Lecture #1 ~ Date__________

Chapter 1: Themes in the Study of Life

Chapter 2:Chemical Context of Life

Chapter 3:Water and the Fitness of the Environment

Themes of AP Biology

Evolution: Organisms, molecules, and the Earth change over time and this is the fundamental idea behind all studies of biology.

Science as a Process: Inquiry and ongoing need for better explanations are the root of all scientific studies.

Energy Transfer: Energy provided by the sun cycles through molecules and organisms.

Continuity and Change: Life and life processes have the ability to be consistent and stay the same and to change over time to adapt.

Relationship of Structure to Function: The structure of a molecule, part, or organism is integral to the functioning of that aspect of life.

Regulation: Organisms have the ability to adapt to changes in the environment.

Interdependence in Nature: Abiotic factors and other organism exert constant influence on an organism.

Science, Technology, and Society: We use technology to study life and we use the study of life to improve technology and society.

Other concepts in Biology

Emergent Properties: smaller parts work together and while they often retain their original properties, there are also new properties based on their synergy.

The Cell~ all organism’s basic structure Heritable Information~ DNA Reproduction: Organisms reproduce to

continue their species through time. This can be asexual or sexual in nature.

Natural Selection: As populations evolve sometimes the cause is due to differential reproductive success which leads to changes in the population to be genetically more like the successful organisms.

Chemical Context of Life

Matter (space & mass) Elements and compounds The atom (element) The molecule (compound) Atomic number (# of protons); mass

number (protons + neutrons): defines what element a substance is.

Isotopes (different # of neutrons); radioactive isotopes (nuclear decay)

Energy (ability to do work); energy levels (electron states of potential energy) Includes Chemical energy stored in bonds.

Radioactive isotopes uses

http://www.chem.duke.edu/~jds/cruise_chem/nuclear/uses.html

http://www.ausetute.com.au/nuclesum.html

CAT scan vs. PET scan

Chemical Bonding

Covalent Double covalent Nonpolar covalentPolar covalentIonicHydrogenDispersion forces

Covalent Bonding

Sharing pair of valence electrons

Number of electrons required to complete an atom’s valence shell determines how many bonds will form

Ex: Hydrogen & oxygen bonding in water; methane

Polar/nonpolar covalent bonds

Electronegativityattraction for electrons

Nonpolar covalent •electrons shared equally

•Ex: diatomic H and O

Ex: C and H in fats/lipids

Polar covalent•one atom more electronegative than the other (small charges)

•Ex: water

Ionic bonding

High electronegativity difference strips valence electrons away from another atom

Electron transfer creates ions (charged atoms)

Cation (positive ion); anion (negative ion)

Ex: Salts (sodium chloride)

Hydrogen bonds

Hydrogen atom covalently bonded to another atom with a different electronegativity.

The Hydrogen atom has a slight positive charge as the electron is spending more time with the other atom.

The positively charged hydrogen atom of one water is attracted to the negatively charged atom on the other molecule.

Hydrogen Bonds continued

Weak (1/100 of a covalent bond)

TemporaryPlentifulCommon in water

Dispersion forces

Weak interactions between molecules or parts of molecules that are brought about by localized change fluctuations.

Due to the fact that electrons are constantly in motion and at any given instant, ever-changing “hot spots” of negative or positive charge may develop.

Very weak and temporary

Acid/Base & pH

Dissociation of water into a hydrogen ion and a hydroxide ion

Acid: increases the hydrogen concentration of a solution

Base: reduces the hydrogen ion concentration of a solution

pH: “power of hydrogen” Buffers: substances that

minimize H+ and OH- concentrations (accepts or donates H+ ions)

pH scale: pH = -log10[H+]

Table 2.2 pH Scale

The hydrogen ion and hydroxyl ion concentrations are given in moles per liter at 25°C.

pH [H+] [OH-]0 (100) 1.0 0.00000000000001 (10-14)

1(10-1)

0.1 0.0000000000001 (10-13)

2(10-2)

0.01 0.000000000001 (10-12)

3(10-3)

0.001 0.00000000001 (10-11)

4(10-4)

0.0001 0.0000000001 (10-10)

5(10-5)

0.00001 0.000000001 (10-9)

6(10-6)

0.000001 0.00000001 (10-8)

7(10-7)

0.0000001 0.0000001 (10-7)

8(10-8)

0.00000001 0.000001 (10-6)

9(10-9)

0.000000001 0.00001 (10-5)

10(10-10)

0.0000000001 0.0001 (10-4)

11 (10-11) 0.00000000001 0.001 (10-3)

12(10-12)

0.000000000001 0.01 (10-2)

13 (10-13) 0.0000000000001 0.1 (10-1)14 (10-14) 0.00000000000001 1.0 (100)

pH = -log10[H+]

Buffering: Bicarbonate Ion

Bicarbonate Ion, HCO3- is the most common form in

blood and ocean. It can absorb or release an H+ based on environment.

This allows the system to buffered from pH change

Acid Precipitation

Carbon dioxide + water makes carbonic acid (pH = 5.6), so “clean” rain water has a pH of 5.6 as carbon dixoide is plentiful.

NOx and SOx are released in exhausts from cars and coal plants. They make sulfuric acid and nitric acid as they blow away which are much stronger acids.

Acid snow and fog become more concentrated as the water evaporates leaving behind a more concentrated acid.

Water Polar~ opposite ends, opposite charges Cohesion~ H+ bonds holding molecules

together Adhesion~ H+ bonds holding molecules to

another substance Surface tension~ measurement of the difficulty

to break or stretch the surface of a liquid Specific heat~ amount of heat absorbed or lost

to change temperature by 1oC Heat of vaporization~ quantity of heat required

to convert 1g from liquid to gas states Density……….

Density

Less dense as solid than liquid

Due to hydrogen bonding

Crystalline lattice keeps molecules at a distance