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COMPACT REGENTS REVIEW 2010-2011
MATTER
Pure matter = substances, which are represented by (s), (l) or (g)
Substances = elements (identical atoms) or compounds (fixed combinations of different elements)
Compounds can be chemically decomposed, Elements cannot.
Solution (aq) = homogeneous mixture = uniformly distributed substances in varying composition
Filtration = method of separating heterogeneous mixtures
Distillation = method of separating homogeneous mixtures using boiling point difference
Chemical Change = atoms are rearranged = combustion or rusting
Physical Change = the substance does not change, only the physical form of the substance
ENERGY
Potential Energy = energy of phase changes
Highest Kinetic Energy = highest temperature
Phase Change = no temperature change: q = m Hf (freezing / melting) or q = m Hv (boiling)
Temperature Change: q = m x c x T (for water, c = 4.18 J/goC)
Exothermic Changes: condensation (g  l), freezing (l  g), deposition (gs)
Endothermic Changes: boiling (lg), melting (ls), sublimation (sg)
Heating Curve plateaus = phase changes: PE increases, KE is constant
Heating Curve slopes = temperature increases = KE increases
ATOMIC STRUCTURE
Subatomic Particles = protons, neutrons and electrons
Orbital = 3D region outside of nucleus where electron is found 90% of the time
Wave-Mechanical Model = most modern, the theory of electrons in orbitals
Mass = p + n (Carbon-14 has a mass of 14, with 6 protons and 8 neutrons)
Net Charge = p – e (An atom of Carbon-14 has a net charge of 0, 8 protons and 8 electrons)
Nuclear Charge = p (Carbon -14 has a nuclear charge of +8)
Isotopes of an Element = same # of protons, different # of neutrons
Bright line spectra = spectral lines = light
Light is produced when an electron transitions from higher to lower energy state
Ground State Electron Configuration = found on Periodic Table: Mg 2 – 8 – 2
Excited State E.C. = an electron moved to a higher level: Mg 2 – 8 – 1 – 1
Average atomic mass = [(mass isotope #1) x (%) + (mass #2) x (%) + …] / 100%
First subatomic particle discovered = electron
Gold Foil Experiment conclusions: atom is mostly empty space with a massive (+) nucleus
NUCLEAR CHEMISTRY
Alpha particle 42He = most massive, positive, least energetic
Beta particle 0-1e = negative
Gamma ray 00y = massless, neutral, most energetic
Transmutation = Nuclear Change
Natural Transmutation = Decay = Table N = 1 reactant: 19779Au  0-1e + 19780Hg
Artificial Transmutation = 2 reactants: 147N +10n  146C + 11p
Radioisotope = unstable nucleus
Atomic # 84 and above = only radioactive isotopes
1  ½  ¼  1/8 = 3 half live events
Half-Life Formula: # of decay events = time / half-life
Medical isotopes = short half lives = 131I (thyroid) and 60Co (cancers)
Dating Isotopes = long half-lives = 238U (rocks) and 14C (organic remains)
Fusion = 21H + 21H  42He + energy = two smaller nuclei are fused together
Fission = “division” = large nucleus is split into two smaller nuclei = 23592U + 10n 
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COMPACT REGENTS REVIEW 2010-2011
Mass Defect = the mass that seemed to be “lost” was converted into energy (fission and fusion)
BONDING
BARF = Break (bonds) Absorb (energy) , Release (energy) Form (bonds)
Metallic Bonding (all metals) = valence electrons are mobile
Covalent Bonding (all non-metals) = valence electrons are shared
Polar Covalent = unequally shared (two DIFFERENT non-metals)
Non-Polar Covalent = equally shared (two of the SAME non-metals)
Ionic Bonding (metal + nonmetal) = valence electrons are transferred
Most Polar Bond = biggest Electronegativity Difference (H – F)
Most Ionic Bond = biggest Electronegativity Difference (Fr F)
Metallic substances: high MP and BP, malleable, very conductive due to mobile electrons
Molecular substances (covalent bonds): low MP, brittle solids, non-conductors
Ionic substances: high MP, brittle, good conductors as (l) or (aq); non-conductors as (s)
Oh, SNAP!
Symmetrical Non-Polar, Asymmetrical Polar (for molecule polarity)
Order of increasing intermolecular forces: (g) < (l) < (s)
Order of increasing intermolecular forces: nonpolar molecules < polar molecules
Order of increasing intermolecular forces: low melting/boiling pt < high melting/boiling pt
PERIODIC TABLE
Electronegativity = attraction for a pair of bonded electrons
Ionization Energy = energy needed to remove a specific electron
Metals: to the left of the staircase, including Al and Po. All solid except for Hg
Metalloids: on the staircase: B, Si, Ge, As, Sb, Te, At: fair conductors but brittle
Non-Metals: to the right: solids, liquids (Br2) and gases (H2, N2, O2, F2, Cl2 + noble gases)
BrINClHOF: the seven diatomic elements
All elements want to have Noble Gas Configurations: Mg (2-8-2) becomes Mg2+ (2-8) Neon!
Metal atoms LOSE ELECTRONS to form smaller, positive ions
Non Metal atoms GAIN ELECTRONS to form larger, negative ions
Down a Group, elements get larger (more shells) and more metallic = lower IE and EN
Across a Period, elements get smaller and less metallic = higher IE and EN
Most metallic element = Fr = loses electrons most readily
Leas metallic element = F = gains electrons most readily
Allotropes = different forms of same elements with different structures and different properties
Elements in same group = same # of valence electrons = similar chemical and physical properties
Elements in same period = same # of occupied energy levels
GAS LAWS
Standard temperature = 273 K (0 oC)
Standard pressure = 1 atm or 101.3 kPa
Table H Normal Boiling Point = temperature where substance curve intersects dotted line (1 atm)
Ideal gas molecules= no attractive forces = negligible volumes = large intermolecular distances
Most Ideal Molecules = H2 and He
Ideal conditions (think Summer) = high temperature, low pressure (except for Boranian, in summer school)
As Pressure increase, Volume decreases (inverse relationship)
P x V = constant
As Temperature increases, Volume increases (direct relationship) V/T = constant
Combined Gas Law: P1V1/T1 = P2V2 / T2
temperature is in Kelvin
Equal Gas Volumes, Equal #’s of (moles of) gas molecules…at the same T and P
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COMPACT REGENTS REVIEW 2010-2011
FORMULAS, EQUATIONS & MOLES
A BALANCED EQUATION demonstrates the Law of Conservation of Mass:
2 H2 + O2 --> 2 H2O (NO ATOMS ARE CREATED OR DESTROYED)
REACTIONS:
SYNTHESIS: C + O2 --> CO2
DECOMPOSITION: CO2 --> C + O2
SINGLE REPLACEMENT: Element + Compound --> Element + Compound
DOUBLE REPLACEMENT: Compound + Compound --> Compound + Compound
WRITING FORMULAS FROM NAMES:
Iron (II) Oxide --> Fe+2 O-2 --> Fe2O2 --> FeO
Iron (III) Oxide --> Fe+3 O-2 --> Fe2O3
MOLECULAR formulas (C6H12) reduce to simplest whole-number EMPIRICAL formulas (CH2)
mole formula:
moles = mass / gram-formula mass
Atomic Mass: The mass of ONE MOLE of an element (Carbon 12.011 grams / mol)
Percent Composition formula: % = (mass of part / mass of whole) x 100%
SOLUTIONS
Characteristics: homogeneous mixtures; do not scatter light; do not settle on standing;
cannot be separated by filtration; can have color
In NaCl (aq) Solute: NaCl
Solvent: H2O
Gaseous solutes are more soluble at LOW TEMPERATURES and HIGH PRESSURES
Solid & Liquid solutes are more soluble at HIGH TEMPERATURES. PRESSURE has no effect.
Water is a POLAR MOLECULE: the hydrogens are (+) the oxygen is (-)
Polar solvents (H2O) dissolve polar solutes (NH3) and ionic solutes (NaCl)
Non-Polar solvents dissolve Non-Polar solutes
Molarity Formula:
Molarity = moles of solute / L of solution
ppm Formula:
ppm = (mass of solute / mass of solution) x 1,000,000
TABLE F is used to determine which IONIC SOLUTES are soluble in water
TABLE G is used to determine if aqueous solutions are saturated, unsaturated or supersaturated
KINETICS & EQUILIBRIUM
The more effective collisions, the faster the reaction rate.
The greater the surface area / concentration / temperature / pressure, the more effective collisions.
IONIC reactants (NaCl + AgNO3) react faster than MOLECULAR ones (CH3OH +CH3COOH)
Heat of Reaction (∆H) = potential energy of products - potential energy of reactants
= the amount of energy absorbed (+) or released (-) in a reaction
TABLE I gives the ∆H for a specific number of moles of product:
H2 (g) + I2 (g) --> 2 HI (g)
+53 kJ
Producing 2 moles of HI absorbs 53 kJ of heat
Catalysts increase reaction rate by lowering activation energy / providing lower energy pathway
Catalysts have NO EFFECT on ∆H (heat of reaction)
Activation Energy = HUMP of PE diagram - REACTANT line of PE diagram
An EXOTHERMIC PE diagram has higher REACTANTS lower PRODUCTS
An ENDOTHERMIC PE diagram has lower REACTANTS higher PRODUCTS
Equilibrium exists only in CLOSED systems when...
CONCENTRATIONS (masses) are CONSTANT
RATES of forward and reverse reactions are EQUAL
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COMPACT REGENTS REVIEW 2010-2011
In SATURATED solutions, the RATES of dissolving and crystallizing are EQUAL
LECHATELIER’S PRINCIPLE:
a system that loses equilibrium due to stress will shift to re-establish equilibrium
STRESSES: change in temperature / change in concentration / change in pressure
CATALYSTS speed up both forward and reverse reactions; have NO EFFECT on equilibrium
Equilibrium SHIFTS AWAY FROM an increase in concentration or temperature (heat)
Equilbrium SHIFTS TOWARDS a decrease in concentration or temperature (heat)
When PRESSURE increases, equilibrium SHIFTS TOWARDS the side with fewer gas molecules
ACIDS, BASES & SALTS
Electrolytes (Acids, Bases and Salts) release ions in water, which then conduct electricity.
The higher the concentration (M) of a solution with electrolytes, the more conductive.
TABLE K: Common Acids = H with (TABLE E) ion or Group 17 Ion
TABLE L: Common Bases = metal hydroxide or NH3
Salts are ionic compounds = metal + nonmetal
Acidic solutions have pH < 7 / more hydronium than hydroxide / react with metals to form H2 (g)
Basic solutions have pH >7 / more hydroxide than hydronium / react with fat to form soap
In a pH of 1, the H3O+ concentration is 10-1 M
In a pH of 13, the H3O+ concentration is 10-13 M
Compared to a pH of 8, a pH of 2 is (8-2 = 6) 106 times more acidic
Compared to a pH of 4, a pH of 6 is (6-4 = 2) 102 times less acidic
TABLE M: BROMTHYMOL BLUE is yellow below 6.0 and above 7.6 is blue
ARRHENIUS ACIDS release H+ as the only positive ion
ARRHENIUS BASES release OH- as the only negative ion
ACIDS are PROTON (H+) DONORS, whiles BASES are PROTON ACCEPTORS
In NH3 + HCl --> NH4+ + Cl- The acids are HCl & NH4+ / The bases are NH3 & ClNeutralization reaction: Acid + Base
---> Salt + Water
HCl
+ NaOH ---> NaCl + H2O
Titration formula: Ma x Va = Mb x Vb
(molarity of acid) x (volume of acid) = (molarity of base) x (volume of base)
When a solution is neutralized in a titration, moles of OH- equal moles of H3O+
REDOX (or REDUCTION-OXIDATION)
The sum of oxidation # in a formula must equal zero
NaCl: (+1) + (-1) = 0
Na2S2O3: 2(+1) + 2x + 3(-2) = 0
oxidation # of sulfur = +2
Free elements have an oxidation # of zero
A chemical reaction in which oxidation numbers change is REDOX
A chemical reaction with a FREE ELEMENT is REDOX
In the oxidation half-reaction, oxidation # increases as electrons are lost [OIL]
Fe --> Fe+2 + 2e(0) = (+2) + (-2)
In the reduction half-reaction, oxidation # is reduced as electrons are gained [RIG]
Ag+ + 1e- --> Ag
(+1) + (-1) = (0)
for OXIDIZING / REDUCING AGENTS: ROA ORA as in “ROW”-a with the “OAR”-a
R(OA) the Reduced is the Oxidizing Agent
O(RA) the Oxidized is the Reducing Agent
TABLE J: If the free metal is higher than the combined metal, the redox is spontaneous
3 Li + Al(NO3)3 --> 3 LiNO3 + Al
If the free nonmetal is higher than the combined nonmetal, the redox is spontaneous
F2 + 2 I- --> I2 + 2 FVoltaic Cell: spontaneous redox in which chemical reaction produces electricity
(-)ANODE: where oxidation takes place...the metal HIGHER on Table J (AN OX)
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COMPACT REGENTS REVIEW 2010-2011
(+)CATHODE: where reduction takes place...the metal LOWER on Table J (RED CAT)
electrons flow from Anode to Cathode through the WIRE
The SALT BRIDGE maintains electrical neutrality by permitting ion flow
Electrolytic Cell: non-spontaneous redox in which electricity drives a chemical reaction
Unlike the Voltaic Cell: (1) no salt bridge (2) requires a battery (3) cathode (-) & anode (+)
Like the Voltaic Cell: (1) cathode reduction & anode oxidation (2) electrons flow from An to Cat
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