Chapter 2 - Life's Chemical Basis
| 1. | Understand how protons, electrons, and neutrons are arranged into atoms and ions. |
| 2. | Explain how the distribution of electrons in an atom or ion determines the number and kinds of chemical bonds that can be formed. |
| 3. | Know the various types of chemical bonds, the circumstances under which each forms, and the relative strengths of each type. |
| 4. | Understand the essential chemistry of water and of some common substances dissolved in it. |
| 5. | Understand the relationships of acids, bases, and salts. |
Impacts, Issues: What Are You Worth?
A. Chemically, how much is a human body really worth?
1. The body is a collection of elements.
2. The smallest units of elements are atoms.
B. Quantities of elements in the body vary.
1. Oxygen, hydrogen, carbon, and nitrogen are the most abundant.
2. Trace elements include selenium, mercury, arsenic, and lead.
C. All of the elements in the human body could be bought for $118.63.
2.1 Start With Atoms
A. An atom is the smallest unit of matter that is unique to a particular element.
1. Atoms are composed of three particles:
a. Protons (p+) are part of the atomic nucleus and have a positive charge. Their quantity is called the atomic number (unique for each element).
b. Neutrons are also a part of the nucleus; they are neutral. Protons plus neutrons = atomic mass.
c. Electrons (e–) have a negative charge. Their quantity is equal to that of the protons. They move around the nucleus.
2. Atomic numbers and mass numbers give us an idea of whether and how substances will react.
B. The periodic table is an arrangement of elements based on their chemical properties.
1. Those in the same column of the table have the same number of electrons available for interactions with other atoms.
2. This allows chemists to predict the chemical behavior of an element.
2.2 Radioisotopes
A. Isotopes are variant forms of atoms.
1. Atoms with the same number of protons (for example, carbon with six) but a different number of neutrons (carbon can have six, seven, or eight) are called isotopes (12C, 13C, 14C ).
2. Some radioactive isotopes are unstable and tend to decay into more stable atoms.
a. They can be used to date rocks and fossils.
b. Some can be used as tracers to follow the path of an atom in a series of reactions.
B. Radioisotopes are used in PET to follow a particular chemical in the body and determine its patterns of metabolism.
2.3 What Happens When Atom Bonds With Atom?
A. Electrons and Energy Levels
1. Electron behavior influences atom bonding.
a. Electrons are attracted to protons but are repelled by other electrons.
b. Orbitals are like volumes of space around the atomic nucleus in which electrons are likely to be at any instant.
c. Each orbital contains one or two electrons.
2. Orbitals can be thought of as occupying shells around the nucleus.
a. The shell closest to the nucleus has one orbital holding a maximum of two electrons.
b. The next shell can have four orbitals with two electrons each for a total of eight electrons.
3. A chemical bond is a union between the electron structures of atoms.
a. Atoms with “unfilled” orbitals in their outermost shell tend to be reactive with other atoms.
b. The number or the distribution of its electrons changes when an atom gives up, gains, or shares electrons.
B. From Atoms to Molecules
1. A molecule is a bonded unit of two or more (same or different) atoms.
2. A compound is a substance in which the relative percentages of two or more elements never vary.
3. In a mixture, two or more elements simply intermingle in proportions the can vary.
2.4 Bonds in Biological Molecules
A. Ion Formation and Ionic Bonding
1. When an atom loses or gains one or more electrons, it becomes positively or negatively charged—an ion.
2. In an ionic bond, (+) and (–) ions are linked by mutual attraction of opposite charges—for example, NaCl.
B. Covalent Bonding
1. A covalent bond holds together two atoms that share one or more pairs of electrons.
2. In a nonpolar covalent bond, atoms share electrons equally.
3. In a polar covalent bond, because atoms share the electron unequally, there is a slight difference in charge between the two poles of the bond; water is an example.
C. Hydrogen Bonding
1. In a hydrogen bond, an atom of a molecule interacts weakly with a hydrogen atom already taking part in a polar covalent bond.
2. These bonds impart structure to liquid water and stabilize nucleic acids and other large molecules.
2.5 Water’s Life-Giving Properties
A. Polarity of the Water Molecule
1. Water is a polar molecule because of a slightly negative charge at the oxygen end and a slightly positive charge at the hydrogen end.
2. Water molecules can form hydrogen bonds with each other.
3. Polar substances are hydrophilic (water loving); nonpolar ones are hydrophobic (water dreading) and are repelled by water.
B. Water's Temperature-Stabilizing Effects
1. Water tends to stabilize temperature because it can absorb considerable heat before its temperature changes.
2. In evaporative processes the input of heat energy increases the molecular motion so much that hydrogen bonds are broken and water molecules escape into the air, thus cooling the surface.
3. In freezing, the hydrogen bonds resist breaking and lock the water molecules in the bonding patterns of ice.
C. Water's Solvent Properties
1. The solvent properties of water are greatest with respect to polar molecules with which they interact.
2. “Spheres of hydration” are formed around the solute (dissolved) molecules.
D. Water's Cohesion
1. Hydrogen bonding of water molecules provides cohesion (capacity to resist rupturing), which imparts surface tension.
2. Cohesion is especially important in pulling water through plants.
2.6 Acids and Bases
A. The pH Scale
1. pH is a measure of the H+ concentration in a solution; the greater the H+ the lower the pH scale.
2. The scale extends from 0 (acidic) to 7 (neutral) to 14 (basic).
B. How Do Acids and Bases Differ?
1. A substance that releases hydrogen ions (H+) in solution is an acid—for example, HCl.
2. Substances that release ions such as (OH–) that can combine with hydrogen ions are called bases.
C. Salts and Water
1. A salt is an ionic compound formed when an acid reacts with a base; example: HCl + NaOH à NaCl + H2O.
2. Salts dissociate into useful ions (examples: Na+ and Ca++) in body fluids.
D. Buffers Against Shifts in pH
1. Buffer molecules combine with, or release, H+ to prevent drastic changes in pH.
2. Bicarbonate is one of the body’s major buffers.
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Introduction Chemistry helps us understand the nature of all substances that make up cells, organisms, and the Earth, its waters, and the atmosphere. Section 2.1 All substances consist of one or more elements. Ninety-two elements occur natrually. These and others are shown in the periodic table. An atom consists of one or more positively charged protons, negatively charged electrons, and (except for hydrogen atoms) one or more uncharged neutrons. Protons and neutrons occupy the core region, or nucleus. All atoms of an element have the same atomic number, but mass number can vary. Section 2.2 Most (and probably all) elements have one or more naturally occurring isotopes. Some isotopes are stable; most are radioactive. The radioisotopes undergo radioactive decay and are often used as tracers in biological studies. Section 2.3 Whether an atom interacts with others depends on the number and arrangement of its electrons, which occupy orbitals (volumes of space) around the atomic nucleus. When an atom has one or more vacancies in orbitals at its highest energy level, it can interact with other atoms by the process of chemical bonding to form molecules. Proportions of elements are always the same in compounds, but can differ in mixtures. Section 2.4 An atom may lose or gain one or more electrons and thus become an ion, which has a positive or negative charge. Generally, a chemical bond is a union between the electron structures of atoms.
Section 2.5 Water's properties make life possible. Polar covalent bonds join together three atoms in a water molecule (two hydrogens and one oxygen). The water molecule's polarity invites extensive hydrogen bonding between molecules in bodies of water. Such bonding is the basis of liquid water's ability to resist temperature changes (more than other fluids do), display internal cohesion, easily dissolve hydrophilic substances, and repel hydrophobic ones. Any substance dissolved in water is a solute. Section 2.6 The pH of a solution indicates its hydrogen ion concentration. A typical pH range is from 0 (highest H+ concentration, most acidic) to 14 (lowest H+ concentration, most basic). At pH 7, or neutrality, H+ and OH- concentrations are equal. Acids release H+ ions in water; bases combine with them. A salt forms when an acid and base combine. Buffer systems help maintain a favorable pH in internal environments. This is important because most biological processes operate only within a narrow range of pH.
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