Unit 1 Atomic Structure
Table O
1. History: The progression of the atomic model from the solid sphere to the “plum pudding model” (discovery of electrons) to the “planetary model” (due to Gold foil experiment) to the excited electrons of the wave mechanical model.
2. Atoms – nucleons (protons + neutrons) and electrons
a. Proton – mass of 1 amu, charge +1, found in nucleus. # of protons = nuclear charge, atomic number
b. Neutron – mass of 1 amu, charge = 0, found in nucleus, varies in isotopes
c. Electron – mass of 0 amu, charge = -1, found outside nucleus in areas of high probability, 1/1836 amu
3. Atomic number - # of protons, identifies the element. Remember atomic number = # of protons or # of electrons in neutral atom
a. mass # = # of protons + # of neutrons (MAN)
4. nuclear charge - =# of protons
5. # of neutrons = mass number – atomic number (MAN)
6. isotopes – same # of protons but different # of neutrons, same atomic number but different mass number
7. Atomic mass – average of all the naturally occurring isotopes by mass and %. Closest to most abundant isotopes mass
8. Ground state – electrons are in lowest energy levels possible
9. Excited state – electrons are not in the lowest energy levels possible, not like the Periodic Table. Ground + heat or electricity à excited, after electron falls back to ground, energy is released as light (spectra is specific, identifies the element)
10. Lewis Dot diagrams – symbol + valence electrons only – maximum of 8
11. Maximum number of electrons in each energy level = 2-8-18-32
Unit 2 Periodic Table
Table S
1. Periodic Table is arranged by atomic number
2. Periods – horizontal rows = same energy level
3. Groups – verticals columns = same number of valence electrons
4. Trends:
a. periods = radius decreases, Ionization energy increases, electronegativity increases
b. groups = radius increases, Ionization energy decreases, electronegativity decreases
c. elements in groups have similar chemical properties
5. Electronegativity – attraction for electrons in a bond
6. Metals, metalloids and nonmetals
a. metals = left of staircase, most reactive = lower left (Cs), malleable, luster, conductors, ductile, positive ions, “sea of mobile electrons”
b. Nonmetals = right of staircase (except H) most reactive = top right (F), nonconductive, mostly gases.
c. metalloids = on staircase, both metal and nonmetal properties
7. Group names – 1 = alkali metals, 2 = alkaline earth metals, 3-11, transition elements, 17 = halogens, 18 = noble gases (monatomic or inert gases)
8. Diatomic elements – N2, O2, F2, Cl2, Br2, I2, H2, (BrINClHOF). Nonpolar bonds, always diatomic by themselves.
9. Liquids – Br2 and Hg, Nonmetal solids – I2, S, P, C
Unit 3 - Mathematics
Table T
1. Molar mass – gfm = from Periodic Table
2. Moles = given mass/gfm
3. equal volumes of gas means equal number of molecules under the same conditions (pressure and temperature)
4. coefficients – molecule/atom ratio or mole ratio; large number in front of formula
5. Subscripts – tells number of atoms in a compound, follows element
6. % by mass = part/whole X 100
7. Compounds and element: pure substances
a. Compounds can be broken down by chemical means
b. Elements cannot be broken down by chemical means
8. Homogeneous and heterogeneous mixtures
a. Heterogeneous mixtures are visibly different
b. Homogeneous mixtures can be separated by physical means (filtration, distillation, paper chromatography)
9. Empirical formula – subscripts of the smallest whole number ratio
10. Finding empirical formula –
a. change % to grams
b. use grams to find # of moles
c. divide all moles by smallest number
11. Reaction types – synthesis, decomposition, single replacement, double replacement or combustion
12. Balancing equations = law of conservation
13. Particle diagrams – shows relationship of particles in solids, liquids, gases, compounds, elements and mixtures
Unit 4 – Bonding
Tables T and E
1. Noble gases do not bond – because they have 8 valence electrons (full shell)
2. Ionic bonds – electrons are transferred from the metal to the nonmetal and ions are formed.
a. Metals give electrons and form (+) ions and nonmetals take electrons to form (-) ions.
b. Properties – high melting point, electrolytes when dissolved or melted due to mobile charged particles, solids at room temperature, dissolve in water.
c. Ionic character is greatest with largest difference in electronegativity
3. Ionic radius:
a. Positive ions – decrease in radius due to loss of electrons and shell
b. Negative ions – increase in radius due to gain of electrons and repelling of more electrons in a shell
4. Covalent bonds – electrons are shared between nonmetals, forms a molecule
a. polar – unequal sharing, difference in electronegativity = .5-1.6
b. nonpolar – equal sharing, difference in electronegativity = < .4
c. properties – poor conductors of electricity, low melting point, soft, does NOT dissolve in water
5. Metallic bonds – “sea of mobile electrons” in metals, allows for conduction of heat electricity, explains properties of metals
6. Dipoles – attraction between oppositely charged ends of polar molecules
a. H – bonds = stronger attraction than dipoles between polar molecules which have H bonded to F, O, or N (small, highly electronegative elements) which accounts for higher boiling points of compounds such as water (H2O)
b. non-dipole forces = Van Der Waals forces, weak forces between nonpolar molecules
c. molecule – ion attraction = attraction between polar and ionic ( salt dissolving in water, water molecules arranging themselves around charged particles)
d. intermolecular forces explain differences in physical properties (strong in solids and weak in gases)
7. The larger molecule, the stronger the force between nonpolar molecules, the higher the boiling point
8. Formula = ratio of atoms in a compound (drop and swap charges)
a. polyatomic ions = a group of atoms (covalently bonded) that act as 1 ion
b. Roman numeral = after a metal indicates its charge (stock system)
c. binary compounds = 2 elements, ending in ~ide
9. When a bond is formed, energy is released (more stable)
10. When a bond is broken, energy is absorbed (less stable)
Unit 5 – Matter and Energy
Tables A, B, F, G, H and T
1. Temperature = average Kinetic energy
2. Heat flows from hot to cold (high to low)
3. Heating and Cooling Curves
a. slope = temp increases, KE increases, one phase, Q= mCΔT
b. plateau = temp remains the same, PE increases, change of phase, Q=mHf or Q=mHv (Hf = melting/freezing, Hv = boiling/condensation)
4. Phase properties
a. gas = indefinite volume and shape
b. liquid = definite volume, indefinite shape
c. solid – definite volume and shape, regular geometric pattern, crystalline
5. Sublimation – solid to gas phase change (CO2 or I2)
6. Increasing vapor pressure increases the boiling point. Substance boils when vapor pressure equals atmospheric pressure. Liquids with lower intermolecular forces have lower boiling points.
7. Gases
a. PV/T = PV/T but make sure to change temp to Kelvin
b. # temp = # pressure, # pressure = i volume, # temp = # volume
c. Kinetic Molecular Theory = all gases have continuous random straight line motion, collisions are elastic, volume of gas particles is negligible, no forces of attraction
8. Ideal gases
a. most ideal conditions are high temp and low pressure ( hot beach)
b. H and He come closest to ideal
9. Solutions = homogeneous mixture of 2 or more elements/compounds physically combined
a. homogeneous – same throughout, usually aqueous
b. saturated = no more solute can dissolve in the solvent, on the line of Table G, at equilibrium
c. unsaturated = more solute can dissolve in the solvent, under the line of Table G
d. Supersaturated = more solute than normally possible is dissolved in the solute. Occurs only rarely
e. most soluble compounds at the top of Table G
f. # solubility of solids = #temp
g. # solubility of gases = i temp # pressure
10. Molarity à M = moles of solute/Liter of solution
11. ppm – parts per million = part/whole x 106
12. % comp – part/whole x 100
13. Colligative properties
a. Boiling point # when a solution is made, Freezing point i when a solution is made
b. Boiling point and freezing point are affected more by an increase in concentration and if the solute is ionic (separate into ions when dissolved) because the more particles there are, the more the boiling point goes up and the more the freezing point goes down.
14. When determining solubility: “Like dissolves like”
Unit 6 – Equilibrium and Kinetics
Table I
1. Kinetics
a. Exothermic = heat is released; heat is a product; heat or energy (Joules) on the right of the arrow; -ΔH; more favorable and the more stable the compound; feels hot to the touch
b. Endothermic = heat is absorbed; heat is a reactant; heat or energy (Joules) on the left of the arrow; +ΔH; feels cold to the touch
c. If the amount of reactants or products is halved or doubled, the same happens to the ΔH.
d. Heat of reaction ΔH = PEproducts - PEreactants
exothermic endothermic
| |
| |
| |
| |
| |
|______|______
2. Rate of reaction is affected by:
a. nature of reactants (ionic or dissolved or both)
b. concentration: # concentration # rate
c. temperature: # temperature # rate
d. surface area: # SA # rate
e. catalyst iactivation energy with alternate pathway #rate
3. Effective collision is needed (contact, correct orientation, speed) # EC # rate
4. Entropy = disorder. Increase in disorder is favorable (phase change SàLàG; making a solution, decomposition of compounds to elements)
5. Enthalpy = energy. Less energy in compounds is favorable (exothermic)
6. Spontaneous reactions will occur based on decreasing enthalpy (ΔH) AND increasing entropy (disorder)
7. Equilibrium = when the forward and reverse reactions occur at equal rates and the amounts of reactants and products are constant (NOT EQUAL)
8. LeChatlier’s Principle = if an equilibrium is disturbed (stress such as change in concentration, temperature or pressure) the system moves to reestablish equilibrium by removing the stress.
a. Remember add-away, take-toward.
b. Increase in pressure of gases system moves toward the lower number of gas particles
Unit 7 - Organic Chemistry
Tables P, Q and R
1. Organic chemistry is the study of carbon compounds = mostly nonpolar covalent
2. Naming hydrocarbons (H and C)
a. alkanes = single bonded = CnH2n+2 = single, saturated, substitution
b. alkenes = double bonded = CnH2n
c. alkynes = triple bonded = CnH2n-2
3. Nomenclature = name branches (numbers tell where, not how many) + longest continuous chain.
4. Isomers = same molecular formula, different structural formula. The more C, the more isomers possible
5. Functional groups = see Table R, look for differences
6. Organic reactions
a. substitution = ane + halogen or H2 à 2 products
b. addition = ene or yne + halogen or H2 à 1 product
c. fermentation = sugar à alcohol and CO2
d. esterification = alcohol + organic acid à ester + water
e. polymerization = many monomers à 1 polymer
f. saponification = make soap from fats and base
g. combustion = fuel + O2 à H2O + CO2
Unit 8 – Redox
Table J
1. Oxidation # of elements = 0, oxidation # of molecule total = 0
2. Redox reaction involves a change in oxidation # and both a loss and gain of electrons. Electrons lost = electrons gained
3. LEO GER
a. oxidation = loss of electrons, increase in oxidation #, electrons are written on the right side of arrow in half reactions
b. reduction = gain of electrons, decrease in oxidation #, electrons are written on the left side of the arrow
4. Voltaic cells = battery, spontaneous, oxidation at anode, reduction at cathode, electrons flow from anode to cathode, salt bridge allows ion migration to balance charges
5. Electrolytic cell = not spontaneous, needs a power source, oxidation at anode, reduction at cathode, electrons flow from anode to cathode, used to electroplate metals
6. Table J – a reaction is spontaneous when a more reactive element reacts with a less reactive element in a compound.
7. Balancing redox reactions
a. balance the whole reaction first
b. assign oxidation #’s
c. find the numbers that change
d. write half reactions for the elements that change using coefficients from balancing
Unit 9 – Acids and Bases
Tables K, L and M
1. Acid = H+ or H3O+ in solution, proton donor, concentration of H+ > OH-
a. electrolyte
b. turns litmus red, turn phenolphthalein colorless
c. sour
d. reacts with metals to produce hydrogen, neutralize bases to form water and a salt
e. pH 0-6, Table K
2. Base = OH- in solution, proton acceptor, H+ < OH-
a. electrolyte
b. turn litmus blue, turn phenolphthalein pink/magenta
c. bitter, slippery
d. neutralize acids to form water and a salt
e. pH 8-14, Table L
3. Titration or neutralization - moles of H+ = moles of OH-
a. acid + base = salt + water
b. calculation of unknown molarity – MAVA=MBVB
4. pH = -log[H+], pH = exponent of H+ without sign. Each change in pH unit is a 10X change in concentration (pH 3 is 100 times more acidic than pH 5)
Unit 10 – Nuclear Chemistry
Tables O and N
1. elements with atomic number 83 have all radioactive isotopes
2. radioactive decay particles
a. alpha = charge +2, mass 4, slowest
b. beta = charge -1, mass 0
c. gamma = charge 0, mass 0
3. half life = the time it takes for ½ of the radioisotope to decay
amount / timeamount or 1 if fraction / 0
divide by 2 / find ½ life Table N
add half lives
4. nuclear equations = sum of mass numbers are equal on both sides of arrow. Sum of atomic numbers are equal on both sides of arrow