2. A. Matter
Material that takes up space.
1. Elements
Pure chemical substances composed
of atoms.
Examples?
How many elements exist?
How many of these elements are
essential to life?
4. 2. Atom
The smallest “piece” of an element
that retains the characteristics of
that element.
Composed of 3 subatomic particles:
Protons
Neutrons
Electrons
7. Atomic number
# protons in nucleus of an atom (establishes
identity of the atom)
Since most atoms are electrically neutral,
atomic number indicates # of electrons
as well.
Mass number
# protons plus # neutrons in nucleus of an
atom
8. How can we determine the number of
neutrons in an atom?
# neutrons = atomic mass - atomic #
Determine # neutrons in a carbon atom
(atomic mass = 12; atomic # = 6).
# neutrons = 12 - 6 = 6
Do all carbon atoms have the same number of
protons?
Do all carbon atoms have the same number of
neutrons?
9. Isotopes
Atoms having the same number of
protons, but differing numbers of
neutrons.
Ex. Carbon isotopes
carbon 12 (12
C) → 6 neutrons
carbon 13 (13
C) → 7 neutrons
carbon 14 (14
C) → 8 neutrons
10. Periodic table information on carbon:
Atomic mass given in table is average
mass of all the element’s isotopes.
11. 3. Compound
A pure substance formed when atoms
of different elements bond.
The number of atoms of each element
is written as a subscript.
Examples:
CO2 carbon dioxide
H2O water
CH4 methane
C6H12O6 glucose
12. 4. Molecule
Smallest piece of a compound that
retains characteristics of that
compound.
The number of molecules is written as
a coefficient.
Examples:
4CO2 4 molecules of carbon dioxide
2C6H12O6 2 molecules of glucose
6O2 6 molecules of oxygen
13. 5. Chemical Bonds
Type of bond formed is determined
by the number of valence electrons
in the interacting atoms [octet rule].
a) Covalent bonds - form when atoms
share electron pairs.
can be nonpolar or polar
form molecules
16. Polar covalent bonds - electrons are
drawn more strongly to 1 atom’s nucleus
than the other.
Form when less electronegative atoms
bond with more highly electronegative
atoms.
Ex. water
17. b) Ionic bonds - form when oppositely
charged ions are attracted to each other.
stronger than covalent bonds
typically form salts
Ex. NaCl
18. c) Hydrogen bonds - form when opposite
charges on two molecules are attracted to
each other.
weakest type of bond*
Ex. DNA H2O
19. B. The Importance of Water
1. Properties
Cohesion - the attraction of water
molecules for each other.
Adhesion - the attraction of water
molecules for other compounds.
High heat capacity – takes a great deal of
heat to raise the temperature of water.
20. High heat of vaporization - a lot of heat is
required to evaporate water.
Exists as solid, liquid or gas - solid (ice) is less
dense than liquid.
2. Solutions
A solution is a mixture of one or more solutes
dissolved in a solvent.
If solvent is water, then it is an aqueous
solution.
Water is a strong solvent because it separates
charged atoms or molecules.
21. 3. Acids & Bases
Acids - substances that
add H+
to a solution.
Bases - substances
that remove H+
from
solution.
pH scale is measure of
acidity/alkalinity based
on H+
concentration.
22. C. Major Organic Molecules
Molecules that contain carbon in
combination with hydrogen.
1. Carbohydrates
contain C, H & O [# C ≅ # O]
function to store energy & provide
support
building blocks (monomers) are
monosaccharides
24. Disaccharides
simple sugars composed of 2
monosaccharides linked together by
dehydration synthesis.
Other common disaccharides: maltose
(seed sugar) & lactose (milk sugar).
26. 2. Lipids
contain C, H, O [# C >> # O]
do not dissolve in water
Triglycerides (fats)
composed of glycerol linked to 3 fatty acid
chains by dehydration synthesis.
function to cushion organs, as insulation &
in long-term energy storage (adipose tissue).
31. Proteins have a 3-dimensional shape
(conformation):
primary (1o
) structure - amino acid sequence
of polypeptide chain
secondary (2o
) structure - coiling & folding
produced by hydrogen bonds
tertiary (3o
) structure - shape created by
interactions between R groups
quarternary (4o
) structure - shape created
by interactions between two or more
polypeptides
34. DNA (deoxyribonucleic acid)
5-carbon sugar is
deoxyribose
nitrogenous bases are A,
G, C & T
double-stranded helix
held together by
hydrogen bonds
is the genetic material
35. RNA (ribonucleic acid)
5-carbon sugar is
ribose
nitrogenous bases are
A, G, C & U
single-stranded
enables information in
DNA to be expressed
Editor's Notes
109 elements have thus far been named (92 are naturally occurring; 17 are synthetic). 25 elements are essential to life.
Bulk elements - needed in large quantities. [H, C, N, O, P, S, Cl, K, Ca, etc] Trace elements - needed in small quantities. [ex. Iodine is required for the thymus gland to produce its hormones; Zinc is required to produce chlorophyll] A few elements ( arsenic , bromine & tin ) are toxic in large amounts, but may be vital in very small amounts.
All carbon atoms have the same number of protons, but they may have differing numbers of neutrons.
Since isotopes have differing numbers of neutrons, they have differing masses. Neutrons determine nuclear stability. Atoms with equal numbers of protons & neutrons are most stable. Unstable isotopes tend to break down into more stable forms. When they break down they release radioactive energy. Often one isotope of an element is very abundant & others are rare (99% of carbon isotopes have 6 neutrons).
NOTE: some molecules are “diatomic”, consisting of two atoms of the same element. Diatomic molecules include oxygen (O 2 ), nitrogen (N 2 ), hydrogen (H 2 ) and chlorine (Cl 2 ). These diatomic molecules are not considered to be compounds by definition.
Lines between the interacting atoms indicate a shared pair of electrons. The compounds seen here are commonly called hydrocarbons because they contain C & H only.
Electronegativity - tendency of an atom to attract electrons. Oxygen is highly electronegative. Hydrogen has a low electronegativity. The oxygen atom pulls hydrogen’s negatively charged electrons away from hydrogen’s nucleus an towards its own. Thus, the oxygen atom has a slight negative charge, while both hydrogens have a slight positive charge.
Sodium ion (Na + ) = sodium atom that has lost an electron. Chlorine ion (Cl - ) = chlorine atom that has gained an electron.
*Hydrogen bonds are individually weak, but a large number of them together will provide a great deal of strength, like the teeth of a zipper. Note: Hydrogen bonds are located between separate water molecules (here we see 4 hydrogen bonds); however, within each water molecule, the H atoms are bonded to the C by covalent bonds!
Movement of water from a plant’s roots to its highest leaves depends upon cohesion of water molecules & their adhesion to the plant’s water-conducting tubes. Water’s high heat capacity allows organisms to be exposed to extremes of temperature. Their body fluids heat or cool slowly.
Because water has a high heat of vaporization , the body cools off quickly as sweat evaporates from its surface. Because ice is less dense than water, it floats on the surface, forming a cap that retains heat in the water below.
Acids have pH values that range below 7.0, with a pH of zero being the most acidic. Bases have pH values that range above 7.0, with a pH of 14 being the most alkaline. Pure water has a pH of 7.0 and is considered to be neutral. * Each unit on the pH scale represents a 10 fold change in H + concentration. Thus, the H+ concentration of soap (pH=10.0) is 100 times greater than that of household ammonia (pH=12.0).
Common monosaccharides include: ribose (5 carbon sugar) deoxyribose (5 carbon sugar) glucose - blood sugar (6 carbon sugar) fructose - fruit sugar (6 carbon sugar) galactose (6 carbon sugar) Note: glucose fructose & galactose have the same molecular formula C 6 H 12 O 6 . Thus, they are isomers.
Glucose + Fructose Sucrose (table sugar) + water Because water is formed, sucrose will have 2 less hydrogen atoms & 1 less oxygen atom that glucose & fructose combined. Note: hydrolysis (splitting by adding water) breaks disaccharides apart yielding component monosacchrides. Glucose + Glucose Maltose Glucose + Galactose Lactose
Common polysaccharides include: cellulose - forms wood & parts of plant cell walls. starch - energy storage form in plants. glycogen - short term energy storage form in animals. cellulose, starch & glycogen are long chains of glucose units; differ in branching patterns chitin - forms the exoskeletons of arthropods & cell wall in many types of fungi. Cellulose is most common organic compound in nature. Chitin is second most common polysaccharide in nature.
Number of carbon atoms always much greater than number of oxygen atoms.
Note: many more C atoms than O atoms. Saturated fats - fatty acids are saturated with H atoms. There are no double-bonded carbons. Tend to be liquid at room temperature. Unsaturated - fatty acids are not completely saturated with H atoms (thus have 1 or more double bonded carbons). Double bonds cause kinks in the fatty acid tails. Tend to be solid at room temperature. Lipids in plants are less saturated than those in animals.
Cortisone is a steroid hormone. Cholesterol is a key component of cell membranes.
An amino acid contains a central carbon atom bonded to: a hydrogen atom a carboxyl group (COOH) an amino group (NH 2 ) an R group - differs for each of the 20 biologically important amino acids.
Conformation of protein is critical to its function.
Antibodies = Function in immunity. Hemoglobin = protein in RBCs that transports oxygen. Insulin & Glucagon = protein hormones that regulate levels of glucose in the bloodstream. Keratin = structural protein found in hair, nails, hooves. Fibrin & thrombin = proteins involved in blood clotting. Enzymes are protein catalysts (speed up chemical reactions without being altered in the process; thus they are reusable).
Each nucleotide is composed of: a 5 carbon sugar (ribose or deoxyribose) a phosphate group a nitrogenous base (guanine, cytosine, thymine, adenine or uracil).
