Human Anatomy & Physiology - The Chemistry of Life

Atoms and Elements

Matter

• Matter is defined as anything that has mass and occupies space. It exists in three states: solid, liquid, or gas.
• Chemistry is the study of matter and its interactions.

Atoms and Atomic Structure

• An atom is the smallest unit of matter that retains its original properties. It's composed of subatomic particles.
• Subatomic particles:
  • Protons (p+): Located in the central core of the atom (atomic nucleus) and have a positive charge.
  • Neutrons (n0): Located in the atomic nucleus, slightly larger than protons, and have no charge.
  • Electrons (e-): Located outside the atomic nucleus and have a negative charge.
• Atoms are electrically neutral, meaning the number of protons and electrons are equal, canceling each other's charge. The number of neutrons does not necessarily have to equal the number of protons.
• Electron shells are regions surrounding the atomic nucleus where electrons exist. Each shell holds a specific number of electrons:
  • 1st shell (closest to the nucleus): Holds 2 electrons.
  • 2nd shell: Holds 8 electrons.
  • 3rd shell: Holds 18 electrons, but is considered "satisfied" with 8.
  • Some atoms may have more than 3 shells.

Elements in the Periodic Table and the Human Body

• Atomic number is the number of protons in the atomic nucleus, defining each element.
  • An element is a substance that cannot be broken down into simpler substances by chemical means.
  • Each element is made of atoms with the same number of protons.
• The periodic table of elements lists elements by increasing atomic numbers.
  • Elements are organized into groups based on certain properties.
  • Each element is represented by a chemical symbol.
• The human body is primarily composed of four major elements:
  • Hydrogen
  • Oxygen
  • Carbon
  • Nitrogen
• The human body also contains 7 mineral elements and 13 trace elements.

Isotopes and Radioactivity

• Mass number is the sum of all protons and neutrons in the atomic nucleus.
• An isotope is an atom with the same atomic number (number of protons) but a different mass number (number of neutrons).
• Radioisotopes are unstable isotopes that release high energy or radiation through radioactive decay, allowing the isotope to assume a more stable form.

\text{Examples of Hydrogen Isotopes:}
^{1}{1}H \text{ (Hydrogen)}, \quad ^{2}{1}H \text{ (Deuterium)}, \quad ^{3}_{1}H \text{ (Tritium)}

Nuclear Medicine

Common applications of radioisotopes include:

• Cancer radiation therapy: Radiation damages the structure of cancer cells and interferes with their function.
• Radiotracers: Injected into a patient and detected by a camera. The image is analyzed by a computer to show the size, shape, and activity of organs and cells.
• Treatment of thyroid disorders: High doses of iodine-131 treat overactive or cancerous thyroid tissue. The radioisotope accumulates and damages the cells.

Matter Combined: Mixtures and Chemical Bonds

Matter Combined

• Matter can be combined physically to form mixtures, where atoms of two or more elements are physically intermixed without changing the chemical nature of the atoms.
• Three basic types of mixtures:
  • Suspensions
  • Colloids
  • Solutions

Mixtures

• Suspension: A mixture containing two or more components with large, unevenly distributed particles that will settle out when left undisturbed.
• Colloids: Mixtures with two or more components that have small, evenly distributed particles that will not settle out.
• Solutions: Mixtures with two or more components that have extremely small, evenly distributed particles that will not settle out.
  • Solute: The substance being dissolved.
  • Solvent: The substance that dissolves the solute.

Chemical Bonds

• Matter can be combined chemically when atoms are combined by chemical bonds.
• A chemical bond is an energy relationship or attractive force between atoms, not a physical structure.
  • Molecule: Formed by chemical bonding between two or more atoms of the same element.
  • Compound: Formed when two or more atoms from different elements combine by chemical bonding.
• Chemical bonds are formed when valence electrons (outermost shell) of atoms interact.
• Valence electrons determine how an atom interacts with other atoms and whether it will form bonds with a specific atom.
• Octet rule: An atom is most stable when it has 8 electrons in its valence shell (e.g., CO_2).
• Duet rule: For atoms with 5 or fewer electrons, the atom is most stable when the valence electron shell holds 2 electrons.

Ions and Ionic Bonds

• Ionic bond: Electrons are transferred from a metal atom to a nonmetal atom, resulting in the formation of ions (cations and anions).
  • Cation: A positively charged ion; a metal loses one or more electrons.
  • Anion: A negatively charged ion; a nonmetal gains one or more electrons.
• Attraction between opposite charges bonds ions to one another, forming a salt.

Covalent Bonds

• Covalent bonds: Strong bonds where two or more nonmetals share electrons.
• Two atoms can share one (single bond), two (double bond), or three (triple bond) electron pairs.
• Protons attract electrons; this is known as electronegativity.
  • An element's electronegativity increases from the bottom left to the upper right of the periodic table; fluorine (F) is the most electronegative element.
  • Electronegative elements strongly attract electrons and pull them away from less electronegative elements.

Nonpolar Covalent Bonds

• Nonpolar covalent bonds: Two nonmetals in a molecule with similar or identical electronegativities pull with equal force and share electrons equally.
  • Nonpolar molecules occur when:
    • Atoms sharing electrons are the same element.
    • The arrangement of atoms makes one atom unable to pull more strongly than another atom (e.g., CO_2).
    • The bond is between carbon and hydrogen.

Polar Covalent Bonds

• Polar covalent bonds: Nonmetals with different electronegativities share electrons unequally, forming polar molecules.
  • The atom with higher electronegativity becomes partially negative (\delta^−), pulling shared electrons closer.
  • The atom with lower electronegativity becomes partially positive (\delta^+), as shared electrons are pulled toward the other atom.
  • Dipoles: Polar molecules with partially positive and partially negative ends.

Hydrogen Bonds

• Hydrogen bonds: Weak attractions between the partially positive end of one dipole and the partially negative end of another dipole.
• Responsible for the key property of water – surface tension.
• Polar water molecules are more strongly attracted to one another than to nonpolar air molecules at the surface.

Determining the Type of Bonds in a Molecule

Basic "rules":

• If a compound contains both a metal and a nonmetal, the bond is ionic.
• If a molecule contains two or more nonmetals, the bond is covalent; hydrogen behaves like a nonmetal:
  • If the molecule contains two identical nonmetals, it is a nonpolar covalent bond (e.g., O_2).
  • If the molecule contains only (or primarily) carbon and hydrogen, it is a nonpolar covalent bond (e.g., CH_4).
  • If the molecule contains two nonmetals of significantly different electronegativities, it is a polar covalent bond (hydrogen and carbon have low electronegativities; oxygen, nitrogen, and phosphorus have high electronegativities); compounds with a proportionally high number of C-O, C-N, P-O, H-N, C-Cl, and H-O bonds are polar (e.g., C5H{10}O5, CH3Cl, H3PO4).

Chemical Reactions

Chemical Notation

• A chemical reaction occurs when a chemical bond is formed, broken, or rearranged, or when electrons are transferred between two or more atoms (or molecules).
• Chemical notation is a series of symbols and abbreviations that demonstrates what occurs in a reaction; a chemical equation (basic form of chemical notation) has two parts:
  • Reactants (left side of the equation): Starting ingredients that will undergo a reaction.
  • Products (right side of the equation): Results of the chemical reaction.
• Reversible reactions proceed in either direction and are denoted by two arrows in opposite directions.
• Irreversible reactions proceed from left to right and are denoted by a single arrow.

CO2 + H2O \rightleftharpoons H2CO3
Reactants (carbon dioxide + water) ⇌ Product (carbonic acid)

Energy and Chemical Reactions

• Energy is the capacity to do work or put matter into motion or fuel chemical reactions; there are two general forms:
  • Potential energy: Stored energy that can be released later to do work.
  • Kinetic energy: Potential energy that has been released or set in motion to perform work; all atoms have kinetic energy (in constant motion); faster movement = greater energy.
• Three forms of energy in the human body: chemical, electrical, and mechanical, each may be potential or kinetic depending on location or process
  • Chemical energy: Energy in bonds between atoms; drives nearly all chemical processes.
  • Electrical energy: Energy generated by the movement of charged particles or ions.
  • Mechanical energy: Energy directly transferred from one object to another.
• Energy, inherent in all chemical bonds, must be invested any time a chemical reaction occurs:
  • Endergonic reactions: Require an input of energy from another source; products contain more energy than reactants (energy invested so the reaction could proceed).
  • Exergonic reactions: Release excess energy; products have less energy than reactants.

Homeostasis and Types of Chemical Reactions

• Three fundamental processes maintain homeostasis (breaking down molecules, converting energy in food to usable form, and building new molecules); these are carried out by three types of chemical reactions:
  1. Catabolic reactions (decomposition reactions)
  2. Exchange reactions
  3. Anabolic reactions (synthesis reactions)
• Catabolic reactions (decomposition reactions): A large substance is broken down into smaller substances.
  • General chemical notation: AB \rightarrow A + B
  • Usually exergonic; chemical bonds are broken.
• Exchange reactions: One or more atoms from reactants are exchanged for one another.
  • General chemical notation: AB + CD \rightarrow AD + BC
  • Example: HCL + NaOH \rightarrow H_2O + NaCl
• Oxidation-reduction reactions (redox reactions): A special kind of exchange reaction where electrons and energy are exchanged instead of atoms.
  • Oxidized: The reactant that loses electrons..
  • Reduced: The reactant that gains electrons.
  • Redox reactions are usually exergonic reactions, releasing large amounts of energy.
• Anabolic reactions (synthesis reactions): Small, simple subunits are united by chemical bonds to make large, complex substances.
  • General chemical notation: A + B \rightarrow AB
  • Reactions are endergonic, fueled by chemical energy.

Reaction Rates and Enzymes

• For a reaction to occur, atoms must collide with enough energy to overcome the repulsion of their electrons.
• Activation energy (E_a): The energy required for all chemical reactions.
• Analogy: Activation energy must be supplied so that reactants reach transition states (i.e., get to the top of the energy “hill”) to react and form products (i.e., roll down the hill).
• The following factors increase the reaction rate by reducing activation energy or increasing the likelihood of strong collisions between reactants:
  • Concentration
  • Temperature
  • Reactant properties
  • Presence or absence of a catalyst
• When reactant concentration increases, more reactant particles are present, increasing the chance of successful collisions between reactants.
• Raising the temperature of reactants increases the kinetic energy of atoms, leading to more forceful and effective collisions between reactants.
• Both particle size and phase (solid, liquid, or gas) influence reaction rates:
  • Smaller particles move faster, having more energy than larger particles.
  • Reactant particles in the gaseous phase have higher kinetic energy than either solid or liquid phases.
• Catalyst: Increases the reaction rate by lowering the activation energy; it is not consumed or altered in the reaction.
• Enzymes: Biological catalysts, most of which are proteins, with the following properties:
  • Speed up reactions by lowering activation energy.
  • Highly specific for individual substrates (the substance that can bind to the enzyme's active site).
  • Do not alter reactants or products.
  • Are not permanently altered in the reactions they catalyze.
• Induced-fit mechanism: The enzyme's interaction with substrate(s).
  • The binding of the substrate causes a small shape change that brings the substrate to a transition state.
  • Reduces the energy of activation, allowing the transition state to proceed to final products.

Enzyme Deficiencies

Examples of common enzyme deficiencies:

• Tay-Sachs Disease: Deficiency of hexosaminidase; gangliosides accumulate around neurons of the brain; death usually occurs by age 3.
• Severe Combined Immunodeficiency Syndrome (SCIDs): May be due to adenosine deaminase deficiency; nearly complete absence of the immune system; patients must live in a sterile “bubble”.
• Phenylketonuria: Deficiency of phenylalanine hydroxylase; converts phenylalanine into tyrosine; resulting seizures and mental retardation can be prevented by dietary modification.

Inorganic Compounds: Water, Acids, Bases, and Salts Bonds

Biochemistry

• Biochemistry: The chemistry of life.
• Inorganic compounds generally do not contain carbon bonded to hydrogen; water, acids, bases, and salts.
• Organic compounds do contain carbon bonded to hydrogen.

Water

• Water (H_2O) makes up 60-80% of the mass of the human body; key properties vital to existence:
  • High heat capacity: Able to absorb heat without significantly changing temperature.
  • Carries heat with it when it evaporates (changing from liquid to gas).
  • Cushions and protects body structures because of relatively high density.
  • Acts as a lubricant between two adjacent surfaces (reduces friction).
• Water serves as the body's primary solvent; a universal solvent because so many solutes dissolve in it entirely or to some degree.
  • Polar covalent molecule:
    • Oxygen pole – partially negative (\delta^−).
    • Hydrogen pole – partially positive (\delta^+).
  • Water molecules interact with certain solutes, surround them, and keep them apart.
• Only able to dissolve hydrophilic solutes (with fully or partially charged ends); "like dissolves like"; water dissolves ionic and polar covalent solutes.
• Hydrophobic solutes: Do not have full or partially charged ends; do not dissolve in water; include uncharged nonpolar covalent molecules (oils and fats).

Acids and Bases

• The study of acids and bases is really the study of hydrogen ions (H^+).
• Water molecules in a solution may dissociate (break apart) into positively charged hydrogen ions (H^+) and negatively charged hydroxide ions (OH^−).

H_2O \rightleftharpoons H^+ + OH^-

• Acid: A hydrogen ion or proton donor; the number of hydrogen ions increases in water when an acid is added.
• Base (alkali): A hydrogen ion acceptor; the number of hydrogen ions decreases in water when a base is added.
• pH scale: Ranges from 0-14.
• A simple way of representing the hydrogen ion concentration of a solution.
• pH is the negative logarithm of the hydrogen ion concentration: pH = -Log[H^+]
• pH = 7: The solution is neutral; the number of hydrogen ions and base ions are equal.
• pH less than 7: Acidic; hydrogen ions outnumber base ions.
• pH greater than 7: Basic or alkaline; base ions outnumber hydrogen ions.
• Buffer: A chemical system that resists changes in pH; prevents large swings in pH when an acid or base is added to a solution.
• Consists of a weak acid and its corresponding anion; a major buffer is the carbonic acid – bicarbonate buffer system.

H2CO3 \rightleftharpoons HCO_3^- + H^+

• Blood pH must remain within its narrow range to maintain homeostasis; most body fluids are slightly basic:
  • Blood pH: 7.35–7.45
  • Intracellular pH: 7.2

Making Sense of the pH scale

• Why does pH decrease if a solution has more hydrogen ions?
  • Smaller pH = bigger negative log.
  • Single-digit changes in negative logarithm (e.g., from 2 to 3) accompany a 10-fold change in hydrogen ion concentration (e.g., from 0.01 to 0.001).
• Example:
  • Solution A has a hydrogen ion concentration of 0.015 M and a pH of 1.82; solution B has a hydrogen ion concentration of 0.0003 M and a pH of 3.52.
  • The solution with a higher hydrogen ion concentration has a lower -log; so the more acidic a solution, the lower its pH, and vice versa.

Salts and Electrolytes

• Salt: Any metal cation and nonmetal anion held together by ionic bonds.
• Can dissolve in water to form cations and anions called electrolytes, capable of conducting electrical current.

Organic Compounds: Carbohydrates, Lipids, Proteins, and Nucleotides

Monomers and Polymers

Each type of organic compound in the body (carbohydrate, lipid, protein, or nucleic acid) consists of polymers built from monomer subunits:

• Monomers: Single subunits; combined to build larger structures (polymers) by dehydration synthesis (links monomers together; makes a molecule of water).
• Hydrolysis: A catabolic reaction; uses water to break polymers into smaller subunits.

Carbohydrates

• Carbohydrates: Composed of carbon, hydrogen, and oxygen; function primarily as fuel; some perform limited structural roles.
  • Monosaccharides: Monomers from which all carbohydrates are made.
    • 3 to 7 carbons each.
    • Glucose, fructose, galactose, ribose, and deoxyribose are the most abundant.
  • Disaccharides: Formed by the union of two monosaccharides through dehydration synthesis.
  • Polysaccharides: Many monosaccharides joined to one another by dehydration synthesis reactions.
    • Glycogen: A storage polymer of glucose; found mostly in skeletal muscle and liver cells.
    • Polysaccharides covalently bound to proteins or lipids form glycoproteins and glycolipids, which have various functions in the body.

Lipids

• Lipids: A group of nonpolar hydrophobic molecules composed primarily of carbon and hydrogen; includes fats and oils.
• Fatty acids: Lipid monomers; 4 to 20 carbon atoms; may have none, one, or more double bonds between carbons in the hydrocarbon chain.
  • Saturated fatty acids: Solid at room temperature; no double bonds between carbon atoms; carbons are "saturated" with the maximum number of hydrogen atoms.
  • Monounsaturated fatty acids: Generally liquid at room temperature; one double bond between two carbons in the hydrocarbon chain.
  • Polyunsaturated fatty acids: Liquid at room temperature; two or more double bonds between carbons in the hydrocarbon chain.

The Good, the Bad, and the Ugly of Fatty Acids

• Good: Omega-3 Fats
  • Flaxseed oil and fish oil; cannot be made by humans; obtained in diet.
  • Polyunsaturated; positive effects on cardiovascular health.
• Bad: Saturated Fats
  • Animal fats; also in palm and coconut oils.
  • Overconsumption is associated with increased cardiac disease risk.
• Ugly: Trans Fats
  • Produced by adding H atoms to unsaturated plant oils (“partially hydrogenated oils”).
  • No safe consumption level; significantly increases the risk of heart disease.
• Triglyceride: Three fatty acids linked by dehydration synthesis to a modified 3-carbon carbohydrate, glycerol; a storage polymer for fatty acids (neutral fat).
• Phospholipids: A glycerol backbone, two fatty acid “tails,” and one phosphate “head” in place of the third fatty acid.
  • A molecule with a polar group (phosphate head) and a nonpolar group (fatty acid tail) is amphiphilic.
  • Makes phospholipids vital to the structure of cell membranes.
• Steroids: Nonpolar; share a four-ring hydrocarbon structure called a steroid nucleus.
  • Cholesterol: A steroid that forms the basis for all other steroids.

Proteins

• Two basic types of proteins by structure:
  • Fibrous proteins: Long, rope-like strands; mostly nonpolar amino acids; add strength and durability to structures.
  • Globular proteins: Spherical or globe-like; mostly polar amino acids; function as enzymes, hormones, and other cell messengers.
• Four levels of complex protein structure :
  • Primary structure: Amino acid sequence of the polypeptide chain.
  • Secondary structure: One or more segments of the primary structure folded in specific ways; held together by hydrogen bonds.
    • Alpha helix: Coiled spring.
    • Beta-pleated sheet: Venetian blind.
  • Tertiary structure: The three-dimensional shape of the peptide chain (twists, folds, and coils, including secondary structure); stabilized by hydrogen bonding.
  • Quaternary structure: Linking together more than one polypeptide chain in a specific arrangement; critical to the function of the protein.
• Protein denaturation: Destroying a protein’s shape by heat, pH changes, or exposure to chemicals.
  • Disrupts hydrogen bonding and ionic interactions that stabilize structure and function.

Nucleotides and Nucleic Acids

• Nucleotides: Monomers of nucleic acids; named for their abundance in the nuclei of cells; make up genetic material.
  • Nucleotide structure:
    • A nitrogenous base with a hydrocarbon ring structure.
    • A five-carbon pentose sugar (ribose or deoxyribose).
    • A phosphate group.
  • Types of nitrogenous bases:
    • Purines: A double-ringed molecule; adenine (A) and guanine (G).
    • Pyrimidines: A single-ringed molecule; cytosine (C), uracil (U), and thymine (T).
  • Adenosine triphosphate (ATP)
    • Adenine attached to ribose and three phosphate groups; the main source of chemical energy in the body.
    • Synthesized from adenosine diphosphate (ADP) and a phosphate group (Pi); energy from the oxidation of fuels (like glucose).
    • ATP synthesis is a highly endergonic reaction due to the negative charges on phosphate groups.
    • Hydrolysis of the bond is highly exergonic because ADP is more stable than ATP.
    • Not stored significantly by cells; the entire supply is exhausted in 60-90 seconds; cells must continually replenish the ATP supply.
    • The production of large quantities of ATP requires oxygen; the reason we breathe air.
  • DNA: An extremely large molecule in the nuclei of cells; composed of two long chains that twist around each other to form a double helix.
    • Contains genes – a recipe (code) for protein synthesis (the process to make every protein).
    • Structural features of DNA:
      • The pentose sugar deoxyribose (lacks the oxygen-containing group of ribose) forms the backbone of the strand; alternates with a phosphate group.
      • Bases: adenine, guanine, cytosine, and thymine.
      • Double helix strands are held together by hydrogen bonding between the bases of each strand.
      • Each base faces inside the double helix as strands run in opposite directions.
      • DNA exhibits complementary base pairing; purine A always pairs with pyrimidine T; purine G always pairs with pyrimidine C.
      • A = T (where = denotes 2 hydrogen bonds) and C ≡ G (where ≡ denotes 3 hydrogen bonds).
  • RNA: A single strand of nucleotides; moves between the nucleus and the cytosol; critical to making proteins.
    • RNA contains the pentose sugar ribose.
    • RNA contains uracil instead of thymine; it still pairs with adenine (A = U).
    • Transcription: RNA copies the recipe for a specific protein (gene in DNA).
    • Translation: RNA exits the nucleus to a protein synthesis location; directs the making of the protein from the recipe.