1. Matter, Elements, and Compounds

2. Matter: Anything that takes up space and has mass. There are 92 naturally occurring elements, of these 25 are essential to life. 4 of these make up 96% of living matter (Carbon, Hydrogen, Oxygen, and Nitrogen). The rest are called trace elements. These are required in minute amounts (Zinc, Cobalt, Iron, and Magnesium). Element: A substance that cannot be broken down by ordinary means. The material making up matter.

3. Atomic Number: The total number of protons in an atom. Atomic Mass: The total number of neutrons and protons in an atom. Small units of matter are called atoms. Protons (+), neutrons (0), and electrons (-) are the subunits of atoms. Their mass is measured in units called Daltons.

4. Isotopes: Different atomic forms caused by varying the number of neutrons. Example: Normal carbon is 12, carbon isotope is 14. Some isotopes are radioactive, they undergo a transformation to gain a stable condition. This transformation is called the half-life of the isotope.

5. Energy Levels: All electrons have the same mass and charge. Electrons closer to the nucleus contain less potential energy. The farther away from the nucleus, the more potential energy they contain.

6. Energy Levels These different states of potential energy are called energy levels. Each level is divided into subunits called orbitals. No more then 2 electrons can occupy the same orbital. The orbital closest to the nucleus contains 2 e-. The next level is divided into 4 orbitals for a total of 8 electrons.

7. Energy Levels Valence Electrons: These electrons occupy the last energy level of an atom. It stands to reason that chemical bonds will occur here in a chemical reaction, since it is here where atoms come in contact with each other.

8. Energy Levels The maximum number of valence electrons any atom can contain is 8. Any number less than 8 will allow that atom to take or give up electrons to become stable. Atoms that give electrons will become positive ions, while ions that receive electrons will become negative ions.

9. Chemical Bonding: Atoms will form bonds depending on their incomplete valence shell. There are several types of chemical bonds: –Covalent Bond: These bonds are the strongest bonds. They are formed by the sharing of valence electrons. –Ionic Bond: These bonds are formed by the taking of electrons. Anion: negative ion, Cl- and (OH). Cation: positive ion, Na+ and (NH4)+. –Hydrogen Bond: This bond is formed when hydrogen that is covalently bonded to an electronegative atom is attracted to another electronegative atom on another molecule.

10. Covalent Bonds

11. Ionic Bonds

12. Hydrogen Bonds

13. Chemical Reactions: The combination of 2 or more elements forming a different product or products. Each reaction contains reactants and products. The reactants are written on the left side of the equation, while the products are written on the right side. The reactants and products must contain the same number of atoms making the reaction balanced. Reactants  Products

15. Water and the Fitness of the Environment

16. Hydrogen Bonding of Water Molecules: Due to polar covalent bonds that hold a water molecule together, hydrogen bonds form where the negative oxygens and the positive hydrogens are located. The results of these bonds are as follows:

17. 1. Cohesion Cohesion is the sticking together of similar molecules. Water is very cohesive.

18. 2: Surface Tension Cohesion allows water to pull together and form droplets or form an interface between it and other surfaces. The measure of how hard it is to break this interface is surface tension.

19. 3: Adhesion Adhesion is the sticking of one substance to another. Water very adhesive. It will cling to many objects and act as a glue. Capillary action is an example of cohesion and adhesion working together to move water up a thin tube.

20. 4: Imbibition Imbibition is the process of soaking into a hydrophilic substance. Some examples are water being taken into a sponge, a seed, or paper towels.

21. 5: High Specific Heat Specific heat of a substance is the heat needed (gained or lost) to change the temperature of 1g of a substance 1 degree Celsius. A Kilocalorie equals 1,000 small calories. It takes 1,000 calories to raise 1,000g of water 1 degree Celsius. This high specific heat allows water to act as a heat sink. Water will retain its temperature after absorbing large amounts of heat, and retain its temperature after losing equally large amounts of heat. The reason for this is that hydrogen bonds must absorb heat to break. They must release heat when they form. The Ocean acts as a tremendous heat sink to moderate the earth’s temperature.

22. High Specific Heat Because of high specific heat capacity, water during the day is cooler than land. Rising air above warm land is replaced by cooler air pushed in from the lake. The reverse happens at night, when the land's temperature has fallen below that of the lake; the lake's temperature drops, too, at night, but not as much as the land's.

23. 6. High Heat of Vaporization Water must absorb a certain amount of additional heat to change from a liquid to a gas. This extra heat is called heat of vaporization. In humans, this value is 576 cal/g. This results in evaporative cooling of the surface. Alcohol has a value of 237 cal/g and chloroform 59 cal/g.

24. 7: Freezing and Expansion of Water Water is most dense at at 4 degrees C. At 0 degrees C it is less dense. Ice floats because maximum hydrogen bonding occurs at 0 degrees C.

25. 8: Versatile Solvent Water is a major solvent in nature. When water and another substance are mixed, the resulting solution is call an aqueous solution. Any solution that contains the following parts: Solute (what’s being dissolved) + Solvent (what’s doing the dissolving) = Solution. Solute Concentration: The concentration of the dissolved materials in relation to the solvent. This is always measured in moles.

26. 8: Versatile Solvent A mole is the amount of a substance whose mass in grams is numerically equivalent to its molecular weight in daltons. One must first find the atomic weights of the substances involved and add them together for the representative molecule and change the value to grams. Molarity occurs when the mole (gram atomic weight of the substance) is placed in a container nd dissolved in one liter of water.

27. pH pH: refers to the dissociation of water molecules. The pH constant is Kw = 1.0 x 10 -14 (mol/L) When we have an even split of H+ and OH- ions, the pH value is 7. Given the pH constant, in a solution with pH 7, the concentration of the H+ ion is 1 x 10 -7 and the concentration of the OH- ion is also 1 x 10 -7. The true definition of pH is pH = - log [H+]

29. Practice problems with pH Problems: 1.) H+ conc = 1 x 10 -10 mol/L. Determine the pH. 2.) OH- conc = 1 x 10 -2 mol/L. Determine the pH.

30. Buffers Abrupt changes in pH are harmful to the cell and any living organism. In order to minimize this harm, cells contain buffering systems. In order to change the pH of a solution, H ions must be added or taken from a solution Buffers maintain equilibrium by adding or removing H+ as needed.