George Mason University
NOTE SET 1 – Spring 2010
Remember….
No make-up exams for any reason.
Cannot take exams in another class section.
Always abide by the Honor Code.
Be kind to each other!
Skeleton notes will be posted by most Fridays based on how much we cover in class. Remember that these are NOT FULL NOTE SETS. They are just the words on the slides. You are responsible for everything we talk about in class even if it’s not in the skeleton notes.
• Read the appropriate chapter BEFORE coming to the lecture.
• Come to all scheduled class meetings.
• Take notes on the slides.
• Study consistently and not just the night before the exam.
• Find a study buddy.
• Email your professor if you have questions!
CHAPTER 1
What is “LIFE?”
1. capacity to reproduce
-all organisms on earth use DNA or RNA
2. harness, store, and use energy
-create complex molecules from simpler ones
-use complex molecules to grow and develop
3. sense and respond to the environment
4. evolve over subsequent generations
-Populations evolve, not individuals!
5. die
What is “Science”? (The scientific method)
-Observation, hypothesis, and experimental based process (observe phenomenon)
• Develop hypothesis (a logical explanation)
• Devise tests of hypothesis (experiments)
– Minimize variables
– Use control groups
– Many observations (reduce sampling error)
– Eliminate or reduce bias
– Carry out tests and analyze results
Scientific Theory – A hypothesis that has been tested for its predictive power many times and has not
yet been found incorrect.
ENERGY AND LIFE
Nothing lives without energy
What is metabolism?
Ability of cells to use energy and small molecules to maintain self, grow and reproduce
-catabolism
-anabolism
What is energy? Light, heat, chemical nutrients that receptors can detect
Organisms often store energy in the form of carbohydrates and lipids (fats and oils)
Life shows amazing diversity yet all species on earth are interdependent on one another and share the same simple building blocks:
Nucleic acids
Proteins
Carbohydrates
Lipids
Levels of Biological Organization
What kinds of organisms are there?
-prokaryotes (simple, single celled, no nucleus)
-Archaea (extremophiles)
-Bacteria (most are bacteria)
-eukaryotes
-protists
-fungi
-plants
-animals
Levels of Biological Organization
Domain ( Kingdom ( Phylum ( Class ( Order ( Family ( Genus ( Species
Naming different forms of life:
First name is genus; Second name is species (Devised by Carl Linnaeus, 1707-1778)
e.g. Escherichia coli, Homo sapiens, Drosophila melanogaster
Levels of Biological Organization
Atoms ( molecules ( cells (life!) ( tissues ( organs ( organ system ( organism
Natural Selection and Evolution
• Individual variation and mutation
• Overproduction and competition
• Unequal reproductive success
• Biological change
Chapter 2 - Chemical Context of Life
• Hierarchy of Biological Order
• Atomic Carbon, hydrogen, oxygen, etc.
• Molecular DNA, proteins, etc.
• Organelle Nucleus, mitochondria, etc.
• Cellular Metabolism, cell signaling, etc.
• Tissue Smooth muscle, bone, etc.
• Organ Heart, brain, liver, etc.
• Organ Systems Circulatory, digestive, etc.
• Organism Mouse, human, maize, etc.
• Higher Levels Ecosystems, populations, etc.
Cell Biology
• Cells - Smallest unit which still retains characteristics attributed to life of organisms
• Cell Properties / Organism Properties
– Ability to reproduce
– Ability to grow
– Ability to use energy
– Ability to respond to the environment
• Organisms are composed of cells
– Some are unicellular; some multicellular
• Multicellular Organism
– Starts as a single cell (zygote)
– Reproduces by division (mitosis)
– Cells differentiate
• Have different forms and/or properties
• Cells may appear different morphologically, but they have many similarities
Basic Chemistry for Cell Biologists
• Matter - Composed of elements and compounds
• Elements required by living organisms
• Atomic structure
• Chemical bonding
• Chemical reactions
Matter
• Definition = Anything which takes up space and has mass
– Exists in many forms - Rock, wood, water, air, plastic, human, etc.
Element
• substance which cannot be broken down into other substances by chemical means
• 92 natural elements + ~20 man-made
• Names are abbreviated by 1 or 2 letters (symbols)
– H = Hydrogen He = Helium
– Na = Sodium (Natrium) Fe = Iron (Ferrum)
– (may be derived from Latin or Germanic name)
Compound
– Substance consisting of two or more elements combined in a fixed ratio
– Water: H2O = 2 Hydrogen : 1 Oxygen
– Table salt: NaCl = 1 Sodium : 1 Chlorine
• Sodium - very reactive (explosive) metal
• Chlorine - poisonous gas
– Combined = edible compound
– A compound has characteristics beyond those of its combined elements
Required Elements
• Number required by living organisms
– Only about 25 of 92 natural elements
– 96% of living matter composed of:
• Carbon, Hydrogen, Oxygen, Nitrogen
– Other 4% composed of:
• Phosphorus, Sulfur, Calcium, Potassium, others
Table 2.1 Naturally Occurring Elements in the Human Body
Trace Elements (~15-20)
• required only in very small quantities
• Iron (Fe) - hemoglobin, other proteins
• Iodine (I) - thyroid hormone
• e.g. Iodine - thyroid hormone - a daily intake of 0.15 mg of iodine is required for normal activity of the human thyroid gland.
Fig 2.3 Goiter due to Iodine deficiency
• "Iodine Deficiency Disorder (IDD) is the single greatest cause of preventable mental retardation. Severe deficiencies cause cretinism, stillbirth and miscarriage ….Even a moderate deficiency, especially in pregnant women and infants, lowers their intelligence by 10-15 IQ points …Today >1 billion people in the world suffer from iodine deficiency, and 38 million babies born every year are not protected from brain damage due to IDD. These … newborns come from families that are the least educated, most isolated and economically disadvantaged.”
-UNICEF Deputy Executive Director Kul Gautam
October 2007
Mich. State Med. Society and doctors proposed iodizing salt
• Worked with salt manufacturers (e.g. Morton Salt); added potassium iodide; 1924
• First example of a FUNCTIONAL FOOD
Basic Chemistry - Atomic Structure
• Atomic structure determines behavior of element
• Atoms combine by chemical bonding to form molecules
• Weak chemical bonds are important
• Shape of molecule is related to function
• Chemical reactions make and break bonds
Atom
• Smallest unit of matter that still retains properties of an element
• Composed of three types of particles
– Neutrons - no electrical charge
– Protons - positive electrical charge (+1)
– Electrons - negative electrical charge (-1)
• Neutrons and Protons
– Packed together in dense core referred to as the Atomic Nucleus
• Electrons
– Located in a cloud or shell around the nucleus
– Attracted to nucleus by positively charged protons
Fig. 2.4 Molecules of helium atom
Nucleus
• Small in comparison to entire atom (5/1000)
• Electrons orbit at a distance from nucleus
• Accounts for nearly all of the mass
– Protons + neutrons (electron mass is negligible)
• Neutrons & protons - about the same mass
– ~1 dalton (~1.7 x 10-24g) = 1 atomic mass unit
– “Dalton” after John Dalton, a British scientist
Basic Chemistry - Atomic Structure
• Atomic Number
– Number of protons in an atom
• Unique for each element
– Written as a subscript to left of chemical symbol - 2He, 8O, 16S, 26Fe
– Number of protons and electrons are generally the same - no net charge
• Mass Number
– Number of protons and neutrons
– Written as a superscript to left of chemical symbol - 4He, 16O, 32S, 55Fe
• Atomic Weight
– approximately the mass number
– proton and neutron both ≈ 1 dalton
• 4He has a mass number of 4
• estimated atomic weight of 4 daltons (4.003)
Number of Neutrons
# of protons = atomic number
# of protons + neutrons = mass number
Mass number - Atomic Number = # of neutrons
Number of Protons – cannot vary in atoms for a given element since it defines the element
• Number of Neutrons – CAN vary!
– Isotopes – atoms of an element which possess different numbers of neutrons
• Carbon - 12C - 6 protons + 6 neutrons
13C - 6 protons + 7 neutrons
14C - 6 protons + 8 neutrons
• Stability
• Unstable isotopes (radioactive) - break down (decay) and lose particles
• 12C and 13C are stable
• 14C is unstable (radioactive isotope)
• Radioactive Isotopes (AKA radioisotopes or radionuclides)
– Useful in scientific research and medicine
– Also pose a hazard to life
• Decay particles damage cellular molecules
– From nuclear reactor accidents or a “dirty” bomb
Fig 2. 5 Using Radioactive Isotopes
Fig 2.6 PET scan
Electrons
• Involved in chemical reactions between atoms
• Energy levels of electrons vary (potential)
– Discrete steps called electron shells
– Low energy near nucleus
– High energy further away
• Electron shells and Orbitals
– Once thought to be like planets orbiting the sun, but don’t know exact path
– Orbital is location of electron 90% of time
– Only 2 electrons per orbital at any time
– Can have more than one orbital in a shell
• Electron orbitals
– First energy shell - max of 2 electrons
• 1s orbital - 2 e-
– Second energy shell - max of 8 electrons
• 2s orbital - 2 e- ; Three 2p orbital - 6 e-
Fig. 2.9 Electron orbitals
Electron configuration
– Determines chemical properties of elements
– Depends upon the number of electrons in its outermost shell
– Referred to as valence shell and valence electrons
– Valence shell and valence electrons
– Completed shell is unreactive
– e.g. Helium, Neon, Argon
– Other atoms transfer or share electrons to complete valence shell
• Chemical bonds
Fig. 2. 8 – Electron configurations of the first 18 elements
Most atoms do not exist singly.
They have incomplete outermost shells and are unstable in that condition.
Seek stability by forming partnerships with other atoms.
- Two or more atoms linked together = molecule
Chemical Bonds
– Covalent bonds
• Nonpolar
• Polar
– Ionic bonds
– Hydrogen bonds
Covalent Bond
• Sharing of a pair of valence electrons by two atoms
– If unshared orbitals overlap, each atom can count both electrons toward its goal of filling the valence shell.
– Strongest of the chemical bonds
Fig 2.10 Formation of a covalent bond
When two hydrogen atoms approach, the electron of each atom is also attracted to the proton in the other nucleus.
The two electrons become shared in a covalent bond, forming an H2 molecule.
Covalent Chemical Bonds
• Sharing of valence electrons
• Strongest of the chemical bonds
– e.g. peptide bonds in proteins
– e.g. glycosidic bonds in sugars
– e.g. phosphodiester bonds in nucleic acids
Molecules
• Two or more atoms held together by covalent bonds constitute a molecule.
• The molecular formula indicates the number and types of atoms present in a single molecule.
– For molecular hydrogen: H2
– Structure abbreviation for molecular hydrogen: H-H
Covalent Bond - O2
• Oxygen has 6 electrons in valence shell
– Needs to add 2 e- to complete its valence shell.
– Two oxygen atoms can form a molecule by sharing two pairs of valence electrons.
– Forms a double covalent bond.
Valence
• Every atom has a characteristic total number of covalent bonds that it can form - an atom’s valence.
– The valence of hydrogen is 1. Hydrogen has 1 valence electron.
– The valence of oxygen is 2. Oxygen has 6 valence electrons.
– The valence of nitrogen is 3. Nitrogen has 5 valence electrons.
– The valence of carbon is 4. Carbon has 4 valence electrons.
– Phosphorus has a valence of 5, forming 3 single covalent bonds and 1 double bond.
Covalent Bonds
• Can form between atoms of the same element or atoms of different elements.
– While both types are molecules, the latter are also compounds.
– Water, H2O, is a compound in which two hydrogen atoms form single covalent bonds with an oxygen atom.
Fig 2.11c,d Covalent Bonds
Covalent Bonds
• Sharing of electrons between atoms
– Equal - nonpolar covalent bond
– Unequal - polar covalent bond
• Some atoms have a greater attraction for an electron
– Electronegativity is a measure of the degree of attraction
• Electronegativity is a measure of the degree of attraction; the > the electronegativity value, the more it attracts electrons to itself
• Values are calculated; Pauling Scale
Polar Covalent Bonds
• Form between atoms with different electronegativity
– Water: Oxygen - high electronegativity
Hydrogen - low electronegativity
– partial negative charge near the strongly electronegative atom
– partial positive charge near the weakly electronegative atom.
Fig. 2.12 Polar covalent bonds in a water molecule
Basic Chemistry - Ionic Bonds
• Transfer of electrons between atoms
– Sodium and Chlorine
• Sodium has one electron in valence shell
• Chlorine has seven
• Transfer one electron from sodium to chlorine
• Sodium becomes positively charged - cation
• Chlorine becomes negatively charged - anion
• Both have complete valence shell
Fig 2.13 Electron Transfer and Ionic Bonding
Ionic Compounds
• Compounds formed by ionic bonds
– Formula indicates ratio of elements
• NaCl - 1 atom of sodium: 1 atom of chlorine
• MgCl2: 1 atom of Mg : 2 atoms of Cl
• Strength of ionic bonds depends on environmental conditions
– dry conditions - salts are hard
– aqueous conditions - salts dissolve
Basic Chemistry - Hydrogen Bonds
• Weak type of bonding
• Hydrogen atom covalently bonded to one atom (strongly electronegative) is attracted to another atom (also strongly electronegative)
• The partially positive charged hydrogen atom in a covalent bond is attracted to negatively charged (partial or full) molecules, atoms, or even regions of the same large molecule.
Function Relates to Shape
• Shape of a molecule
– determined by the arrangement of electron orbitals that are shared by the atoms involved in a bond.
– A molecule with two atoms is always linear.
– However, a molecule with more than two atoms has a more complex shape.
– For atoms with electrons in both s and p orbitals, the formation of a covalent bond leads to hybridization of the orbitals to form four new orbitals in a tetrahedron shape.
• In a water molecule the hybrid orbitals that oxygen shares with hydrogen atoms are spread in a V shape.
• A methane molecule (CH4)
– all four hybrid orbitals are shared
– hydrogen nuclei are at the corners of the tetrahedron.
Fig 2.17 A molecular mimic
• Molecules with similar shapes can interact in similar ways.
• Binding to the receptors produces euphoria and relieves pain.
Chemical Reactions
• Chemical bonds are broken and reformed, leading to new arrangements of atoms.
• Making and breaking of chemical bonds
• All of the atoms in the reactants must be accounted for in the products. The reactions must be “balanced”.
Basic Chemistry - Chemical Reactions
Reactants ( Products
In Photosynthesis (summarized reaction):
6 CO2 + 6 H2O ( C6H12O6 + 6 O2
Some reactions can go forward or reverse
Reactants Products e.g. 3 H2 + N2 3 NH3
Reversible reactions indicated by the opposite headed arrows
Chemical Equilibrium
• Rate of formation of products is the same as the rate of breakdown of products
– At equilibrium
• products and reactants are continually being formed, but there is no net change in the concentrations of reactants and products.
– At equilibrium
• the concentrations of reactants and products are typically not equal, but their concentrations have stabilized.
Chapter 3 - Water and the Fitness of the Environment
Water - Polar molecule; Hydrogen bonding
Water - Has Polarity
– Hydrogen Bonding - between water molecules
– gives water its special properties
Cohesion
Adhesion
Surface Tension
Cohesion
Binding together of like molecules by H bonds
High in water
H-bonds constantly breaking and reforming
most water molecules are bonded to neighboring molecules at any instant
contributes to water transport in plants (See Fig 3.3)
Adhesion
Clinging of one substance to another
also involves H-bonds
also contributes to water transport in plants
water adheres to molecules of the walls of the xylem vessels in plant stems (trunks) helps counter the effects of gravity
Surface Tension
Measure of how difficult it is to stretch or break the surface of a liquid
Higher in water than other liquids
Ordered arrangement of water molecules at air-water interface
Fig 3.6 Hydrogen bonds in ice
Are more “ordered” than in liquid water, making ice less dense
Insulates the water it floats on
Water - Solvent of Life
Solvent - dissolving agent (water)
Solute - substance that is dissolved (sugar)
Aqueous solution - water is the solvent
Solution = solvent + solute (e.g. sugar in water)
What happens when a solute dissolves in a solvent?
– δ- (O) and δ+ (H) charges on water molecule have affinity for positive and negative parts of solute
• Other compounds besides ionic ones can also be solubilized
• Most will have regions that are polar - δ+ and δ- regions (like water)
– May have some ionic regions, too, Proteins, DNA, etc.
Two categories of substances:
– Hydrophilic; Hydrophobic
Hydrophilic - water-loving; compound with an affinity for water; polar or ionic compounds; can form H-bonds with water
Hydrophobic - water-fearing; compounds which lack an affinity for water; not miscible in water; non-polar compounds; lipids of the cell membrane; no H-bonding with water
Fig 3.8 A water-soluble protein
Polar and ionic groups are on the exterior of the protein and interact with water molecules.
Hydrophobic groups are on the interior of the protein (protected from water molecules).
Solutes are dissolved in solvent (water)
How are concentrations of solutes calculated?
Can’t weigh individual molecules, so must use units called MOLES
Solute Concentration - Moles
– Mole (mol) - Number of molecules in a mole is constant; always 6.022 x 1023 molecules / mol
– Avogadro’s number = 6.022 x 1023 molecules / mol
– (similar to concept of a “dozen”)
How to calculate the “weight” (mass) of a Mole (mol) of substance:
– It’s the molecular weight of a substance, but in gram units instead of amu.
– One mole of a substance “weighs” the sum of the atomic masses of each atom (molecular weight) of that substance in grams; e.g…..
– H2O 2(1) + 16 = 18 amu, so 18 g/mol
– sucrose C12H22O11 = 342 amu
12(12) + 22(1) + 11(16) = so 342 g/mol
Steps to Calculate amount of substance in 1 mole
• Examine molecular formula
• E.g. C6H12O6 (glucose)
• Add atomic masses for every atom present
• Carbon atoms - 6 X 12 amu = 72 amu
• Hydrogen atoms – 12 X 1 amu = 12 amu
• Oxygen atoms - 6 X 16 amu = 96 amu
180 amu for one molecule of glucose
Molarity
M stands for molarity
M = 1 mole/per 1 liter of solution present
A 1M solution means that you have 1 mole of the substance for every liter of solution present.
How many grams of glucose are needed to make a 1M solution?
How many grams for a 2M solution of glucose?
How many molecules of glucose in 1L of 1M?
To determine the amount of chemical substance that contains 6.02 X 1023 molecules of that substance, figure out the molecular weight.
Change the units on molecular weight from amu (daltons) to grams.
If you take 180 GRAMS of glucose, you will have 6.02 X 1023 molecules of glucose (aka one MOLE of glucose).
Substance Mol. Wt. Amt # molecules
A 10 g/mol 10g 6.022 x 1023
B 100 g/mol 100g 6.022 x 1023
Dissociation of Water
• Hydrogen atom shared between two water molecules may shift from one molecule to the other; end up with one hydronium ion and one hydroxide ion
Acids and Bases / pH
• In pure H2O, [H+] and [OH-] = 10-7 M
• How can an imbalance of [H+] and [OH-] occur?
• Acid - substance that increases the H+ concentration in solution
• Base - substance that reduces the H+ concentration in solution
• Product of [H+] and [OH-] = 10-14 M2 (Holds true for any aqueous solution)
[10-7 M] [10-7 M] = 10-14 M2 for pure water
• If [H+] increases, [OH-] decreases
10-7 10-7 = 10-14
10-5 10-9 = 10-14
– [H+] can vary greatly ( by a factor of 100 trillion (1014)
– pH scale used to express more conveniently; ranges from 0 to 14
• pH = - log [H+] ; For neutral solution: pH = - log [10-7] = -(-7) = 7
Calculating M of H + and OH -
pH 8
Alkaline (basic)
10-8 moles/L of H+ = 0.00000001 m/L
What is the concentration of OH - ?
14 – 8 = 6 [OH - ] = 10 -6 moles/L
10. -6 moles/L of OH - = 0.000001 m/L
Acidic solution pH < 7
• Basic solution pH > 7
– Each pH unit represents a 10-fold difference in [H+]
Fig 3.9 The pH of some common aqueous solutions
How can pH changes be minimized?
Buffers - Minimize changes in [H+] in solution; Present in all biological fluids
• Human blood maintained at pH 7.4
How do buffers work?
– accept H+ ions from the solution when in excess
– donate H+ ions to the solution when depleted
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