Big Chem Outline:
U01: Scientific Skills
U02: Matter
U03: Atomic Structure
U04: Periodic Table
U05: Bonding
U06: Reactions
U07: Energy and phases

Unit 3: Atomic Structure

* Protons are positively charged (+).

* Neutrons have no charge.

* Electrons are small and are negatively charged (-).

* Protons & neutrons are in an atom’s nucleus (nucleons).

* Electrons are found in “clouds” (orbitals) around an atom’s nucleus.

* The mass number is equal to an atom’s number of protons and neutrons added together.

* The atomic number is equal to the number of protons in the nucleus of an atom.

* The number of neutrons = mass number – atomic number.

* Isotopes are atoms with equal numbers of protons, but differ in their neutron numbers.

* Cations are positive (+) ions and form when a neutral atom loses electrons. They are smaller than their parent atom.

* Anions are negative ions and form when a neutral atom gains electrons. They are larger than their parent atom.

* Ernest Rutherford’s gold foil experiment showed that an atom is mostly empty space with a small, dense, positively-charged nucleus.

* J.J. Thompson discovered the electron and developed the “plum-pudding” model of the atom.

* Dalton’s model of the atom was a solid sphere of matter that was uniform throughout.

* The Bohr Model of the atom placed electrons in “planet-like” orbits around the nucleus of an atom.

* The current, wave-mechanical model of the atom has electrons in “clouds” (orbitals) around the nucleus.

* Electrons emit energy as light when they jump from higher energy levels back down to lower (ground state) energy levels. Bright line spectra are produced.

* Isotopes are written in a number of ways: C-14 is also Carbon-14, and is also like this:

* The distribution of electrons in an atom is its electron configuration.

* Electron configurations are written in the bottom center of an element’s box on the periodic table in your reference tables. Example: Calcium has an electron configuration of 2-8-8-2. This means it has 2 electrons in the first principal energy level, 8 in the second principal energy level, 8 electrons in the third principal energy level, and 2 electrons in the third principal energy level.

* Orbital notation is a way of drawing the electron configuration of an atom. Fluorine's orbital notation looks like this:

Each arrow represents an electron spinning on one direction. The adjecnt arrow represents another electron spinning in the opposite direction. Each spinning electron makes a magnetic field that's opposite the other's, and together they lock together like magnets.

* Polyatomic ions (Table E) are groups of atoms with an overall charge.

* Draw one dot for each valence electron when drawing an element’s or ion’s Lewis diagram.

* The kernel of an atom includes everything in an atom except the atom’s valence electrons.

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Unit 4: Periodic Table

* The Periodic Law states that the properties of elements are periodic functions of their atomic numbers.

* Periods are horizontal rows on the Periodic Table.

* Groups are vertical columns on the Periodic Table.

* Metals are found left of the “staircase” on the Periodic Table, nonmetals are above it, and metalloids border it.

*General properties of metals: malleable, ductile, lustrous, good conductors of heat & electricity, low ionization energy low electronegativity, tend to form positive ions.

*General properties of nonmetals: brittle when solid, are mostly gases at STP, dull, good insulators, high ionization energy and electronegativity, tend to form negative ions.* Noble gases (Group 18) are inert and stable due to the fact that their valence level of electrons is completely filled.

* Ionization energy increases as you go up and to the right on the Periodic Table.

* Atomic radii decrease left to right across a period due to increasing nuclear charge.

* Atomic radii increase as you go down a group.

* Electronegativity is a measure of an element’s attraction for electrons.

* Electronegativity increases as you go up and to the right on the Periodic Table.

* The elements in Group 1 are the alkali metals.

* The elements in Group 2 are the alkaline earth metals.

* The elements in Group 17 are the halogens.

* The elements in Group 18 are the noble gases.

* With the exception of groups 3 - 12, the last digit of an element’s group number is equal to its number of valence electrons.

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Unit 5: Bonding

* Binary compounds are substances made up of only two kinds of atoms.

*Chemical formulas are written so that the charges of cations and anions neutralize one another. Aka, the crossover method. All oxidation states must add to zero. Example: ammonium carbonate:

* When naming binary ionic compounds, write the name of the positive ion (cation) first, followed by the name of the negative ion (anion) with the name ending in “-ide.” Example: potassium chloride, magnesium sulfide.

* When naming compounds containing polyatomic ions, keep the name of the polyatomic ion the same as it is written in Table E. ammonium chloride, ammonium nitrate.

* Metallic bonds can be thought of as a crystalline lattice of kernels surrounded by a “sea” of mobile valence electrons.

* Atoms are most stable when they have 8 valence electrons (an octet) and tend to form ions to obtain such a configuration of electrons.

* Covalent bonds form when two atoms share a pair of electrons.

* Ionic bonds form when one atom transfers an electron to another atom when forming a bond with it.

* Nonpolar covalent bonds form when two atoms of the same element bond together.

* Polar covalent bonds form when the electronegativity difference between two bonding atoms is between 0.4 and 1.7.

* Ionic bonds form when the electronegativity difference between two bonding atoms is greater than 1.7.

* Substances containing mostly covalent bonds are called molecular substances.

* Substances containing mostly ionic bonds are called ionic compounds.

* Ionic substances tend to be: hard, high melting and boiling points, conduct electricity when molten or when they are in aqueous solutions.

* Covalent (Molecular) substances tend to be: soft, low melting and boiling points, do not conduct electricity (are good insulators).

* Hydrogen bonds form when hydrogen bonds to the elements N, O, or F and gives the compound unusually high melting and boiling points.

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Unit 6: Reactions

* The seven diatomics: they only exist in pairs of themselves at STP. “BrINClHOF” (Br2, I2, N2, CL2, H2, O2, F2).

* Coefficients are written in front of the formulas of reactants and products in chemical equations. They give us the ratios of reactants and products in a balanced chemical equation.

* Reactants are on the left side of the reaction arrow and products are on the right.

*Only coefficients can be changed when balancing chemical equations!

* Synthesis reactions occur when two or more reactants combine to form a single product: A + B -> AB

* Decomposition reactions occur when a single reactant forms two or more products: AB -> A + B

* Single replacement reactions occur when one element replaces another element in a compound: A + BC -> AC + B

* Double replacement reactions occur when two compounds react to form two new compounds: AB + CD -> AD + CB

* Law of Conservation of Mass - the masses of the reactants in a chemical equation is always equal to the masses of the products. Its why equations are balanced.

* The gram formula mass of a substance is the sum of the atomic masses of all of the atoms in it.

* Know how to calculate the percentage composition of a compound using the following formula. From the reference tables, Table T:

* Avogadro’s number is 6.02 E 23 and is the number of particles in 1 mole of a substance.

*Energy is released when a chemical bond forms. The more energy that is released, the more stable the bond is.

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Unit 7: Energy, Phses of Matter

* Endothermic reactions absorb heat. The energy value is on the left side of the reaction arrow in a forward reaction.

* Exothermic reactions release energy and the energy is a product in the reaction.

* The particles in a solid are rigidly held together.

* Solids have a definite shape and volume.

* Liquids have closely-spaced particles that easily slide past one another.

* Liquids have no definite shape, but have a definite volume.

* Gases have widely-spaced particles that are in random motion.

* Gases are easily compressed and have no definite shape or volume.

* Be able to read and interpret the heating/cooling curve as pictured below:

* Substances that sublime turn from a solid directly into a gas.

* Degrees Kelvin = °C + 273

*Be able to calculate heat absorbed/released by substance using: q = m * c * (final temp - initial temp)
q = heat absorbed or released (Joules), m = mass of substance in grams, c = specific heat (4.18 for water) , temperature must be in degrees Celsius

* The heat absorbed or released when 1 gram of a substance changes between the solid and liquid phases is the substance’s heat of fusion. (334 J/g for water, on reference table).

* The heat absorbed or released when 1 gram of a substance changes between the liquid and gaseous phases is the substance’s heat of vaporization. (2260 J/g for water, on reference table).

* Liquids boil when their vapor pressure is equal to the atmospheric pressure.

* The normal boiling point of a substance is the temperature at which it boils at 1 atm of pressure. (Take note of Table H)

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Unit 8: Gases

* “STP” means “Standard Temperature and Pressure.” (273 Kelvin & 1 atm).

* As the pressure on a gas increases, the volume decreases proportionally.

* As the pressure on a gas increases, temperature increases.

* As the temperature of a gas increases, volume increases.

* Always use Kelvins for temperature when using the combined gas law: (P1V1)/T1 = (P2V2)/T2

* Real gas particles have volume and are attracted to one another, and thus do not always behave like ideal gases.

* Real gases behave more like ideal gases at low pressures and high temperatures.

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Unit 9: Kinetics and Equilibrium

* Covalently bonded substances tend to react more slowly than ionic compounds.

* Increasing the concentration of reactants will increase reaction rate.

* Increasing the surface areas of the reactants will increase reaction rate.

* Increasing the pressure on gases increases reaction rate.

* Catalysts speed up reactions by lowering their activation energies. They are not changed themselves and can be reused many times over.

* Increasing temperature increases reaction rate.

* DH is (+) for endothermic reactions and is (-) for exothermic reactions.

* The rates of the forward and reverse reactions are equal at equilibrium.

* Adding any reactant or product to a system at equilibrium will shift the equilibrium away from the added substance.

* Removing any reactant or product from a system at equilibrium will shift the equilibrium point toward that removed substance.

* Be able to recognize and read potential energy diagrams:

* An increase in temperature shifts an equilibrium system in the endothermic direction.

* A decrease in temperature shifts an equilibrium system in the exothermic direction.

* Increasing the pressure on a gaseous equilibrium will shift the equilibrium point toward the side with fewer moles of gas.

* Decreasing the pressure on a gaseous equilibrium will shift the equilibrium point toward the side with more moles of gas.

* Catalysts have no effect on an equilibrium. It just establishes itself quicker.

* Enthalpy (H) is the heat energy gained or lost in a reaction.

* Entropy (S) is high in a highly unorganized system, such as a gas, a messy room, etc.

* For the hypothetical reaction: wA + xB ->yC + zD, Keq is expressed by:

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Unit 10: Solutions

* Solutions are the best examples of homogeneous mixtures. (Air, salt water, etc.)

* A solute is the substance being dissolved, while the solvent is the substance that dissolves the solute. (Water is the solvent in Kool-Aid, while sugar is the solute.)

* Use Table F to predict the solubilites of compounds.

* Remember substances tend to be soluble in solvents with similar bond types. Polar solutes dissolve in polar solvents. Nonpolar solutes dissolve in nonpolar solvents. “Like dissolves like”

* As temperature increases, solubility increases for most solids.

* At low temperatures and high pressures solubility increases for most gases.

* Use Table G to determine whether a solution is saturated, unsaturated, or supersaturated. Here's how to interpret it:

* Molarity is a way to measure the concentration of a solution. Molarity is equal to the number of moles of solute divided by the number of liters of solution. The formula is on the back of the reference tables.

* Percent by mass = mass of the part / mass of the whole x 100%

* Parts per million (ppm) = grams of solute / grams of solution x 1,000,000

* Solutes raise the boiling points and lower the melting points of solvents.

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Unit 11: Acids and Bases

* Acids and bases are both good electrolytes. Their solutions conduct electricity well.

* Weak acids taste sour.

* Weak bases taste bitter.

* Acids and bases turn indicators different colors. They’re listed on Table M.

* Acids have a pH < 7.

* Bases have a pH > 7.

* Tables K & L list names and formulas of common acids and bases asked about on the Regents.

* The metals above H2 on Table J will react with acids to make H2 gas bubbles.

* Arrhenius says: “Acids give off H+ or H3O+ ions in solution.” “Bases give off OH- ions in solution.”

* Brønsted says: “Acids donate protons.” “Bases accept protons.”

* Acids and bases react in neutralization reactions to make water and a salt.

* Titrations are controlled neutralization reactions used to find the concentration of an acid or base sample. Note the formula for it on Table T.

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Unit 12: Oxidation and Reduction

* Oxidation is the loss of electrons by an atom or ion. The oxidation number increases as a result. The electrons are on the right side of the reaction arrow.

* Reduction is the gain of electrons by an atom or ion. The oxidation number decreases (is reduced!) as a result. The electrons are on the left side of the reaction arrow.

* Redox reactions always involve the exchange of electrons.

* Remember…. “LEO says GER!”, Lose electrons oxidation, gain electrons reduction.

* Identify redox reactions by seeking an uncombined element on one side of a reaction that is in a compound on the other side: Zn + 2HCl -> ZnCl2 + H2, Uncombined Zn is combined with Cl.

* Oxidizing agents are what get reduced in a redox reaction. Reducing agents are what get oxidized in a redox reaction.

* Electrochemical cells produce electricity with a spontaneous redox reaction.

* The left electrode is usually the site of oxidation in an electrochemical cell diagram.

* Memorize this saying… “I have AN OX and a RED CAT.”

* In electrochemical cells, the ANode gets OXidized and REDuction occurs at the CAThode.

* Electrolytic cells use an applied electrical current to force a nonspontaneous redox reaction to take place.

* Electrolytic cells are usually used for metal plating of objects.

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Unit 13: Organic Chemistry

* All organic compounds contain the element carbon.

* Carbon ALWAYS makes four bonds in molecules.

* Saturated hydrocarbons have all single bonds within them (alkanes).

* Unsaturated hydrocarbons have double or triple bonds in them (alkenes & alkynes).

* Hydrocarbons contain ONLY the elements hydrogen and carbon.

* The homologous series of hydrocarbons’ formulas are on Reference Table Q.

* The functional groups on organic molecules are listed on Reference Table R.

* Structural isomers of organic compounds have different structural formulas but the same molecular formula.

* Number the parent carbon chain in an organic molecule from the end closest to the alkyl group(s).

* Combustion reactions occur when a hydrocarbon reacts with oxygen to make CO2 and H2O.

* Organic substitution reactions occur when an alkane and a halogen (Group 17) reacts so that one or more hydrogen atoms on the alkane are replaced with oxygen.

* Organic addition reactions occur when an alkene or alkyne combine with a halogen to make one product (halide).

* Esterification occurs when an organic acid and an alcohol react to make water and an ester.

* Saponification occurs when an ester reacts with a base to make alcohol and a soap.

* Fermentation reactions occur when yeast catalyze a sugar (C6H12O6) to make carbon dioxide and ethanol.

* Polymers are long chains of repeating units called monomers.

* Polymers form by polymerization reactions.

* Addition polymerization occurs when unsaturated monomers join in a long polymer chain.

* Condensation polymerization occurs when monomers join to form a polymer by removing water. Water is a product!

* Natural polymers include starch, cellulose, and proteins. Synthetic polymers include plastics such as nylon, rayon, and polyester.

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Unit 14: Nuclear Chemistry

* Unstable atoms that are radioactive are called radioisotopes. (Table N)

* Radioisotopes can decay by giving off any of the particles/emanations listed in Table J.

* Alpha particles (see Table J) are positively charged (+).

* Beta particles (see Table J) are negatively charged (-).

* The half life of a radioisotope is the length of time it takes for one half of the atoms in a sample to radioactively decay. (Table N)

* C-14 is used to determine the ages of organic material up to 23,000 years old.

* U-238 is used to determine the ages of rocks.

* I-131 is used to treat thyroid disorders.

* Co-60 is used to treat cancer tumors.

* Radiation can be used to kill bacteria on foods to slow the spoilage process.

* Disposal of radioactive waste is a problem associated with nuclear reactors.

* The sum of the mass numbers and atomic numbers must be equal on both sides of the reaction arrow for nuclear equations.

* Fission reactions split heavy nuclei into smaller ones.

* Fusion reactions occur when light nuclei combine to form a heavy nucleus and a lot of energy.

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4.12.08