Download Just a Few Things 2012

Just a Few Things
To Remember
for the
2012 Regents Chemistry
Exam
—Florian König
Hastings High School
2011/2012
1.
Homogeneous mixtures are called solutions. example: NaCl (aq), air
Solid solutions are often between metals  alloys
Liquid solutions are transparent.
2.
Temperature reprents (NOT: is equal to) the average molecular kinetic
energy.
Heat is the amount of total molecular kinetic energy.
Objects with the same temperature but different mass or heat capacity
have different amounts of total energy.
3.
Endothermic reaction: A + B + energy 
C
Exothermic reaction: A + B

C + energy
4.
# of Joules = mass of water (in gram) x T x 4.2 J / (g K)
5.
Kelvin scale: no negative numbers, based on molecular motion
K = oC + 273
0 K = -273 oC
6.
Boyle’s Law: The volume of a gas at constant temperature is inversely
proportional to its pressure.
(pV = constant  p1V1 = p2V2)
7.
Charles’ Law: The volume of a gas at constant pressure is directly
proportional to its Kelvin temperature.
(V/T = constant  V1/T1 = V2/T2)
8.
Real gases differ from ideal gases when the molecules are close
together, i.e. under conditions of high pressure and low temperatures
(best approximations to ideal gas: H2, He)
9.
vapor pressure: depends only on temperature of a substance, not on its
amount (Ref. H)
10.
Boiling occurs when the vapor pressure = atmospheric pressure
“normal boiling point” = boiling point at 1 atm pressure (101.3 kPa)
11.
energy for (solid  liquid) = “heat of fusion” (change of EPOTENTIAL )
energy for (liquid  gas) = “heat of vaporization”(change of EPOTENTIAL)
12.
Sublimation: solid  gas (no liquid phase)
13.
nucleons: protons and neutrons
mass number = # (protons) + # (neutrons)
atomic number = # (protons)
Energy is absorbed.
Energy is released.
examples: dry ice CO2,
iodine I2
14.
Mass number - atomic number = # (neutrons)
15.
In a neutral atom the positive and negative charges balance each other.
#(protons) = #(electrons)
16.
Isotopes are different kinds of atoms of the same element.
(same number of protons, different number of neutrons)
e.g. 12C/14C, 235U/238U
17.
Atomic mass unit (amu) = 1/12 of the mass of
18.
Atomic mass of an element = weighed average of isotope masses
19.
Ground state:
all lower energy levels are filled, e.g., 2-8-18-4 (Germanium Ge)
(electron configuration on Periodic Table: ground state)
Excited state:
one or more lower energy levels are not filled, e.g., 2-7-18-2
20.
Spectral lines are produced when electrons fall back from higher to lower
energy levels (e.g., PEL 3  PEL 2) and emit energy
21.
Period number = principal energy level (PEL)
22.
Valence electrons: (max. 8) electrons of the outermost principal energy
level (valence shell)
23.
Ionization energy is the energy to remove the 1st, 2nd, 3rd, ... valence
electron (Ref. S)
24.
Electronegativity is the attraction of an atom to other electrons (Ref. S)
High electronegativity = nonmetals (fluorine 4.0), low = metals
25.
Radioactivity (Ref. N & O)
alpha emission (alpha decay)

beta emission (beta decay)

12
C
atomic number decreases by 2,
mass number by 4
atomic number increases by 1,
mass number constant
26.
The rate of radioactive decay (half-life!) is independent of any conditions.
27.
Nonpolar covalent bonds: EN = 0 - 0.4
Polar covalent bonds: EN = 0.4 - 1.7
Ionic bonds: EN > 1.7
Nonpolar molecules: e.g. H2, Cl2, N2, ... (only nonpolar bonds)
CO2, CBr4, CCl4 (polar bonds, but symmetrical)
28.
29.
Polar molecules (dipoles): e.g. HCl, H2O, NH3, CHCl3, CH3COOH
30.
Network solids = covalent bonds, no distinct molecules in the solid
phase, e.g., C (graphite, diamond), SiO2 (quartz), asbestos
31.
Molecular solids = covalent compounds, distinct molecules in the solid
phase, e.g., H2O, CO2, CH4, SO3, etc.
32.
Metallic bonds = cations (kernels) in a “sea of valence electrons”
(“chocolate chip cookie” model)
33.
Van der Waals’ forces (vdW) = weak attraction between nonpolar
molecules, e.g., He, CO2, hydrocarbons
34.
Dipole-dipole interactions = attraction between polar molecules
H bonds = dipole-dipole interactions between H2O/HF/NH3 molecules
35.
Ionic solids
=
high m.p and b.p, hard, conductivity only when
molten or dissolved
Molecular solids
=
low m.p. and b.p., soft, no conductivity
Metallic solids
=
high m.p. and b.p., varying hardness, high
conductivity
36.
Empirical formula =
Molecular formula =
lowest atom-atom ratio
multiple of empirical formula
e.g. CH (empirical formula) 
C2H2, C6H6 (molecular formula)
37.
nonmetals:
metals
r(atom) < r(anion)
r(atom) > r(cation)
38.
Most active nonmetal (strongest oxidizing agent): fluorine F2
Most active metal (strongest oxidizing agent): cesium Cs (or francium Fr)
(Ref. J)
39.
Chemical groups are the columns in the Periodic Table. Elements in
groups have similar chemical properties.
Group 1 (I A)
Group 2 (II A)
Group 17 (VII A)
Group 18 (VIII A)
alkali metals
alkaline earth metals
halogens
noble gases
40.
Transition metals = metals with e- in d-orbitals
Transition metal compound are often colored (CuSO4 blue, K2CrO4
yellow) and have more than one oxidation state.
41.
Monoatomic molecules: noble gases He, Ne, Ar, Kr, Xe, Rn
Diatomic molecules: Br2, I2, N2, Cl2, H2, O2, F2
42.
1 mole = 6.02 x 1023 particles (atoms, ions, molecules, formula units)
mass of 1 mole = atomic (formula) mass in gram
43
Percentage composition =
mass of element
__________________________________
(Ref. T)
formula mass of compound
44.
Molar volume = volume of 1 mole of any gas at STP = 22.4 liters
(STP = 0 oC and 1 atm)
45.
Avogadro’s Law: Equal volumes of gas at equal temperature and
pressure contain equal numbers of molecules.
46.
You should be able to solve mass-mass, mass-volume and volumevolume problems (or do the same thing with moles)
47.
Balanced equation gives: a) mole ratio
b) volume ratio (for gases only)
moles solute
48.
Molarity M =
___________________________
(Ref. T)
liters solution
(solute: smaller part of solution; solvent: bigger part)
49.
Dissolved solute LOWERS freezing point and RAISES boiling point.
Freezing and boiling point constants for water: (first page of Ref. tables)
50.
Heat of reaction = H = Hproducts – Hreactants
H < 0 exothermic
(see Ref.I for examples)
H > 0 endothermic
51.
Catalyst lowers activation energy
 for both forward and reverse reaction
 speed up reaction
 do not change position of equilibrium
52.
LeChatêlier’s Principle: A system at equilibrium reacts to an external
stress in such a way that the stress is relieved.
e.g.
temperature increases
temperature decreases
pressure increases 
pressure decreases 
reactant is added 
etc. pp.

endothermic reaction favored

exothermic reaction favored
side with fewer gas molecules favored
side with more gas molecules favored
shift towards product side
53.
Two driving forces: a) enthalpy H tends to be minimal
b) entropy S tends to be maximal
54.
Entropy S increases with:



55.
phase:
solid < liquid < solution < gas
temperature:
S (20 oC) < S (70 oC)
# of particles, amount
Acids and bases (Ref. K & L)
Strong acids: hydrochloric acid HCl, nitric acid HNO3, sulfuric acid H2SO4
Weak acids: carbonic acid H2CO3, acetic acidCH3COOH
Strong bases:
Weak bases:
alkali hydroxides (LiOH, NaOH, KOH)
ammonia NH3
56.
Brønsted acid = proton donor
Brønsted base = proton acceptor
BAAD (Bases Accept, Acids Donate)
57.
acid
58.
Amphoteric substances can act as acids or bases.
59.
60.
Titration:
MA x VA = MB x VB
Acidic salts: NH4Cl, (NH4)2SO4
conjugate base + H+
Basic salts: K2CO3, Na2SO3
Neutral salts: KNO3, Li2SO4
61.
Strong acids: H2SO4, HNO3, HCl, HI
62.
pH = -log[H+]
pH
<7
=7
>7
[H+]
> 10-7 M
10-7 M
< 10-7 M
The solution is…
acidic
neutral
basic
[OH-]
< 10-7 M
10-7 M
> 10-7 M
63.
Oxidation numbers (see separate handout with rules and common
oxidation states)
64.
-4
-3
-2
-1
0
———————— oxidation
<———————— reduction
65.
Hydrides: Compounds with negatively polarized hydrogen (ox. state = -1)
(LiH, NaH, KH)
66.
Oxidizing agents = oxidize something else, get reduced
Reducing agents = reduce something else, get oxidized
67.
Reduction strengths (Ref. J)
Strong oxidizing agents are on top right (F2, Cl2), strong reducing
agents on top left (Li, Rb, K, …)
68.
Electrochemical cell:
Electrolytical cell:
69.
In BOTH types of cell the same types of reaction occur at the same
electrode:
ANode — OXidation
+1
+2
+3
+4
———————————>
———————————
chemical energy
electrical energy


electrical energy
chemical energy
REDuction — CAThode
70.
Objects being plated in an electrolytical cellare always wired as the
(negative) cathode.
71.
The farther away two metals are from each other in Ref. J, the bigger the
voltage in an electrochemical cell with them.
72.
Organic compounds = covalent, molecular compounds based on carbon
(Ref. P, Q, R)
usually nonpolar, low m.p./b.p., low reactivity
73.
Isomers = same molecular formula, but different structure
74.
Saturated = only single bonds (alkanes)
Unsaturated = multiple bonds (alkenes or alkynes)
75.
Alkanes
Alkenes
Alkynes
Benzenes
CnH2n+2
CnH2n
CnH2n-2
CnH2n-6
(cyclic, alternating single/double bonds)
76.
Ethyne = acetylene = C2H2 H—CC—H
77.
Benzene series:
78.
Alcohols R — OH do not ionize (OH is covalently bound)
benzene C6H6
Type of alcohol
# of OH groups
monohydroxy ~
1
dihydroxy ~
trihydroxy ~
2
3
Type of alcohol
Structural
characteristic
Primary
|
R — C — OH
|
secondary
R1
|
2
R — C — OH
|
tertiary
R1
|
R2 — C — OH
|
R3
toluene C6H5—CH3
Example
CH3OH
methanol
CH3CH2OH ethanol
C2H4(OH)2 ethylene glycol
C3H5(OH)3 glycerol
Example
ethanol CH3CH2OH
propanol CH3CH2CH2OH
2-propanol
(CH3)2CH — OH
2-methyl-2-propanol
(CH3)3C — OH
79.
Organic acids (carboxylic acids) = R — COOH
80.
Acid + Alcohol

e.g. CH3CH2COOH + CH3CH2OH 
81.
Aldehydes
Ketones
Ethers
Amines
82.
Addition
Ester + water
CH3CH2COOCH2CH3 + H2O
R—CHO
R1—CO—R2
R1—O—R2
R—NH2
Substitution
Polymerization
Cracking
Saponification
to multiple bonds (e.g. hydrogenation of unsaturated
fats)
at single bonds (saturated compounds)
small building blocks  large molecules
large molecules  small building blocks
hydrolysis of fats with lye  fatty acids (soaps) +
glycerol
83.
Nuclear Chemistry (Ref. N & O):
a)
All isotopes with an atomic number bigger than 83 are unstable.
b)
Neutrons (no charge!) are not accelerated in a particle accelerator.
c)
Mass defect = sum of neutron and proton masses is less than
nucleus mass
84.
Nuclear reaction:
a) fission = splitting of nuclei (naturally or after neutron bombardment)
b) fusion = two nuclei are combined
4
c) types of radiation 
He nuclei (mass = 4, charge = +2)

electrons (mass = 0, charge = -1)

high-energy radiation (no mass or charge)
d) nuclear reactors 
U-235 or U-238 as fuel

moderator slow down neutrons

control rods absorb neutrons
85.
Radioisotopes
86.
Significant figures: see separate handout
87.
Do as you oughta, acid to water!
88.
Percentage error: Ref. T
89.
Applications of chemical principles
Tc-99 to detect brain tumors
I-131 to treat thyroid gland
Haber Process
N2 + 3 H2
2 NH3
(Ammonia is used for fertilizers and explosives.)
90.
Refining of metal ores
– or –
CuS + O2

Fe3O4 + 2 C 
reaction of metal oxide with carbon
reaction of metal sulfide with oxygen
Cu + SO2
3 Fe + 2 CO2
91.
Fractional distillation = separates petroleum into different parts with
different boiling points
92.
READ THE QUESTIONS CAREFULLY!
93.
Read the Reference Tables. Sing the Reference Tables. Be the
Reference Tables.
94.
Do not study the night before. Get a good night’s sleep. Have breakfast.