Chemistry

1. Chemistry • At the bottom, biology is nothing but applied chemistry • All matter is composed of atoms • Elements such as carbon and oxygen are a group of atoms of the same type. For instance, a nail made of iron is just a large group of iron atoms. • There are 92 naturally occurring elements, plus about 25 artificially-created elements. • Living things are mainly composed of the elements carbon, hydrogen, oxygen, and nitrogen. Another dozen or so elements are also used: phosphorus, iron, magnesium, sodium, potassium, chlorine, to name a few.

2. Atoms • Atoms have 3 components: protons, neutrons, and electrons • protons and neutrons are in the nucleus • Electrons circle around the nucleus • NOTE! This nucleus is NOT the same as the cell’s nucleus

3. Properties of Protons, Neutrons, and Electrons • Protons have a mass of 1 unit. If you had 600,000,000,000,000,000,000,000 protons you would have 1 gram. Protons also have an electrical charge of +1: protons are positively charged. • Neutrons also have a mass of 1 unit. Most of the mass of any object is in its protons and neutrons. Neutrons are neutral (not electrically charged). • Electrons weigh about 1/2000 of a unit, very light. They have a -1 charge (negatively charged).

4. proton protons neutrons Atomic Number and Weight • The atomic number of an atom is the number of protons it has. The atomic number defines the basic identity of the atom: what element it is. For example. All hydrogen atoms have 1 proton (atomic number = 1), ↑ and all carbon atoms have 6 protons (atomic number = 6). All gold atoms have 79 protons; uranium atoms have 92 protons.

5. atomic weight Atomic Number and Weight • The atomic weight of an atom is the number of protons plus the number of neutrons. Most hydrogen atoms have 1 proton and no neutrons, so the atomic weight is 1. Most carbon atoms have 6 protons and 6 neutrons, so the atomic weight is 12. • Electrons are very light, so they don’t add significantly to the atomic weight.

6. Isotopes • Isotopes are atoms with the same atomic number (same element) but different atomic weights: that is, isotopes have the same number of protons but different numbers of neutrons. • Different isotopes of the same element all have the same chemical properties. Chemical properties are determined by the number of protons only. • Heavier isotopes (too many neutrons) are usually radioactive.

7. Electron Shells • Atoms are usually neutral in electrical charge, which means that they have the same number of electrons (- charge) as protons (+ charge). • Electrons circle the nucleus at defined positions called shells. • The innermost shell of every atom holds 2 electrons (except hydrogen, which only has 1 electron). • All other shells hold up to 8 electrons • Electrons fill in from the nucleus out, so the inner shells are always full. • In most atoms, the outermost shell is not full. • Most chemistry is caused by the electrons of the outermost shell: atoms have a “desire” to have a full outer shell with 8 electrons in it.

8. Ions • The number of electrons in an atom is usually the same as the number of protons. Since electrons have a -1 charge and protons have a +1 charge, atoms are electrically neutral. • Ions are atoms where the number of electrons is different from the number of protons. • Ions have an electrical charge: positive charge if more protons than electrons, and negative charge if more electrons than protons. • For example, sodium atoms have 11 protons. Neutral sodium atoms also have 11 electrons, but sodium ions have only 10 electrons: 11 protons plus 10 electrons, giving a total charge of +1.

9. Chemical Bonds • Atoms can combine with each other to form molecules. • A molecule is a defined number of atoms grouped into a defined spatial relationship. For example, water, H2O, is 2 hydrogen atoms connected to an oxygen atom. The oxygen is in the middle, and the hydrogens are attached at an angle to it. • A large group of the same molecule is called a compound (just as a large group of the same atom is called an element). • Molecules are held together by chemical bonds. • Chemical bonds are the result of 2 forces: • 1. The octet rule, which means that atoms want to have 8 electrons in their outer shell (2 in the case of hydrogen). • 2. The attraction between atoms of opposite electrical charge. • The three main types of chemical bond are; ionic bond, covalent bond, and hydrogen bond.

10. Ionic Bonds • In an ionic bond, one atom gives an electron to another atom. This makes both atoms ions, and they are held together because their opposite charges attract each other. ↑ In sodium chloride (table salt), sodium starts out with 1 electron in its outer shell. The next shell down has 8 electrons, so by giving 1 electron away, the sodium atom gets a full outer shell. It then has a +1 charge. • Chlorine starts out with 7 electrons in its outer shell. By gaining one more electron, it gets 8 in the outer shell, and a -1 charge. ↑ The + charged sodium and the – charged chlorine attract each other, and they pack together in salt crystals.

11. Ionic Bonds, pt. 2 • Na = sodium atom; Cl = chlorine atom (called chloride when in a compound). • NaCl = one sodium plus one chlorine combined into a compound, sodium chloride (table salt). • The properties of sodium chloride are very different from sodium, a soft flammable metal, and chlorine, a poisonous gas. • Ionic bonds are very strong, but they are easily brokenby water.

12. Covalent Bonds • Covalent bonds occur when 2 atoms share a pair of electrons. The electrons spend part of their time with both atoms, so the octet rule is satisfied sufficiently. • A molecule of hydrogen gas, H2, has 2 hydrogen atoms. Each atom provides 1 electron, so in the bond each atom shares 2, a complete shell for hydrogen. • The bond is symbolized as a line connecting the 2 H’s: H-H ↑ In water (H2O), the oxygen has 6 electrons in its outer shell, and it shares one with each of the 2 hydrogens, giving 8 shared electrons for oxygen and 2 for each hydrogen. • Covalent bonds are the most common type in biological molecules.

13. Single, Double, and Triple Bonds • In a single bond, a pair of electrons (2 electrons) is shared. H2 gas and water are examples of this. Most covalent bonds are single bonds. ↑ In a double bond, 2 pairs of electrons (total of 4 electrons shared) are shared. In oxygen gas (O2), each atom has 6 electrons, a total of 12. Each atom contributes 2 electrons to the bond, so 4 are shared. Thus each atom has 4 unshared and 4 shared electrons, satisfying the octet rule. Carbon dioxide is another example. ↑ In a triple bond, 3 pairs of electrons (6 electrons) are shared. Nitrogen gas is an example. Cyanide (CN) is another example.

14. Polar Covalent Bonds • Sometimes the electrons in a covalent bond aren’t shared equally, because one atom attracts electrons more strongly than the other. When this happens, the electrons spend more time with one atom, and that atom becomes slightly negatively charged. The other atom becomes slightly positively charged. This is a polar covalent bond, because the atoms form positive and negative poles. • Rule: Carbon and hydrogen share electrons equally. Oxygen and nitrogen also share equally. But, oxygen and nitrogen attract electrons more strongly than carbon or hydrogen. • Water is a polar compound, because the oxygen is slightly negative and the hydrogens slightly positive. • Note that the total charge on the molecule is balanced, same number of electrons as protons, but within the molecule the charges are slightly separated. (Bonds where the electrons are shared equally are called non-polar.) • Polar molecules attract each other: the opposite charges attract.

15. Hydrogen Bonds • The slight + and – charges in polar bonds attract each other. In biological molecules, it is common for the partial + charge on a hydrogen to attract the partial – charge on a nearby oxygen or nitrogen. This attraction is called a hydrogen bond. ↑ Hydrogen bonds are very weak compared to covalent bonds, but large numbers of them can add up to a strong bond. The strands of DNA are held together by hydrogen bonds. • Hydrogen bonds also form between different parts of the same molecule, and between water and other molecules. • NaCl dissolves as H-bonds form between water molecules and the Na+ and Cl- ions. Water molecules surround each ion completely, separating the ions from the salt crystal and dispersing them throughout the water.

16. Water • All life occurs in water. Most molecules are dissolved in water: an aqueous solution. • Solution: a homogeneous mixture of 2 or more types of atom or molecule. • solute: what is being dissolved e.g. NaCl. • solvent: the liquid that is doing the dissolving e.g. water. • Water, H2O, is a polar compound. The 2 hydrogens are held at an angle to each other, and so the oxygen end of the molecule is partially negative and the hydrogen end is partially positive. ↑ Water forms many hydrogen bonds with other water molecules and with other polar substances. This causes water molecules to stick together (causing surface tension) and stick to other things (causing capillary action, how water gets from the roots to the top of trees).

17. Water • Polar substances dissolve in water, because water forms hydrogen bonds with the polar molecules. Thus, polar substances are called hydrophilic, or “water-loving”. ↑ Non-polar substances don’t dissolve in water because they can’t form hydrogen bonds, so they are called hydrophobic, or ‘water-fearing”. Oils and fats are examples of non-polar substances. Hydrophilic coating reduces friction by trapping a thin layer of water next to the boat’s hull.

18. Water • The membrane that surrounds each cell is hydrophobic, which is what keeps the water inside separated from the water outside.

19. Other Properties of Water • Water absorbs heat very effectively, slowing temperature changes within the organism. This is especially true when it evaporates, which makes it useful for cooling—sweating on a hot day. • Water expands when it freezes: ice floats, insulating the water below and preventing freezing all the way to the bottom. Many organisms live below the ice, protected from harsher conditions above.

20. Acids and Bases • Water dissociates into hydrogen ions (H+) and hydroxide ions (OH-). They ions then re-associate back into water: it’s a constant back-and-forth process, with an equilibrium between the H2O form and the H+ and OH- forms. • Acids produce H+ ions when dissolved in water • Bases absorb H+ ions when dissolved in water. A substance that releases OH- ions is a base because the OH- combine with the H+ in the water. • Water is neutral, neither acid nor base, because it always has equal numbers of H+ ions and OH- ions. • Both conditions can speed up various chemical reactions; acids and bases cause most of the chemical transformations that occur in living things. Much of the digestion that occurs in the stomach is due to strong acids (hydrochloric acid) there. ↑

21. Acids and Bases • To create an acidic condition, dissolve an ionic substance in water that releases H+ ions but not OH-. • Similarly, an ionic substance that releases OH- but not H+ is a base. The OH- absorbs the H+ ions already present in the water. An example is lye: sodium hydroxide: NaOH, which dissociates into Na+ and OH-. • Bases are also called alkaline. Ammonia and other compounds that contain NH2 also work as bases.

22. pH Scale • Acidity is measured on the pH scale, which indicates the amount of H+ ions present. The scale runs from 0 (very acidic) to 14 (very basic). • Water, which is neutral, has a pH of 7 • Acids have lower pHs: the hydrochloric acid in your stomach has a pH of about 2. Eating food stimulates your stomach to secrete more acid. Antacids (like Tums or Rolaids) neutralize some of this acid. • Bases have a higher pH: the lye in oven cleaner has a pH of about 14. • Body fluids are slightly basic, pH 7.4

23. Buffers and Salts • Too much acid or base is harmful. The body needs to protect itself against large pH shifts. It uses buffers, pairs of weak acids and weak bases to absorb excess H+ and OH- and keep the body’s pH near neutral. • The main buffer in the body is carbonic acid, which dissociates into H+ and HCO3-. If H+ is added by an acid, it gets converted into the neutral H2CO3. Similarly, excess OH- combines with the H+, leaving the much less basic HCO3-. These opposing reactions keep the pH at the proper level.

25. Cl- Na+ Buffers and Salts • Salts are ionic compounds that don’t release H+ or OH- when they dissolve. Thus, sodium chloride is a salt because it dissolves to form Na+ and Cl- , while hydrogen chloride (hydrochloric acid, HCl) is an acid because it dissolves to form H+ and Cl-. • Acids usually have an associated salt. An example of this is glutamic acid, an important component of proteins, and monosodium glutamate (MSG), which is used as a flavor enhancer in food. They are the same thing chemically, except that MSG has a sodium where glutamic acid has a hydrogen.