Lecture: Basic Chemistry
I. Matter and Energy
A. Matter - fundamental building blocks of nature
1. elements - basic units of matter
B. Energy - capacity to do work (put matter into motion
1. potential energy - energy stored in a structure
a. water stored in a lake uphill
b. chemical bonds of glucose molecule
2. kinetic energy - energy in an object in motion
a. water in a stream - allows mill to grind corn
b. broken glucose bonds -> ATP -> muscles work
3. Forms of Energy
a. chemical energy - energy in chemical bonds
I. ATP (adenosine triphosphate) - stores energy
b. electrical energy - energy of separated charges
I. battery - + pole and - pole separate charge
ii. nervous impulse run just like a battery
c. mechanical energy - energy of matter in motion
I. bowling ball transfers energy to move pins
ii. muscle motion - ATP -> contraction of muscle
d. electromagnetic energy - energy traveling in waves (light, X-rays, UV rays)
I. electromagnetic spectrum - visible light, UV light, radio waves, X-rays
C. First Law of Thermodynamics
1. "Energy can change from one form to another, but it can never be created or
destroyed" (Total Energy In = Total Energy Out)
examples:
Car Engine vs. Human Body
a. Car Engine - gasoline used to run motor to move car
Chemical Energy (gas) ---> motion (20%) + heat (79%) + sound (1%)
b. Human Body - food used to move body, digest, think, etc.
Chemical Energy (food/glucose) ---> physiology (80%) + heat (20%)
II. Organization of Matter (Atoms - Elements)
A. Atomic Particles
Mass
Charge
Characteristics
proton
1
+1
defines element
neutron
1
neutral
defines isotopes
electron
0
-1
determines element bonding properties
B. Organization of Periodic Table
1. # protons = atomic number (unique for each element)
2. # protons + # neutrons = atomic mass
3.
isotope - same element; different # neutrons
# protons
Carbon-12 (99%)
Carbon-13 (0.9%)
Carbon-14 (0.1%)
4.
+
# neutrons =
6
6
6
6
7
8
atomic mass
12
13
14
# electrons - dictates the NET CHARGE of an atom
H
H+
H-
# protons
# electrons
1
1
1
1
0
2
NET CHARGE
0
+1
-1
ion - any atom with a positive or negative charge
anion - atom with a NEGATIVE charge
cation - atom with a POSITIVE charge
III. Electron Shells, the Periodic Table, and Chemical Bonds
A. Electron Shells - electrons occupy "shells" as they orbit around the nucleus (2, 8, 8, ...)
B. The Periodic Table of Elements is organized by electron shells
H1
Li3 Be4
Na11 Mg12
He2
B5 C6 N7 O8 F9 Ne10 SHELL 2
Al13 Si14 P15 S16 Cl17 Ar18 SHELL 3
SHELL 1
8 e8 e-
2 e-
C. Chemical Bonds are formed so that each atom can have the outermost electron shell filled
1. Ionic Bond - one atom donates electron(s) to another
a. Example: Sodium Chloride (table salt) Na+Cl-
2. Covalent Bond - two atoms share one/more electrons
a. Example: Methane (CH4), Carbon Dioxide (CO2), and Ammonia (NH3)
a. Polar Molecule - electron sharing is unequal in the bonds
Example: Water (H2O)
b. Non-polar Molecule - electron sharing is almost equal
Example: Methane (C02)
IV. Elements other than C,H,O,N in Humans
Primary Elements (3% of all body weight)
Calcium
Ca
Bones,teeth, muscle and nerve action, blood clotting
Phosphorus
P
Bones and Teeth, DNA, RNA, ATP. Important in energy transfer
Trace Elements (Less than 1% of body weight altogether)
Potassium
K
Osmotic balance; cell voltage, muscle and nerve action
Sulfur
S
Component of proteins (cysteine) and other organic molecules
Sodium
Na
Osmotic balance; cell voltage, muscle and nerve action
Chlorine
Cl
Osmotic balance; cell voltage, muscle and nerve action
Magnesium
Mg
Co-factor for many enzymes
Iron
Fe
Hemoglogin and many enzymes
Copper
Cu
Co-factor of many enzymes
Zinc
Zn
Co-factor of many enzymes
Manganese
Mn
Co-factor of many enzymes
Cobalt
Co
Co-factor of many enzymes and vitamin B12
Chromium
Cr
Co-factor of many enzymes and potentiates Insulin
Selenium
Se
Required for normal liver function
Molybdenum Mo
Co-factor of many enzymes
Flourine
F
Tooth and bones
Tin
Sn
Promotes growth (unknown mechanism)
Silicon
Si
Growth, bone mineralization, connective tissue synthesis
Vanadium
V
Promotes growth and reproduction
V. Chemical Reactions
A. Patterns of Chemical Reactions
1. Chemical Equation - # of atoms of each element same for reactants and products
C6H12O6 + 6O2
---------->
6H2O + 6CO2
2. Synthesis - smaller molecules form larger molecule
A + B ----> AB (anabolic process)
amino acid 1 + amino acid 2 + ........
---->
peptide (protein)
sugar 1 + sugar 2 + sugar 3 + ..........
----->
polysaccharide (glycogen)
3. Decomposition - larger molecule broken down into smaller molecules
AB ----> A + B
glycogen ----> glucose + glucose + glucose + ..........
4. Displacement - one part is exchanged
AB + C -----> A + BC
glucose + adenosine-P-P-P (ATP) ------->
glucose-P + adenosine-P-P (ADP)
B. Exergonic vs. Endergonic Reactions
1. Exergonic - energy is released during the reaction
A + B ------> C + D +
ENERGY
glucose + oxygen ----> water + carbon dioxide +ENERGY (trapped by ATP)
2. Endergonic - energy required for reaction to proceed
A +
B
+
ENERGY ------>
C
amino acid 1 + amino acid 2 + ... + ENERGY ---> peptide (protein)
C. Chemical Equilibrium
1. Reversible Reactions
A + B --------->
AB
and
AB --------> A + B
2. Chemical Equilibrium
A + B
<=====>
AB
D. Rates of Chemical Reactions
1. size of reactants species (smaller means faster)
2. temperature (speeds up the particles)
3. concentration (more likely to come together)
4. catalysts (enzymes) - make reacting more convenient
VI. Acid- Base Chemistry and the pH Scale
A. Water normally exists in an equilibrium reaction with some dissociation
H2O
<======>
H+
+
-
OH
in a beaker of pure water, the ratio of H+ to H2O is about 1/10,000,000
pH = -log10[H+]
= -log10[10-7]
= -(-7)
=
7
pH = relative concentration of H+ in a solution of water
B. Acids - compounds which increase the concentration of H+ (pH = 1 to 6)
C. Bases - compounds which decrease the concentration of H+ (pH = 8 to 14)
D. Buffer - compound that prevents large changes in pH of a solution (pH “shock absorber”)