A. Chemical Elements and Compounds
1. Matter consists of chemical elements in pure form and in combinations called compounds.
a. Organisms are composed of matter.
1. Matter is anything that takes up space and has mass.
b. An element is a substance that cannot be broken down to other substances by chemical reactions.
1. There are 92 naturally occurring elements. (Table 2.1)
2. Each element has a unique symbol, usually from the first one or two letters of the name, often from Latin or German.
c. A compound is a substance consisting of two or more elements in a fixed ratio. (Fig. 2.2)
1. Table salt (sodium chloride or NaCl) is a compound with equal numbers of chlorine and sodium atoms.
2. While pure sodium is a metal and chlorine is a gas, their combination forms an edible compound.
2. Life requires about 25 chemical elements (Table 2.1)
a. About 25 of the 92 natural elements are known to be essential for life.
Four elements - carbon (C), oxygen (O), hydrogen (H), and nitrogen (N) - make up 96% of living matter.
b. Most of the remaining 4% of an organism�s weight consists of phosphorus (P), sulfur (S), calcium (Ca), and potassium (K).
c. Trace elements are required by an organism but only in minute quantities.
1. Some trace elements, like iron (Fe), are required by all organisms.
2. Other trace elements are required only by some species.
a. For example, a daily intake of 0.15 milligrams of iodine is required for normal activity of the human thyroid gland. (Fig. 2.4)
B. Atoms and Molecules
1. Atomic structure determines the behavior of an element.
a. Subatomic particles
1. An atom is the smallest unit of matter that still retains the properties of an element.
2. Atoms are composed of even smaller parts, called subatomic particles.
a. Two of these, neutrons and protons, are packed together to form a dense core, the atomic nucleus, at the center of an atom.
b. Electrons orbit the nucleus.
3. Each electron has one unit of negative charge. (Fig. 2.5)
4. Each proton has one unit of positive charge.
5. Neutrons are electrically neutral.
6. The attractions between the positive charges in the nucleus and the negative charges of the electrons keep the electrons in the vicinity of the nucleus.
7. A neutron and a proton are almost identical in mass, about 1.7 x 10-24 gram per particle.
8. The mass of an electron is about 1/2000th that of a neutron or proton.
b. Atomic number and mass number
1. All atoms of a particular element have the same number of protons in their nuclei. (Table 2.1)
a. Each element has a unique number of protons, its atomic number.
b. The atomic number is written as a subscript before the symbol for the element (for example, 2He).
2. Unless otherwise indicated (see ions below), atoms have equal numbers of protons and electrons - no net charge.
a. Therefore, the atomic number tells us the number of protons and the number of electrons that are found in a neutral atom of a specific element.
3. The mass number is the sum of the number of protons and neutrons in the nucleus of an atom.
a. Therefore, we can determine the number of neutrons in an atom by subtracting the number of protons (the atomic number) from the mass number.
b. The mass number is written as a superscript before an element�s symbol (for example, 4He).
c. Energy levels of electrons
1. If the nucleus of an atom were the size of a golf ball, the electrons would be about 1 kilometer from the nucleus.
2. When two elements interact during a chemical reaction, it is actually their electrons that are involved.
3. Electrons have potential energy because of their position relative to the nucleus.
a. The negatively charged electrons are attracted to the positively charged nucleus.
b. The farther electrons are from the nucleus, the more potential energy they have.
4. However, electrons cannot occupy just any location away from the nucleus.
a. Changes in potential energy can only occur in steps of a fixed amount, moving the electron to a fixed location.
b. An electron cannot exist between these fixed locations.
5. The different states of potential energy that the electrons can have are called energy levels or electron shells. (Fig. 2.9)
a. The first shell, closest to the nucleus, has the lowest potential energy.
b. Electrons in outer shells have more potential energy.
c. Electrons can only change their position if they absorb or release a quantity of energy that matches the difference in potential energy between the two levels.
d. Chemical behavior
1. The chemical behavior of an atom is determined by its electron configuration - the distribution of electrons in its electron shells.
2. The first 18 elements, including those most important in biological processes, can be arranged in 8 columns and 3 rows. (Fig. 2.10)
a. Elements in the same row use the same shells.
b. Moving from left to right, each element has a sequential addition of electrons (and protons).
3. The first electron shell can hold only 2 electrons.
4. Atoms with more than two electrons must place the extra electrons in higher shells.
5. The second shell can hold up to 8 electrons.
6. The chemical behavior of an atom depends mostly on the number of electrons in its outermost shell, the valence shell.
a. Electrons in the valence shell are known as valence electrons.
7. Atoms with the same number of valence electrons have similar chemical behavior.
a. An atom with a completed valence shell is unreactive.
b. All other atoms are chemically reactive because they have incomplete valence shells.
e. Electron orbitals are not all spherical. (Fig. 2.11)
2. Atoms combine by chemical bonding to form molecules.
a. General
1. Atoms with incomplete valence shells interact by either sharing or transferring valence electrons.
2. These interactions result in the atoms remaining close together, held by attractions called chemical bonds.
b. Covalent bonds
1. A covalent bond is the sharing of a pair of valence electrons by two atoms. (Fig. 2.12)
a. If two atoms come close enough that their unshared orbitals overlap, each atom can count both electrons toward its goal of filling the valence shell.
b. For example, if two hydrogen atoms come close enough that their orbitals overlap, then they can share the single electron that each contributes.
2. Two or more atoms held together by covalent bonds constitute a molecule.
3. Structural formula-abbreviate the structure of a molecule by substituting a line for each pair of shared electrons.
a. H-H is the structural formula for the covalent bond between two hydrogen atoms.
4. Molecular formula-indicates the number and types of atoms present in a single molecule. (For example, H2)
5. Double bond-oxygen needs to add 2 electrons to the 6 already present to complete its valence shell.
a. Two oxygen atoms can form a molecule by sharing two pairs of valence electrons.
b. These atoms have formed a double covalent bond.
6. Every atom has a characteristic total number of covalent bonds that it can form - an atom�s valence.
a. The valence of hydrogen is 1.
b. Oxygen is 2.
c. Nitrogen is 3.
d. Carbon is 4.
7. Covalent bonds can form between atoms of different elements.
a. Water is a compound in which two hydrogen atoms form single covalent bonds with an oxygen atom-satisfies the valences of both elements.
8. If electrons in a covalent bond are shared equally, then this is a nonpolar covalent bond.
a. A covalent bond between two atoms of the same element is always nonpolar. (O2)
b. A covalent bond between different atoms can also be nonpolar (CH4).
9. If the electrons in a covalent bond are not shared equally by the two atoms, then this is a polar covalent bond (H2O). (Fig. 2.13)
a. Compounds with a polar covalent bond have regions that have a partial negative charge near the larger atom and a partial positive charge near the smaller atom. (Covalent bonds movie-textbok activity 2E)
c. Ionic bonds (Fig. 2.14)
1. An ionic bond can form if two atoms are so unequal in their attraction for valence electrons that one atom strips an electron completely from the other.
a. Sodium with one valence electron in its third shell transfers this electron to chlorine with 7 valence electrons in its third shell.
b. Now, sodium has a full valence shell (the second) and chlorine has a full valence shell (the third).
c. After the transfer, both atoms are no longer neutral, but have charges and are called ions.
2. Atoms with positive charges are cations.
3. Atoms with negative charges are anions.
4. Because of differences in charge, cations and anions are attracted to each other to form an ionic bond. (Fig. 2.15)
5. Compounds formed by ionic bonds are ionic compounds or salts, like NaCl.
6. The formula for an ionic compound indicates the ratio of elements in a crystal of that salt. (Ionic bonds movie-textbok activity 2G)
3. Weak chemical bonds play important roles in the chemistry of life.
a. Hydrogen bonds form when a hydrogen atom that is already covalently bonded to a nitrogen or oxygen atom is attracted to another nitrogen or oxygen atom.
1. These bonds result because the polar covalent bond with hydrogen leaves the hydrogen atom with a partial positive charge and the nitrogen or oxygen atom with a partial negative charge.
2. Example: ammonia and water (Fig. 2.16)
b. Even molecules with nonpolar covalent bonds can have partially negative and positive regions.
1. Because electrons are constantly in motion, there can be periods when they accumulate by chance in one area of a molecule.
2. This creates ever-changing regions of negative and positive charge within a molecule.
4. Chemical reactions make and break chemical bonds.
a. In chemical reactions chemical bonds are broken and reformed, leading to new arrangements of atoms.
b. The starting molecules in the process are called reactants and the end molecules are called products.
c. In a chemical reaction, all of the atoms in the reactants must be accounted for in the products.