Home

Class Schedule

Syllabus

Lessons

Quizzes

Links

Grass Home

E-mail Dr. Grass

Basic Chemistry

Concepts of Matter and Energy

1. Matter
   1. Matter is anything that occupies space and has weight.

   2. Matter exist in three states: gas, liquid, and solid.

2. Energy
   1. Energy is the capacity to do work or to put matter into motion. Energy has kinetic (active) potential (stored) work capacities.

   2. Types of energy that are important in body functions include, chemical, electrical, mechanical, and radiant.

   3. Energy can be converted from one form into another, but some energy is always unuseable (lost as heat) in such transformations.

Return to top

Composition of Matter

1. Elements and atoms
   1. Each element is a unique substance that can not be decomposed into simpler substances by ordinary chemical methods. A total of 112 elements exist; they differ from one another in their chemical and physical properties.

   2. Four elements (carbon, hydrogen, oxygen, and nitrogen) make up 96% of living matter. Several other elements are present in small or trace amounts.

   3. The building blocks of elements are atoms. Each atom is described by an atomic symbol consisting of one or two letters.

2. Atomic structure

   1. Atoms are composed of three subatomic particles: protons, electrons, and neutrons. Because all atoms are electrically neutral, the number of protons in any atom is equal to its number of electrons.

   2. The planetary model of the atom portrays all the mass of the atom (protons and neutrons) concentrated in a minute central nucleus (Figure 2.1). Electrons orbit the nucleus along specific orbits. The orbital model also locates protons and electrons in a central nucleus, but it depicts electrons as occupying areas of space called orbitals and forming an electron cloud of negative charge around the nucleus.

   3. Each atom can be identified by an atomic number, which is equal to the number of protons contained in the atom's nucleus.

   4. The atomic mass number is equal to the sum of the protons and neutrons in the atom's nucleus.

   5. Isotopes are different atomic forms of the same element; they differ only in the number of neutrons in the nucleus. Many of the heavier isotopes are unstable and decompose to a more stable form by ejecting particles of energy from the nucleus, a phenomenon called radioactivity. Such radioisotopes are useful in medical diagnosis and treatment and in biochemical research.

   6. The atomic weight is approximately equal to the mass number of the most abundant isotope of any element.

Return to top

Molecules and Compounds

1. A molecule is the smallest unit resulting from the binding of two or more atoms. If the atoms are different, a molecule of a compound is formed.

2. Compounds exhibit properties different from those of the atoms they comprise.

Return to top

Chemical Bonds and Chemical Reactions

1. Bond formation
   1. Chemical bonds are energy relationships. Electrons in the outermost energy level (valence shell) of the reacting atoms are active in the bonding.

   2. Atoms with a full valence shell (2 electrons in shell 1, or 8 in the subsequent shells) are chemically inactive. Those with an incomplete valence shell interact by losing, gaining, or sharing electrons to achieve stability (that is, to fill the valence shell).

   3. Ions are formed when valence-shell electrons are completely transferred from one atom to another. The oppositely charged ions formed attract each other, forming an ionic bond. Ionic bonds are common in salts (Figure 2.2).

   4. Covalent bonds involve the sharing of electron pairs between atoms (Figure 2.3). If the electrons are shared equally, the molecule is a nonpolar covalent molecule. If the electrons are not shared equally, the molecule is a polar covalent molecule. Polar molecules orient themselves toward charged particles.

   5. Hydrogen bonds are fragile bonds that bind together different parts of the same molecule (intramolecular bonds). They are common in large, complex organic molecules, such as proteins and nucleic acids and between water molecules (Figure 2.4).

2. Patterns of chemical reactions
   1. Chemical reactions involve the formation or breaking of chemical bonds. They are indicated by the writing of a chemical equation, which provides information about the atomic composition (formula) of the reactant(s) and product(s).

   2. Chemical reactions that result in larger, more complex molecules are synthesis reactions; they involve bond formation (Figure 2.5).

   3. In decomposition reactions, larger molecules are broken down into simpler molecules or atoms. Bonds are broken (Figure 2.6).

   4. Exchange reactions involve both the making and breaking of bonds. Atoms are replaced by other atoms (Figure 2.7).

Return to top

Biochemistry: The Chemical Composition of Living Matter

1. Inorganic compounds
   1. Inorganic compounds making up living matter do not contain carbon. They include water, salts, acids, and bases.

   2. Water is the single most abundant compound in the body. It acts as a universal solvent in which electrolytes (salts, acids, and bases) ionize and in which chemical reactions occur, and it is the basis of transport and lubricating fluids. It slowly absorbs and releases heat, thus helping to maintain homeostatic body temperature, and it protects certain body structures (e.g., the brain) by forming a watery cushion. Water is also a reactant in hydrolysis reactions.

   3. Salts in ionic form are involved in nerve transmission, muscle contraction, blood clotting, transport of oxygen by hemoglobin, cell permeability, metabolism, and many other reactions. Additionally, calcium salts (as bone salts) contribute to bone hardness (Figure 2.8).

   4. Acids are proton donors. When dissolved in water, they release hydrogen ions (H+). Strong acids dissociate completely; weak acids dissociate incompletely.

   5. Bases are proton acceptors. The most important inorganic bases are hydroxides (OH-). Bicarbonate ions are important bases in the body. When bases and acids interact, neutralization occur that is, a salt and water are formed.

   6. pH is a measure of the relative concentrations of hydrogen and hydroxyl ions in various body fluids. Each change of one pH unit represents a 10-fold change in hydrogen (or hydroxyl) ion concentration. A pH of 7 is neutral (that is, the concentrations of hydrogen and hydroxyl ions are equal). A pH below 7 is acidic; a pH above 7 is alkaline (basic) (Figure 2.9).

   7. Normal blood pH ranges from 7.35 to 7.45. Slight deviations outside this range can be fatal.

2. Organic compounds
   1. Organic compounds are the carbon-containing compounds that living matter comprises. Carbohydrates, lipids, proteins, and nucleic acids are examples. They all contain carbon, oxygen, and hydrogen. Proteins and nucleic acids also contain substantial amounts of nitrogen.

   2. Carbohydrates contain carbon, hydrogen, and oxygen in the general relationship (CH2O) their building blocks are monosaccharides. Monosaccharides include glucose, fructose, galactose, deoxyribose, and ribose, disaccharides include sucrose, maltose, and lactose; and polysaccharides include starch and glycogen (Figure 2.10). Carbohydrates are ingested as sugars and starches. Carbohydrates, and in particular glucose, are the major energy source for the formation of ATP.

   3. Lipids include the neutral fats or triglycerides (glycerol plus three fatty acid chains) (Figure 2.11), phospholipids, and steroids (most importantly, cholesterol). Neutral fats are found primarily in adipose tissue, where they provide insulation and reserve body fuel. Phospholipids and cholesterol are found in all cell membranes. Cholesterol also forms the basis of certain hormones, bile salts, and vitamin D. Like carbohydrates, the lipids are degraded by hydrolysis and synthesized by dehydration synthesis.

   4. Proteins are constructed from building blocks called amino acids; 20 common types of amino acids are found in the body. Amino acid sequence determines the proteins constructed (Figure 2.12). Fibrous, or structural, proteins are the basic structural materials of the body. Globular proteins are functional molecules; examples of these include enzymes, some hormones, and hemoglobin. Disruption of the hydrogen bonds of functional proteins leads to their denaturation and inactivation.

   5. Enzymes increase the rates of chemical reactions by combining specifically with the reactants and holding them in the proper position to interact. They do not become part of the product. Many enzymes are produced in an inactive form or are inactivated immediately after use (Figure 2.13).

   6. Nucleic acids include deoxyribonucleic acid (DNA) and ribonucleic acid (RNA). The building unit of nucleic acids is the nucleotide; each nucleotide consists of a nitrogenous base, a sugar (ribose or deoxy ribose), and a phosphate group. DNA (the "stuff" of the genes) maintains genetic heritage by replicating itself before cell division and contains the code-specifying protein structure is a double-stranded helix (Figure 2.14). RNA acts in protein synthesis to ensure that instructions of the DNA are executed and is single-stranded (Figure 2.15).

   7. ATP (adenosine triphosphate) is the universal energy compound used by all cells of the body. When energy is liberated by the oxidation of glucose, some of that energy is captured in the high-energy phosphate bonds of ATP molecules and is stored for later use (Figure 2.16).