Chapter 2: Chemical Foundations for Cells  (The Chemistry Inside You!) Leafy Clean-up Crews: plants take pollutants into their tissues! A. Life depends on chemical reactions! B. Phytoremediation is the use of living plants to withdraw harmful substances from the environment. I. What are Atoms? A. Structure of Atoms 1. An atom is the smallest unit of matter that is unique to a particular element. 2. 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&emdash;) have a negative charge. Their quantity is equal to that of the protons. They move around the nucleus. 3. Atomic numbers and mass numbers give us an idea of whether and how substances will react. B. Isotopes&endash;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 or to diagnose disease.   II. Focus On Science: Using Radioisotopes to Track Chemicals and Save Lives: Thyroid gland and radioactive Iodine III. 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. B. The Nature of Chemical Bonds 1. A chemical bond is a union between the electron structures of atoms. 2. Atoms with "unfilled" orbitals in their outermost shell tend to be reactive with other atoms. 3. The number or the distribution of its electrons changes when an atom gives up, gains, or shares electrons. C. 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. IV. Important 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&endash;an ion. 2. In an ionic bond, (+) and (&emdash;) ions are linked by mutual attraction of opposite charges&endash;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. V. Properties of Water 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 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. D. 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. VI. Acids, Bases, and Buffers 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 Differ From Bases? 1. A substance that releases hydrogen ions (H+) in solution is an acid&endash;for example, HCl. 2. Substances that release ions such as (OH-) that can combine with hydrogen ions are called bases. C. 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. D. Salts 1. A salt is an ionic compound formed when an acid reacts with a base; example: HCl + NaOH &endash;&endash;> NaCl + H2O. 2. Salts dissociate into useful ions (examples: Na+ and Ca++) in body fluids.