Chapter
2: Atoms, Molecules, and Water
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One of the hardest concepts to get across in this chapter is the mechanisms involved in the different types of bonds.
Ionic bonds are the easiest to understand in that they have the simplest structure. In an ionic bond one atom has a very high attraction for electrons, while the other has a very weak ability to retain its few outer shell electrons. It is important to remember that we are only concerned with the outer shell in these atoms, any filled shells are considered stable and out of play. Ionic bonds occur between atoms with a large vertical separation on the periodic table, typically between atoms from column IA or IIA and those in VIIA (such as Na and Cl). The VIIA atoms have a very strong pull on electrons, a property called electronegativity, while those in IA and IIA have only one or two weakly attached outer shell electrons. The ionic bond results from the strong attraction of the VIIA atom pulling one of the weakly held electrons of a IA or IIA atom away, resulting in ions. The VIIA atom becomes a negatively charged ion and the IA or IIA atom becomes a positively charged ion. These opposite charges attract each other (similar to putting two magnets next to each other with opposite poles) and this attraction is the ionic bond.
Covalent bonds are about sharing electrons. The most important concept to get here is that electrons enter the shells in a specific order, forming first single electrons, and then pairs. Pairs of electrons are more stable than singles. A covalent bond is when two atoms contribute one electron each to form a pair, they share the two electrons between them. To add a wrinkle to this nice easy picture we have to introduce the concept of polar and nonpolar bonds. If the atoms involved in the bond are the same element, they will both exert the same pull on the shared electrons, resulting in an equal sharing. This equal sharing results in an equal distribution of the charges involved and is called a nonpolar covalent bond. However, if the elements involved are different, they will exert an unequal force on the electrons and produce a lopsided sharing arrangement. This lopsided sharing of electrons results in an unequal charge separation that is referred to as a polar covalent bond. Covalent bonds are on a continuum between very polar covalent bonds (high charge separation, though NOT ionic) and completely nonpolar bonds (between two identical elements). The degree of polar nature can be determined by how separated elements are on the periodic table, the more separated the more polar the bond between them as a general rule.
Hydrogen bonds are the result of polar covalent bonds. Highly polar bonds (such as O-H or N-H) result in a partial charge forming on each side of the bond. These are not true ions, rather a mild charge we refer to as a partial charge (denoted either as δ+ or δ-). These partial charges are attracted to their opposite charge and result in very weak, transient bonds between adjacent O-H and N-H bonds. Individual hydrogen bonds are very weak, but in mass are one of the strongest forces in biology. This importance is because all of biology is based on water, which is nothing but O-H bonds. In a quantity of water even if a million transient hydrogen bonds break a million more can instantly form.
I. What are atoms?
A. Atoms and their substructure
1. Atoms are the fundamental structural units of matter
2. Atoms are composed of still smaller structures
a. Atomic nucleus (plural: atomic nuclei)
i. Protons – positively charged particles
ii. Neutrons – electrically neutral particles
b. Electrons – negatively charged particles that orbit the atomic nucleus.
3. A stable atom has an equal number of protons and electrons (electrically neutral)
4. There are 92 naturally occurring types of atoms, called elements.
a. Each element has a specific number of protons (i.e. hydrogen is 1, carbon is 6) – called the atomic number
b. Atoms of the same element may have differing numbers of neutrons; these different atoms are called isotopes.
5. Some atoms spontaneously break apart releasing energetic particles, a property called radioactivity.
B. Electrons, orbits, and energy levels
1. Electrons occupy areas close to the nucleus
2. There are discrete energy levels that the electrons can occupy
3. Only a set number of electrons can occupy any one level
4. These three dimensional orbits are called electron shells
5. Electrons are responsible for the interactions between atoms, also called bonds.
6. Only the outer most shells can interact with other atoms.
II. How do Atoms Interact to form Molecules?
A. Atoms react with other atoms when there are vacancies in their outermost electron shells.
1. Two or more atoms held together by interactions of their outermost electron shells are called a molecule.
2. Only atoms whose outermost electron shells have vacancies will bond with other atoms. These atoms are termed reactive.
3. Atoms whose outer shells are full are termed inert and they do not readily react with other atoms.
4. An atom is in a chemical bond as the result of gaining, losing, or sharing electrons with another atom.
5. Chemical reactions are the making and breaking of chemical bonds to create new molecules.
6. There are three types of chemical bonds.
a. Ionic Bond
b. Covalent Bond
c. Hydrogen Bond
B. Ionic Bonds
1. Ions (singular: ion) are atoms that have lost or gained electrons, upsetting the balance between protons and electrons in the atom.
2. Ions either have a net positive charge (protons>electrons) or a net negative charge (electrons>protons).
3. The electrical attraction between oppositely charged ions can result in an ionic bond.
4. A common example of this type of bond is table salt, sodium chloride (NaCl). A positively charged sodium atom is in an ionic bond with a negatively charged chlorine atom.
5. These types of bonds mostly occur with the atoms that have outer shells that are either almost full or almost empty.
C. Covalent Bond
1. A covalent bond is one where two atoms share electrons in their outer shells.
2. Covalent bonds are stronger than ionic bonds.
3. Some bonds can share more than one electron from each atom, resulting in double and triple bonds.
D. Most biological molecules utilize covalent bonds.
E. Polar and Nonpolar Covalent Bonds
1. If all nuclei of a molecule have roughly the same positive charge than the electrons are shared equally, giving a nonpolar covalent bond.
2. If the nuclei of a molecule are significant different in charge you will get the more positively charged atom pulling the electrons towards it, resulting in an unequal sharing of electrons. This unequal sharing of electrons gives a net electrical pole to the bond and it is called a polar covalent bond.
F. Free Radicals
1. A free radical is an atom that has one or more unpaired electrons in their outer shells.
2. Free radicals cause extensive damage to biological molecules because they are very unstable and will react with nearby molecules.
3. Free radical damage has been implicated in a number of diseases and conditions.
4. A variety of materials can cause free radical formation in the body including: chemicals from exhaust, radiation, and heavy metals.
5. Antioxidants are chemicals that react with free radicals and render them harmless.
G. Hydrogen bonds
1. Hydrogen bonds are the result of electrical attractions between opposite poles of polar covalent bonds.
2. The strongest hydrogen bonds occur between water molecules.
3. The hydrogen bond, though weak, is essential to biology and can be very strong when a number of them are working together.
III. Why is water so important to life?
A. Water interacts with many other molecules
1. Water is a good solvent, a liquid with the capacity to dissolve a wide variety of other molecules.
2. Water dissolves ionic molecules by interacting with the charged components and shielding them from their opposite charge.
3. Water also dissolves molecules with polar covalent bonds through electrostatic interactions.
4. Ionic and Polar covalent materials are called hydrophilic (“water loving”) due to their ability to dissolve in water.
5. Molecules that are uncharged and nonpolar are considered hydrophobic (“water-fearing”) due to their inability to dissolve in water.
6. Hydrophobic molecules tend to aggregate in water, pushing there most hydrophobic sections inside the aggregate away from water; this is described as hydrophobic interaction.
B. Water molecules stick together
1. Because of hydrogen bonds, water molecules tend to have a tendency to stick together, this is termed high cohesion.
2. Cohesion is exemplified at the surface of water by its tendency to remain unbroken and support a significant amount of weight. This tendency not to be broken is called surface tension.
3. Water also exhibits adhesion, the tendency to stick to polar surfaces.
C. Water can form acids, bases, or be neutral
1. Some water molecules in a collection are in ionic form, one positively charged hydrogen ion (H+), and one negatively charged hydroxide ion (OH-).
2. An equal number of hydroxide to hydrogen ions makes a solution neutral.
3. An excess of hydrogen ions makes a solution an acid.
4. An excess of hydroxide ions makes a solution a base.
5. The degree to which a solution is acidic, basic or neutral is registered on a pH scale, where the value of 7 represents neutral, greater than 7 is basic, and less than 7 is acidic. For example, Coke has a pH of 2, and is therefore an acid, while Clorox has a pH of 12 and is a base.
6. Buffers are chemicals that can be added to water to help maintain a constant pH. They release or absorb hydrogen ions in response to changes in the hydrogen ion concentration of the solution.
D. Water moderates the effects of temperature change
1. Water has a high specific heat.
a. Specific heat is defined as the amount of energy required to raise the temperature of one gram of material one degree Celsius.
b. For water this is defined as one calorie.
c. In comparison it takes 0.2 calorie to affect salt in a similar manner, and only 0.02 calories for something such as granite.
2. Water has a high heat of vaporization, the energy required to turn one gram of a material from liquid to gas.
3. Water has a high heat of fusion, the energy required to be released to turn one gram of a material from liquid to solid.
E. Water forms an unusual solid.
1. Ice is less dense than water; in contrast most other materials become denser when they solidify.
2. This has implications when you think about ponds freezing in the winter.