Chemistry

Chemistry Summary

Matter
1. Matter takes up space and has mass.
2. All living and nonliving matter is composed of 92 naturally-occurring basic elements.
3. Elements cannot be broken down to substances with different chemical or physical properties.
4. Six elements (C, H, N, O, P, S) make up 98% of living things.

Atomic Structure
1. Chemical and physical properties of atoms (e.g., mass) depend on the subatomic particles.
    a. Different atoms contain specific numbers of protons, neutrons, and electrons.
    b. Protons and neutrons are in nucleus of atoms; electrons move around nucleus.
    c. Protons are positively charged particles; neutrons have no charge; both have about 1 atomic mass
unit of weight.
    d. Electrons are negatively charged particles.
2. The atomic mass of an atom is about equal to the sum of its protons and neutrons.
3. All atoms of an element have the same number of protons, the atom’s atomic number.

The Periodic Table
1. The periodic table shows how various characteristics of atoms recur.
2. The table is arranged in order of atomic number, with periods in horizontal rows and groups in vertical columns.

Isotopes
1. Isotopes are atoms with the same number of protons but differ in number of neutrons; e.g., a carbon
atom has six protons but may have more or less than usual six neutrons.
2. A carbon with eight rather than six neutrons is unstable; it releases rays and subatomic particles and is
a radioactive isotope.
3. Low levels of radiation such as radioactive iodine or glucose allow researchers to trace the location and
activity of the atom in living tissues; therefore these isotopes are called “tracers.”
4. High levels of radiation can cause cancerous tissues and destroy cells; careful use of radiation in turn
can sterilize products and kill cancer cells.

Electrons and Energy
1. Electrons occupy an orbital at some level near or distant from the nucleus of the atom.
2. An orbital is a volume off space where an electron is most likely to be found; an orbital contains no more than two electrons.
3. The more distant the orbital, the more energy it takes to stay in the orbital.
4. When atoms absorb energy during photosynthesis, electrons are boosted to higher energy levels.
5. The innermost shell of an atom is complete with two electrons; all other shells are complete with eight electrons.

Compounds
1. When two or more different elements react or bond together, they form a compound (e.g., H2O).
2. A molecule is the smallest part of a compound that has the properties of the compound.
3. Electrons possess energy and bonds that exist between atoms in molecules contain energy.

Ionic Bonding
1. Ionic bonds form when electrons are transferred from one atom to another.
2. Losing or gaining electrons, atoms participating in ionic reactions fill outer shells, and are more stable.
3. Example: sodium with one less electron has positive charge; chlorine has extra electron that has
negative charge. Such charged particles are called ions.
4. Attraction of oppositely charged ions holds the two atoms together in an ionic bond.

Covalent Bonding
1. Covalent bonds result when two atoms share electrons so each atom has octet of electrons in the outer
shell.
2. Hydrogen can give up an electron to become a hydrogen ion (H+
) or share an electron with another
atom to complete its outer shell of two electrons.
3. Structural formulas represent shared atom as a line between two atoms; e.g., single covalent bond
(H–H), double covalent bond (O=O), and triple covalent bond (N = N).
4. Three dimensional shape of molecules is not represented by structural formulas but shape is critical in
understanding the biological action of molecules: action of insulin, HIV receptors, etc.

Nonpolar and Polar Covalent Bonds
1. In nonpolar covalent bonds, sharing of electrons is equal.
2. With polar covalent bonds, the sharing of electrons is unequal.
    a. In water molecule (H2O), sharing of electrons by oxygen and hydrogen is not equal; the oxygen
atom with more protons dominates the H2O association.
    b. Attraction of an atom for electrons in a covalent bond is called electronegativity; oxygen atom is
more electronegative than hydrogen atom.
    c. Oxygen in water molecule, more attracted to electron pair, assumes small negative charge.

Hydrogen Bonding
1. A hydrogen bond is weak attractive force between slightly positive hydrogen atom of one molecule
and slightly negative atom in another or the same molecule.
2. Many hydrogen bonds taken together are relatively strong.
3. Hydrogen bonds between complex molecules of cells help maintain structure and function.

Chemistry of Water
1. All living things are 70–90% water.
2. Because water is a polar molecule, water molecules are hydrogen bonded to each other.
3. With hydrogen bonding, water is liquid between 0º C and 100 C which is critical for life.
4. The temperature of liquid water rises and falls more slowly than that of most other liquids..
    a. Calorie is amount of heat energy required to raise temperature of one gram of water 1º C.
    b. Because water holds more heat, its temperature falls more slowly than other liquids; this protects
organisms from rapid temperature changes and helps them maintain normal temperatures.

Heat of Vaporization
    a. Hydrogen bonds between water molecules require a large amount of heat to break.
    b. This property moderates earth’s surface temperature; permits living systems to exist here.
    c. When animals sweat, evaporation of the sweat takes away body heat, thus cooling the animal.

Cohesion, Adhesion
    a. Cohesion allows water to flow freely without molecules separating, due to hydrogen bonding.
    b. Adhesion is ability to adhere to polar surfaces; water molecules have positive, negative poles.
    c. Water rises up tree from roots to leaves through small tubes.
    d. Adhesion of water to walls of vessels prevents water column from breaking apart.
    e. Cohesion allows evaporation from leaves to pull water column from roots.

Acids and Bases
1. Covalently bonded water molecules ionize; the atoms dissociate into ions.
2. When water ionizes or dissociates, it releases a small (107 moles/liter) but equal number of H+ and OH ions; thus, its pH is neutral.
3. Water dissociates into hydrogen and hydroxide ions: H – O –H → H+ + OH-
4. Acid molecules dissociate in water, releasing hydrogen ions (H+) ions: HCl → H+ + Cl-
5. Bases are molecules that take up hydrogen ions or release hydroxide ions. NaOH → Na+ + OH-
6. The pH scale indicates acidity and basicity (alkalinity) of a solution.
    a. Measure of free hydrogen ions as a negative logarithm of the H+ concentration (-log [H+]).
    b. pH values range from 0 (100 moles/liter; most acidic) to 14 (1014 moles/liter; most basic).
1) One mole of water has 107 moles/liter of hydrogen ions; therefore, has neutral pH of 7.
2) Acid is a substance with pH less than 7; base is a substance with pH greater than 7.
3) As logarithmic scale, each lower unit has 10 times the amount of hydrogen ions as next higher pH unit; as move up pH scale, each unit has 10 times the basicity of previous unit.
7. Buffers keep pH steady and within normal limits in living organisms.
    a. Buffers stabilize pH of a solution by taking up excess hydrogen (H+) or hydroxide (OH-) ions.
    b. Carbonic acid helps keep blood pH within normal limits: H2CO3 → H+ + HCO3-.

Organic Chemistry
1. Carbon has four electrons in outer shell; bonds with up to four other atoms.
2. Addition of an –OH (hydroxyl group) turns a carbon skeleton into an alcohol.
3. The ethanol alcohol is hydrophilic, it dissolves in water, because the –OH functional group is polar.
4. A hydrocarbon is hydrophobic except when it has an attached ionized functional group such as carboxyl (acid) (— COOH); then the molecule is hydrophilic.
5. Carboxyl groups ionize in solution and release hydrogen ions, being both polar and acidic.
6. Isomers are molecules with identical molecular formulas but differ in arrangement of their atoms.

Polymers
1. Polymers are the large macromolecules composed of three to millions of monomer subunits.
2. Polymers build by different bonding of different monomers; mechanism of joining and breaking these bonds is dehydration synthesis and hydrolysis.
3. Cellular enzymes carry out condensation synthesis and hydrolysis of polymers.
4. During dehydration synthesis, a water is removed and a bond is made (synthesis).
5. Hydrolysis reactions break down polymers in reverse of dehydration; a hydroxyl (— OH) group from water attaches to one monomer and hydrogen (— H) attaches to the other.

Monosaccharides
1. Monosaccharides are simple sugars with a carbon backbone of three to seven carbon atoms.
2. Glucose and fructose isomers have same formula (C6H12O6) but differ in structure.
3. Ribose and deoxyribose are five-carbon sugars (pentoses); the backbones of DNA.
4. Disaccharides contain two monosaccharides joined by dehydration synthesis.
5. Polysaccharides are polymers of monosaccharides.
6. Starch is straight chain of glucose molecules with few side branches.

Triglycerides
1. Lipids contain two molecular units: glycerol and fatty acids.
2. Glycerol is a water-soluble compound with three hydroxyl groups.
3. Triglycerides are glycerol joined to three fatty acids by dehydration synthesis.
4. A fatty acid is a long hydrocarbon chain with a carboxyl (acid) group at one end.
    a. Most fatty acids in cells contain 16 to 18 carbon atoms per molecule.
    b. Saturated fatty acids have no double bonds between their carbon atoms.
    c. Unsaturated fatty acids have double bonds in the carbon chain where there are less than two hydrogens per carbon.

Phospholipids
1. Phospholipids are like neutral fats except the third fatty acid is replaced by phosphate group or a group with both phosphate and nitrogen.
2. The phosphate group bonds to another organic group (R) and is the polar head; hydrocarbon chains become nonpolar tails.
3. Phospholipids arrange themselves in a double layer in water, so the polar heads face outward toward water molecules and nonpolar tails face toward each other away from water molecules.

Chemistry Review Notes

Physical Chemistry Equilibria Redox Thermodynamics Rate Equations Equilibrium Constant Electrode Potentials Acid-base Equilibria Periodicity Groups Halogens Period Qualitative Analysis Enthalpy Reaction Rates Equilibrium Transition metals Aqueous Ions Organic Chemistry Alkanes Halogenoalkanes Alkenes Alcohols Organic Analysis Isomerism Aldehydes and Ketones Carboxylic acids Arene Chemistry Amines Polymers Proteins and DNA Organic Synthesis NMR Chromatography

Atomic structure Substances Acids Redox Electronic Structure Bonding Hydrocarbons Alkanes Alkenes Alcohols Haloalkanes Organic Synthesis Analytical Techniques Acids & Bases Energy Redox Potentials Transition Metals Qualitative analysis Aromatic Compounds Carbonyls Carboxylic Acids and Esters Amines Amino Acids Polyesters and Polyamides Carbon Bonds Organic Synthesis Chromatography Spectroscopy

Free Practice Tests

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Chemical Reactions

Try clicking on one of the links below. This is an excellent reference list to study from, as it returns visual details and the mechanism of both the specified chemical reaction, as well as related reactions.

Chemical Reactions A - J	Chemical Reactions K - Z
Acetoacetic Ester Condensation	Kabachnik-Fields Reaction
Acetoacetic Ester Synthesis	Kindler Reaction
Acyloin Condensation	Knoevenagel Condensation
Alder-Ene Reaction	Kochi Reaction
Aldol Addition	Kolbe Electrolysis
Aldol Condensation	Kolbe Nitrile Synthesis
Appel Reaction	Kolbe-Schmitt Reaction
Arbuzov Reaction	Koser's Reagent
Arndt-Eistert Synthesis	Kulinkovich Reaction
Azide-Alkyne 1,3-Dipolar Cycloaddition	Kumada Coupling
Azo Coupling	Lawesson's Reagent
Baeyer-Villiger Oxidation	Leuckart Thiophenol Reaction
Baker-Venkataraman	Luche Reduction
Balz-Schiemann Reaction	Malonic Ester Synthesis
Bamford-Stevens Reaction	Mannich Reaction
Barton Decarboxylation	Markovnikov's Rule
Barton-McCombie Reaction	McMurry Reaction
Baylis-Hillman Reaction	Myers-Saito Cyclization
Beckmann Rearrangement	Michael Addition
Benzilic Acid Rearrangement	Michaelis-Arbuzov Reaction
Benzoin Condensation	Mitsunobu Reaction
Bergman Cyclization	Miyaura Borylation Reaction
Bestmann-Ohira Reagent	Mukaiyama Aldol Addition
Biginelli Reaction	Nazarov Cyclization
Birch Reduction	Nef Reaction
Bischler-Napieralski Reaction	Negishi Coupling
Blaise Reaction	Newman-Kwart Reaction
Blanc Reaction	Nitroaldol Reaction
Bohlmann-Rahtz Synthesis	Nozaki-Hiyama Coupling
Boronic Acid Mannich Reaction	Nucleophilic Substitution
Bouveault-Blanc Reduction	Ohira-Bestmann Reagent
Brook Rearrangement	Olefin Metathesis
Brown Hydroboration	Oppenauer Oxidation
Bucherer-Bergs Reaction	Overman Rearrangement
Buchwald-Hartwig Cross Coupling	Oxy-Cope Rearrangement
Cadiot-Chodkiewicz Coupling	Ozonolysis
Cannizzaro Oxidation Reduction	Paal-Knorr Furan Synthesis
CBS Reduction	Paal-Knorr Pyrrole Synthesis
Chan-Lam Coupling	Passerini Reaction
Claisen Condensation	Paterno-Büchi Reaction
Claisen Rearrangement	Pauson-Khand Reaction
Clemmensen Reduction	Pechmann Condensation
Click Chemistry	Petasis Reaction
Cope Elimination	Peterson Olefination
Cope Rearrangement	Pinacol Coupling Reaction
Conia-Ene Reaction	Pinacol Rearrangement
Corey-Bakshi-Shibata	Pinner Reaction
Corey-Chaykovsky Reaction	Prilezhaev Reaction
Corey-Fuchs Reaction	Prins Reaction
Corey-Kim Oxidation	Pschorr Reaction
Corey-Seebach Reaction	Ramberg-Bäcklund Reaction
Corey-Winter Olefin Synthesis	Reformatsky Reaction
Coumarin Synthesis	Ring Closing Metathesis
Criegee Ozonolysis	Ring Opening Metathesis
Cross Metathesis	Ritter Reaction
Curtius Reaction	Robinson Annulation
Dakin Reaction	Rosenmund Reduction
Darzens Condensation	Rosenmund-von Braun
Darzens Reaction	Rubottom Oxidation
Delépine Reaction	Sakurai Reaction
Dess-Martin Oxidation	Sandmeyer Reaction
Diazotisation	Saytzeff's Rule
Dieckmann Condensation	Schiemann Reaction
Diels-Alder Reaction	Schlosser Modification
1,3-Dipolar Cycloaddition	Schmidt Reaction
Directed ortho Metalation	Schotten-Baumann Reaction
Doebner Modification	Seebach Umpolung
Eglinton Reaction	Seyferth-Gilbert Homologation
Ene Reaction	Sarett Reagent
Enyne Metathesis	Shapiro Reaction
Epoxidation	Sharpless Aminohydroxylation
Eschweiler-Clarke Reaction	Sharpless Dihydroxylation
Ester Pyrolysis	Sharpless Epoxidation
Fischer Esterification	Shi Epoxidation
Favorskii Reaction	Simmons-Smith Reaction
Finkelstein Reaction	Sonogashira Coupling
Fischer Esterification	Staudinger Cycloaddition
Fischer Indole Synthesis	Staudinger Reaction
Fleming-Tamao Oxidation	Staudinger Reduction
Friedel-Crafts Acylation	Staudinger Synthesis
Friedel-Crafts Alkylation	Steglich Esterification
Friedlaender Synthesis	Stetter Reaction
Fries Rearrangement	Stille Coupling
Fukuyama Coupling	Strecker Synthesis
Fukuyama Reduction	Suzuki Coupling
Gabriel Synthesis	Swern Oxidation
Gewald Reaction	Tamao-Kumada Oxidation
Glaser Coupling	Tebbe Olefination
Griesbaum Coozonolysis	Tishchenko Reaction
Grignard Reaction	Tsuji-Trost Reaction
Grubbs Reaction	Trost Allylation
Haloform Reaction	Ugi Reaction
Hay Coupling	Ullmann Reaction
Heck Reaction	Upjohn Dihydroxylation
Hell-Volhard-Zelinsky Reaction	Van Leusen Oxazole Synthesis
Henry Reaction	Van Leusen Reaction
Hiyama Coupling	Nucleophilic Substitution
Hiyama-Denmark Coupling	Vilsmeier Reaction
Hofmann Elimination	Wacker-Tsuji Oxidation
Hofmann's Rule	Weinreb Ketone Synthesis
Horner-Wadsworth-Emmons Reaction	Wenker Synthesis
Hosomi-Sakurai Reaction	Willgerodt-Kindler Reaction
Huisgen Cycloaddition	Williamson Synthesis
Hunsdiecker Reaction	Wittig-Horner Reaction
Hydroboration	Wittig Reaction
Ireland-Claisen Rearrangement	Wohl-Ziegler Reaction
Itsuno-Corey Reduction	Wolff-Kishner Reduction
Iwanow Reaction (Reagent)	Wolff Rearrangement
Jacobsen Epoxidation	Woodward cis-Hydroxylation
Jacobsen-Katsuki Epoxidation	Woodward Reaction
Jocic Reaction	Wurtz Reaction
Jones Oxidation	Wurtz-Fittig Reaction
Julia-Lythgoe Olefination	Yamaguchi Esterification

Chemistry Compound Database - searchable by partial name, melting point, boiling point, index of refraction, molecular formula, UV absorption, mass spectral peaks, and chemical types.

Chemical Analysis

Oxidating Agents	Reducing Agents
Acetone	ADH
Ammonium cerium nitrate	Alcohol dehydrogenase
Ammonium peroxydisulfate	Boranes
1,4-Benzoquinone	Catecholborane
Benzoyl peroxide	Carrots
Bleach	Copper hydride
N-Bromosaccharin	Copper (low valent)
(E)-But-2-enenitrile	Chromium (low valent)
tert-Butyl nitrite	Daucus Carota
CAN	Decaborane
Cerium ammonium nitrate	DIBAL-H
3-Chloroperoxybenzoic acid	Diborane
Chromium trioxide	Dimethylsulfide borane
Collins Reagent	DMSB
Corey-Suggs Reagent	Fe
CMHP	Formaldehyde
Copper compounds	Formic acid
Crotononitrile	Hantzsch Ester
Cumene hydroperoxide	Hydrazine
DBDMH	Hydrogen
DDQ	Indium (low valent)
DEAD	Iron
Dess-Martin periodinane	Isopropanol
DIAD	LAH
DIH	Lithium
Dimethyl sulfoxide	Magnesium
Ferric chloride	Manganese
Ferric nitrate	3-Mercaptopropionic acid
Formic acid	3-MPA
Hydrogen peroxide	NBSH
Hypervalent bromine	Nickel
Hypervalent iodine	Nickel borohydride
HTIB	Niobium (low valent)
IBX	Phenylsilane
Iodine	Phosphorous acid
Iodobenzene dichloride	Pinacolborane
Iodosobenzene diacetate	PMHS
N-Iodosuccinimide	Potassium
Iodosylbenzene	Potassium iodide
2-Iodoxybenzoicacid	2-Propanol
Iron(III), (V) and (IV)	Sodium cyanoborohydride
Jones Reagent	Sodium dithionite
Koser's Reagent	Sodium hydrosulfite
Manganese compounds	Strontium
MCPBA	Titanium (low valent)
Methyltrioxorhenium	TMDSO
Molybdenum compounds	Tributyltin hydride
MTO	Trichlorosilane
Nitric Acid	Triethylphosphine
Oxygen	Trimethylphoshpine
Ozone	Triphenylphosphine
Sodium bromate	Triphenylphosphite
Water	Triethylsilane

        source:   Organic-Chemistry.org