College Biology Chapter 2

Organic Chemistry

• Carbon containing molecules

• 4 electrons in outer shell - up to 4 covalent bonds

- usually doesn’t make ionic bonds

Hydrocarbons

• Made up of hydrogen and carbon

• Simplest version = methane

• Can make long chains

- branched or unbranched

- -ane = single bond, -ene = double bond, -yne = triple bond

• Aliphate - linear, or cyclic with all single bonds

• Aromatic - cyclic and alternating single/double bonds

Isomers

• Same formula, different function

- structural: different placements of covalent bonds

- geometric: different organization of functional groups

- enantiomers: “non-superimposable”

Functional Groups

• Groups of atoms that give a molecule specific properties

• Common and important in biological macromolecules

Acids

Lower number, add hydrogen ions to solution

Bases

Higher number, add hydroxide ions to solution

Water is:

Neutral (pH of 7)

Buffers

• Can absorb excess Hydrogen or Hydroxide ions, which maintain pH levels

- examples - antacids

Polar Molecule

• Uneven sharing of electrons leads to “slight” positive and negative charges

• Hydrophilic (polar):

- compare to hydrophobic (non-polar)

States of Matter

• Solid is actually less dense (ice floats, rare!)

• High specific heat makes is resistant to change

- amount of heat to raise 1 gram of 1 water degree celcius = “calorie”

- heat of vaporation = amount of energy needed to turn 1 gram of water to gas

Solvent

• Because its polar, polar/ionic things can dissolve and bind to the water molecules

• Dissociation:

- when atoms/ groups of atoms break off from molecules and form ions (NaCI, for example)

Other Properties

• Cohesion - sticking together

- surface tension - paper clips can float on water

• Adhesion - sticking to other things

- capillary action - changed surfaces attract water more than other water molecules

- example = roots

Ionic Bonds

Form between the positive and negative charges of ions.

Covalent Bonds

• “Sharing” electrons

• Polar

- shared electrons are unequally pulled towards the atoms

• Nonpolar

- made between atoms of the same element, or elements that share electrons equally

Hydrogen Bonds

• Slight positive/negative charges caused by polar covalent bonds

• Weak “bond” - technically a strong “intermolecular force”

Van Der Waals Interactions

• Based on electrondensity around an atom

• also “intermolecular forces”

Periodic Table

• Compounds are 2+ atoms from 2+ elements

• Molecules are 2+ atoms

Electrons

• Fill “orbitals”

- S,P,D,F=2,6,10,14

Octet Rule

• Atoms are more stable with 8 valence electrons

- 0=8

- naturally having 8 valence electrons=”noble gas”

• Valence electrons are the outermost shell

Ions

• Gaining electron=negative charge (anion)

• Losing electron=positive charge (cation)

Atom

• Has a nucleus

- hold protons and nuetrons

• Has “electron cloud”

- general areas where you’ll find electrons

• Mostly made up of empty space

Subatomic Particles

• Proton

- positive charge, 1 “amu”

• Nuetron

- nuetral charge. 1 “amu”

• Electron

- negative charge, “neligible mass”

Atomic Number

Number of protons

Mass Number

• Average number of protons + nuetrons for this element when found in nature

- isotopes have different # of nuetrons

Number of Nuetrons

Can be calculated by subtracting atomic number from mass number.