Fundamentals of Anatomy & Physiology - Chapter 2 Vocabulary

Chemical Level of Organization

Learning Outcomes

• 2-1 Describe an atom and how atomic structure affects interactions between atoms.
• 2-2 Compare the ways in which atoms combine to form molecules and compounds.
• 2-3 Distinguish among the major types of chemical reactions that are important to physiology.
• 2-4 Describe the role of enzymes in metabolism.
• 2-5 Distinguish between inorganic compounds and organic compounds.
• 2-6 Explain how the chemical properties of water make life possible.
• 2-7 Explain what pH is and discuss its importance.
• 2-8 Describe the physiological roles of acids, bases, and salts and the role of buffers in body fluids.
• 2-9 Describe monomers and polymers, and the importance of functional groups in organic compounds.
• 2-10 Discuss the structures and functions of carbohydrates.
• 2-11 Discuss the structures and functions of lipids.
• 2-12 Discuss the structures and functions of proteins.
• 2-13 Discuss the structures and functions of nucleic acids.
• 2-14 Discuss the structures and functions of high-energy compounds.

Introduction to the Chemical Level

• Chemistry: The science that deals with the structure of matter.
• Matter: Anything that takes up space and has mass.

Atoms and Atomic Structure

• Atom: The smallest stable unit of matter.
• Atoms are composed of subatomic particles:
  • Protons: Positively charged.
  • Neutrons: Neutral.
  • Electrons: Negatively charged.
  • Protons and neutrons have a similar mass and are found in the nucleus.
  • Electrons are much lighter and orbit around the nucleus.
• Atomic Structure
  • Atomic number: the number of protons in an atom.
  • Electron cloud: area around the nucleus that contains electrons.
  • Electron shell: two-dimensional representation of an electron cloud.
• Element: A pure substance composed of atoms of one kind, which cannot be broken down into simpler substances by ordinary physical means.
  • 92 naturally occurring elements.
  • The atomic number (number of protons) is unique for each element.
  • The human body consists of 13 main elements and 14 trace elements (found in very small amounts).
• Principal Elements in the Human Body
  • Oxygen (O): 65% of total body weight; component of water and other compounds; essential for respiration.
  • Carbon (C): 18.6% of total body weight; found in all organic molecules.
  • Hydrogen (H): 9.7% of total body weight; component of water and most other compounds.
  • Nitrogen (N): 3.2% of total body weight; found in proteins, nucleic acids, and other organic compounds.
  • Calcium (Ca): 1.8% of total body weight; found in bones and teeth; important for membrane function, nerve impulses, muscle contraction, and blood clotting.
  • Phosphorus (P): 1.0% of total body weight; found in bones and teeth, nucleic acids, and high-energy compounds.
  • Potassium (K): 0.4% of total body weight; important for proper membrane function, nerve impulses, and muscle contraction.
  • Sodium (Na): 0.2% of total body weight; important for blood volume, membrane function, nerve impulses, and muscle contraction.
  • Chlorine (Cl): 0.2% of total body weight; important for blood volume, membrane function, and water absorption.
  • Magnesium (Mg): 0.06% of total body weight; a cofactor for many enzymes.
  • Sulfur (S): 0.04% of total body weight; found in many proteins.
  • Iron (Fe): 0.007% of total body weight; essential for oxygen transport and energy capture.
  • Iodine (I): 0.0002% of total body weight; a component of hormones of the thyroid gland.
  • Trace elements: silicon (Si), fluorine (F), copper (Cu), manganese (Mn), zinc (Zn), selenium (Se), cobalt (Co), molybdenum (Mo), cadmium (Cd), chromium (Cr), tin (Sn), aluminum (Al), boron (B), and vanadium (V); some function as cofactors; the functions of many trace elements are poorly understood.
• Isotopes are versions of an element that have different numbers of neutrons.
  • Mass number: The number of protons plus the number of neutrons of an atom; unique for each isotope.
  • Radioisotopes: Isotopes that have unstable or radioactive nuclei.
    • The breakdown of the nucleus is called radioactive decay.
    • The decay rate is expressed as half-life – the time required for half of a given amount of isotope to decay.
    • Radioactive isotopes are used in diagnostic procedures.
• Atomic weight: An average of the different atomic masses and proportions of different isotopes.
• Atomic mass: The actual mass of an atom, taking into account the masses of its protons, neutrons, and electrons.
• Mole (mol): A specific quantity that has a weight in grams equal to the atomic weight of the element.
  • 1 mol of a given element always contains the same number of atoms, called Avogadro’s number.
• Electrons and energy levels:
  • Atoms are electrically neutral because they have equal numbers of protons and electrons.
  • Electron cloud contains electron shells, or energy levels, that can hold a limited number of electrons.
    • The first electron shell can hold up to 2 electrons.
    • The next electron shells can hold up to 8 electrons.
    • The outermost shell is the valence shell, and the number of electrons in this shell determines the chemical properties of an element.
    • If the valence shell is not full, the atom is unstable and will react with other atoms.

Molecules and Compounds

• Inert elements: Have a full valence shell and are not reactive.
• Chemical reactions allow reactive atoms to become stable by gaining, losing, or sharing electrons.
• Chemical bonds: Hold participating atoms together once the reaction has ended.
• Molecule: Two or more atoms joined by shared electrons.
• Compound: Two or more atoms of different elements.
• Molecular weight of a molecule or compound is the sum of the atomic weights of its atoms.
• Types of chemical bonds:
  1. Ionic bonds
  2. Covalent bonds
  3. Hydrogen bonds
• Ionic bonds: Bonds created by the attractions between negative and positive ions.
  • An ion is an atom with an electric charge.
  • During ionic bond formation:
    • One atom is an electron donor and loses one or more electrons and becomes a cation (positively charged ion).
    • Another atom is an electron acceptor and gains those same electrons and becomes an anion (negatively charged ion).
    • The cation and anion attract each other and stay together in an ionic compound.
• Covalent bonds: Strong bonds created when atoms share electrons.
  • Sharing one pair of electrons is a single covalent bond.
  • Sharing two pairs of electrons is a double covalent bond.
  • Sharing three pairs of electrons is a triple covalent bond.
  • Nonpolar covalent bonds: Equal sharing of electrons between two atoms that have equal pull on the electrons.
  • Polar covalent bonds: Unequal sharing of electrons between two atoms.
    • Results in the formation of polar molecules in which one side is partially positive and the other partially negative.
• Hydrogen bonds: Weak electrical attractions between the partial positive charge of a hydrogen atom in a polar covalent bond and the partial negative charge of an atom of oxygen, nitrogen, or fluorine in another polar covalent bond.
  • Hydrogen bonds form between water molecules and are responsible for the high surface tension of water.
  • Hydrogen bonds also form in large molecules such as DNA.
• States of matter:
  • Solid: Has constant volume and shape.
  • Liquid: Has constant volume but no fixed shape.
  • Gas: No constant volume and no fixed shape.

Chemical Reactions

• Chemical reactions: Result in the formation of new bonds or the breaking of existing bonds.
• Reactants: The reacting substances that enter the reaction.
• Products: The resulting substances at the end of the reaction.
• Metabolism: All chemical reactions that are occurring in the cells and tissues of the body at any given time.
• Work: The movement of an object or a change in the physical structure of matter.
• Energy: The capacity to do work.
  • Kinetic energy: The energy of motion.
  • Potential energy: Stored energy an object possesses because of its position or its physical or chemical structure.
  • Energy cannot be destroyed, only converted from one form to another.
  • Each time energy is converted, some is released in the form of heat. Cells cannot capture heat or use it for work.
• Types of chemical reactions:
  • Decomposition reactions
  • Synthesis reactions
  • Exchange reactions
  • Reversible reactions
• Decomposition reactions break molecules into smaller fragments.
  • Hydrolysis reactions are decomposition reactions that involve water.
  • Decomposition reactions are referred to as catabolism; they release energy that can be harnessed to power cellular processes.
• Synthesis reactions assemble larger molecules from smaller molecules.
  • Dehydration synthesis (condensation) reactions are synthesis reactions that result in the production of a molecule of water.
  • Synthesis reactions are referred to as anabolism; they use energy to create chemical bonds.
• Exchange reactions rearrange existing components of molecules into new products.
• Reversible reactions
  • Many important biological reactions are freely reversible.
  • At equilibrium, the amounts of chemicals do not change even though the forward and reverse reactions are occurring.
    • Reversible reactions seek equilibrium, balancing opposing reaction rates.
    • When reactants are added or removed, the reaction rates change until they reach a new equilibrium.

Enzymes

• Biochemical reactions are reactions that happen in living organisms.
• Most biochemical reactions would not occur quickly enough if left to happen spontaneously.
• Activation energy is the amount of energy needed to start the reaction.
• Enzymes are catalysts that lower the activation energy of reactions and therefore speed up the reaction.
  • The enzyme is not changed or consumed in the reaction.
  • The enzyme does not affect the nature of the reaction, just its rate.
  • Reactions often happen in a series of steps, and each is controlled by a specific enzyme.
• Exergonic reactions: The amount of energy released is greater than the activation energy, so net release of energy.
• Endergonic reactions: The activation energy is more than the energy produced, so the reaction absorbs energy.

Inorganic and Organic Compounds

• Nutrients: Substances from food that are necessary for normal physiological function; carbohydrates, proteins, lipids, vitamins, minerals, etc.
• Metabolites: Substances that are involved in metabolism.
• Inorganic compounds: Do not contain carbon in bonds to hydrogen as a primary structural component; carbon dioxide, oxygen, water, inorganic acids, bases, and salts.
• Organic compounds: Contain carbon in bonds to hydrogen; carbohydrates, proteins, lipids, and nucleic acids.

Properties of Water

• Water:
  • The most important substance in the body.
  • Accounts for up to two-thirds of total body weight.
• Properties of water:
  1. Universal solvent
  2. Reactivity
  3. High heat capacity
  4. Lubrication
• Universal solvent: Many molecules and compounds are water-soluble.
  • Solution: Uniform mixture of two or more substances.
  • Solute: The dissolved substance.
  • Solvent: The liquid in which the solutes are distributed.
  • Aqueous solution: Solution in which water is the solvent.
• Reactivity:
  • In our body, chemical reactions take place in water.
  • Water also actively participates in reactions, such as dehydration synthesis and hydrolysis.
• High heat capacity:
  • Heat capacity is the quantity of heat required to raise the temperature of a unit mass of a substance by 1°C.
  • Water has an unusually high heat capacity because of hydrogen bonding, so
    • It remains liquid over a broad range of temperatures.
    • Freezing and boiling temperatures are far apart.
    • Carries a lot of heat when it evaporates.
    • Changes temperature slowly = thermal inertia.
• Lubrication: Water reduces friction between opposing surfaces.
• Properties of aqueous solutions:
  • Water is a polar molecule.
  • Many inorganic compounds held together by ionic bonds split into smaller molecules via dissociation in water.
  • Ionization is dissociation into ions.
  • Polar water molecules form hydration spheres around ions and small polar molecules and so keep them in solution.
  • Aqueous solutions which contain ions conduct electrical current.
• Electrolytes and body fluids:
  • Electrolytes are soluble inorganic substances whose ions conduct electricity in solution.
  • Electrolyte imbalance seriously disturbs vital body functions.
  • The concentration of electrolytes in body fluids is carefully regulated.
• Important Electrolytes That Dissociate in Body Fluids
  • NaCl (sodium chloride) \rightarrow Na^+ + Cl^-.
  • KCl (potassium chloride) \rightarrow K^+ + Cl^-.
  • CaPO4 (calcium phosphate) \rightarrow Ca^{2+} + PO4^{2-} .
  • NaHCO3 (sodium bicarbonate) \rightarrow Na^+ + HCO3^-.
  • MgCl_2 (magnesium chloride) \rightarrow Mg^{2+} + 2Cl^-.
  • Na2HPO4 (sodium hydrogen phosphate) \rightarrow 2Na^+ + HPO_4^{2-} .
  • Na2SO4 (sodium sulfate) \rightarrow 2Na^+ + SO_4^{2-} .
• Hydrophilic and hydrophobic compounds:
  • Hydrophilic: Interact readily with water; includes ions and polar molecules.
  • Hydrophobic: Do not interact readily with water; includes nonpolar molecules, fats, and oils.
• Colloids and suspensions:
  • Colloid: A solution containing dispersed proteins or other large molecules; example: blood plasma.
  • Suspension: Contains large particles in solution, but if undisturbed, the particles will settle out of solution; example: whole blood.

pH and Homeostasis

• pH: The negative logarithm of the hydrogen ion concentration of a solution in moles per liter.
• Neutral pH:
  • Equal numbers of H^+ and OH^- ions.
  • pH of pure water = 7.0.
• Acidic pH (between 0 and 7):
  • Higher H^+ concentration.
  • Lower OH^- concentration.
• Basic (or alkaline) pH (between 7 and 14):
  • Lower H^+ concentration.
  • Higher OH^- concentration.
• pH of human blood: Ranges from 7.35 to 7.45.
• pH scale:
  • Has an inverse relationship with H^+ concentration.
  • More H^+ ions mean lower pH, fewer H^+ ions mean higher pH.

Acids, Bases, and Salts

• Acid (proton donor): Any solute that adds hydrogen ions to a solution.
  • Strong acids dissociate completely in solution (e.g., HCl).
  • Weak acids do not dissociate completely in solution.
• Base (proton acceptor): Any solute that removes hydrogen ions from a solution.
  • Strong bases dissociate completely in solution (e.g., NaOH).
  • Weak bases do not dissociate completely in solution.
• Salt: Ionic compound that dissociates in water into cations and anions other than hydrogen ions and hydroxide ions.
• Buffers: Compounds that stabilize the pH of solutions.
  • Buffer systems often involve a weak acid and its related salt (weak base).
  • Can neutralize acids or bases and prevent fluctuations in pH.
  • The carbonic acid–bicarbonate buffer system is very important in humans.

Functional Groups of Organic Compounds

• Amino group (-NH2): Acts as a base, accepting H^+ depending on pH; can form bonds with other molecules; found in amino acids.
• Carboxyl group (-COOH): Acts as an acid, releasing H^+; found in fatty acids and amino acids.
• Hydroxyl group (-OH): May link molecules through dehydration synthesis (condensation); hydrogen bonding between hydroxyl groups and water molecules; affects solubility; found in carbohydrates, fatty acids, and amino acids.
• Phosphate group (-PO4): May link other molecules to form larger structures; may store energy in high-energy bonds; found in phospholipids, nucleic acids, and high-energy compounds.

Comparison of RNA and DNA

Classes of Inorganic and Organic Compounds