.

  • Hydrophilic/hydrophobic: polarity-driven solubility and interactions with water.

  • Van der Waals: weak, transient interactions between close atoms.

II. WATER AND AQUEOUS SOLUTIONS

• Water is essential; living organisms are up to ~95% water.

• Water properties and their relevance:
  1) Hydrogen bonding and cohesiveness: enables transport in plants (xylem) and seed imbibition.
  2) High specific heat: resistance to temperature change; contributes to stable environments.
  3) High heat of vaporization: supports evaporative cooling.
  4) Ice expansion and lower density: ice floats; preserves aquatic life in cold climates.
  5) Versatile solvent for polar/charged molecules: facilitates nutrient transport.
  6) Medium for biochemical reactions: most reactions occur in aqueous environments; water participates as reactant/product.

• Water’s role in photosynthesis: uses water to extract electrons to fix CO2 into carbohydrates.

• Aqueous solutions: two major properties are solute concentration and hydrogen ion concentration.

• Solute concentration concepts:

  • Molarity (M): one mole of solute per liter of solution: $M = \frac{\text{moles of solute}}{\text{L}}$

  • Molecular weight (MW) / formula weight (FW): sum of atomic weights; MW expressed in Daltons (Da).

  • One mole contains $N_A = 6.023 \times 10^{23}$ molecules.

  • Example: NaOH MW = 40; 1 M NaOH = 40 g per liter.

  • Fractions and unit conversions (examples shown):

  • 1 mM = 1 × 10^{-3} M, 1 µM = 1 × 10^{-6} M, 1 nM = 1 × 10^{-9} M, 1 pM = 1 × 10^{-12} M, 1 fM = 1 × 10^{-15} M.

  • Making molar solutions from solid:

  • Formula: MW × M × L (with unit conversions as needed) to obtain grams.

  • Example: Make 0.5 M NaOH in 100 mL: grams = $MW \times M \times \text{volume (L)} = 40 \times 0.5 \times 0.1 = 2 \text{ g}$

  • Making solutions from stock to diluted (\n C1 V1 = C2 V2

  • Variables: C1 = stock concentration, V1 = volume of stock; C2 = desired concentration, V2 = final volume.

  • Example: To make 100 mL of 2% NaCl from 10% NaCl:

    • 10% × V1 = 2% × 100 mL ⇒ V1 = 20 mL; remainder is water.

• Hydrogen ion concentration, acids, bases, and buffers:

  • In pure water at equilibrium: [H^+] = [OH^-].

  • pH is the negative logarithm of hydrogen ion concentration: $\mathrm{pH} = -\log_{10} [\mathrm{H^+}]$

  • pH scale ranges from 1 to 14 in practical terms; as pH changes by 1 unit, the [H^+] changes by a factor of 10.

  • Acid: dissociates to increase [H^+] (proton donor); e.g., $\mathrm{HCl} \rightarrow \mathrm{H^+} + \mathrm{Cl^-}$

  • Base: accepts protons or increases [OH^-] (proton acceptor); e.g., $\mathrm{NH<em>3} + \mathrm{H^+} \rightarrow \mathrm{NH</em>4^+}$; $\mathrm{NaOH} \rightarrow \mathrm{Na^+} + \mathrm{OH^-}$

  • Neutral pH is around 7; acidic < 7; basic > 7; examples: stomach acid pH ~2, blood ~7.4, bleach ~12.5.

  • Buffer: substance that minimizes pH changes by buffering against added acids or bases; most buffers are weak acids or weak bases.

  • Buffering range and pK:

  • pK is the pH at which the acid form and base form are in equal concentration: [acid] = [base].

  • Example: carbonic acid/bicarbonate system; at pH = pK, concentrations of H2CO3 and HCO3^- are equal.

  • Tris buffer has a pK around 8.1; buffers are chosen based on the desired buffering range (acidic, neutral, or basic).

• Carbon compounds: carbon-based chemistry is central to biology; emphasis on structure and functional groups.

III. CARBON COMPOUNDS

• Carbon-based life: main elements include C, H, O, N, P, S; similar elemental percentages across living systems.

• Organic molecules vary greatly in size and complexity:

  • Simple organic compounds: hydrocarbons (e.g., methane CH4, ethane CH3-CH3).

  • Complex macromolecules: proteins, polysaccharides, DNA, RNA can contain millions of carbons.

  • Shapes: linear (aliphatic, e.g., glycerol), branched (e.g., isoleucine), circular/aromatic (e.g., phenol, cholesterol).

  • Saturation: saturated (no C=C) vs unsaturated (one or more C=C).

  • Structural representations: molecular formula (e.g., CH4) and structural formula; sometimes drawn as line structures.

• Isomers: same molecular formula, different structures and properties.

  • Structural isomers: different arrangements (e.g., leucine vs isoleucine).

  • Geometric (cis/trans) isomers: around double bonds; cis means same side; trans means opposite sides.

  • Optical isomers (enantiomers): mirror images; L- versus D- forms; organisms often use one form (e.g., L-amino acids).

• Functional groups (key to properties and reactivity): groups covalently bonded to carbon skeleton.

  • Hydroxyl group (-OH): polarity; increases water solubility; found in alcohols, sugars, glycerol.

  • Carbonyl group (-C=O): polar; aldehydes (end) vs ketones (middle); present in simple sugars, some proteins and nucleotides; can form ring structures with -OH groups.

  • Carboxyl group (-COOH): acidic; ionizes to -COO^-; present in fatty acids and amino acids; forms conjugate bases such as acetate, formate, citrate, malate.

  • Amino group (-NH2): bases; accepts protons to form -NH3^+; present in amino acids and nucleotides; buffers by donation/acceptance of protons.

  • Sulfhydryl group (-SH): reactive; cysteine contains -SH; can form disulfide bridges (-S-S-) to stabilize proteins; active-site catalysis in some enzymes.

  • Phosphate group (-OPO3^{2-}): present in ATP, nucleotides, DNA/RNA; acidic and highly reactive; conjugate bases of phosphoric acid; Pi and PPi denote inorganic phosphate and pyrophosphate.

  • Methyl group (-CH3): non-polar; influences hydrophobicity and bioactivity; involvement in methylation and membrane permeability; found in many biomolecules.

• Functional groups and biological relevance:

  • Recognizing functional groups helps explain properties and reactivity of biomolecules.

  • Methylation can regulate DNA function and drug behavior; methyl groups influence solubility and interactions.

• Visual summaries:

  • Amphipathic molecules: phospholipids in membranes; proteins often have both hydrophobic and hydrophilic regions.

  • Phospholipids arrange into bilayers with hydrophilic heads and hydrophobic tails; membranes formed by amphipathic molecules.

  • Summary of functional groups presented in a consolidated table (molecular formula, properties, examples).

IV. ISOMERS AND FUNCTIONAL GROUPS (RECAP)

• Isomer types recap:

  • Structural isomers: same formula, different connectivity.

  • Geometric isomers: cis/trans around double bonds.

  • Optical isomers: mirror images with different biological activity.

• Functional groups recap (examples and properties):

  • Hydroxyl (-OH): polar; increases solubility; found in sugars, glycerol, alcohols.

  • Carbonyl (-C=O): polar; aldehydes/ketones; in simple sugars, some proteins and nucleotides.

  • Carboxyl (-COOH or -COO^-): acidic; deprotonates to form -COO^-; in amino acids and fatty acids.

  • Amino (-NH2): base; forms -NH3^+ at physiological pH; in amino acids and nucleotides.

  • Sulfhydryl (-SH): forms disulfide bridges; cysteine-containing active sites.

  • Phosphate (-OPO3^{2-}): important in high-energy compounds (ATP); in DNA/RNA and phospholipids.

  • Methyl (-CH3): non-polar; affects solubility and biological activity; methylation effects on DNA and drugs.

• Functional group table (condensed):

  • Hydroxyl: $-OH$; Alcohol; in sugars, glycerol, ethanol.

  • Carbonyl: $-C=O$; Aldehyde (end) vs Ketone (middle); in simple sugars, some proteins, nucleotides.

  • Carboxyl: $-COOH$ or $-COO^-$; Acid or conjugate base; in amino acids and fatty acids.

  • Amine: $-NH2$ or $-NH3^+$; Base; in amino acids, nucleotides.

  • Phosphate: $-OPO_3^{2-}$; in ATP, nucleotides, phospholipids.

  • Sulfhydryl: $-SH$; in cysteine; forms disulfide bonds.

  • Methyl: $-CH_3$; non-polar; in many biomolecules; affects bioactivity.

• Notation for some key formulas and concepts:

  • pH: $\mathrm{pH} = -\log_{10} [\mathrm{H^+}]$

  • Water ion product: $K_w = [\mathrm{H^+}][\mathrm{OH^-}] = 10^{-14}$ at standard conditions.

  • Avogadro's number: $N_A = 6.023 \times 10^{23}$ molecules per mole.

  • Molarity: $M = \frac{\text{moles of solute}}{\text{L}}$

  • Dilution (C1V1 = C2V2): $C<em>1 V</em>1 = C<em>2 V</em>2$

• Key takeaways for exam readiness:

  • Understand how bonds and interactions govern biomolecular structure and function.

  • Be able to identify functional groups in given molecules and predict properties/reactivity.

  • Recall examples of model systems and the distinctions between in vivo/in vitro/in situ/in silico approaches.

  • Apply the pH, buffering concepts, and dilution calculations to practical laboratory scenarios.

Note: Figures referenced (e.g., Figure 1.2, Figure 1.3, Concept Map 1.1) are not reproduced here but are cited as part of the source material.