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This is a review I made for an introductory chemistry class
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Chem 105 Final ReviewDR. SHIRTS-WINTER 2013
JOANNA WILLIAMS
Disclaimer
All the problems on here come from your textbook! If you need more, I suggest working through micro-exams and practice sheets/parallel example problems, seeing your Learning Community Mentor to run through previous exams, and looking in the book for more problems to work through over specific things your struggling with.
Measurements
Accurate: numbers close to the actual value
Precise: numbers close to each other
Significant Figures:
All non-zeros are sig. figs.
Zeros between two non-zeros are sig. figs.
Zeros left of first non-zero are NOT sig. figs.
If #>or=1, all zeros right of decimal are sig. figs.
If #<1, all zeros at end of # and between non-zeros are sig. figs.
Trailing zeros may or may not be sig. figs. (That’s why we use scientific notation)
Nomenclature
Metal + Nonmetal = Ionic compound
Charges designate formula, name the elements and add –ide to the end
Nonmetal + Nonmetal = Covalent molecule
Use prefixes and add –ide to the end
Polyatomic Ions
Organic Functional Groups
Organic Prefixes
Acids
If you “–ate” too much you feel “–ic”ky
“-ite”s like Nephites and Lamanites are people like “-ous”
Hypo-ous, ous, ic, per-ic increasing O
Hydro-ic
Dimensional Analysis
Use the Mole to Mole Ratio from stoichiometric coefficients in balanced chemical equation
Find limiting reactant
Practice Problems
3.51) Determine the empirical and molecular formulas of each of the following substances:
Styrene, a compound used to make Styrofoam cups and insulation, contains 92.3% C and 7.7% H by mass and has a molar mass of 104 g/mol
Caffeine, a stimulant found in coffee, contains 49.5% C, 5.15% H, 28.9% N, and 16.5% O by mass and has a molar mass of 195 g/mol
Monosodium glutamate (MSG), a flavor enhancer in certain foods, contains 35.51% C, 4.77% H, 37.85% O, 8,29% N, and 13.60% Na, and has a molar mass of 169 g/mol
Practice Problems
3.69) A piece of aluminum foil 1.00 cm square and 0.550 mm thick is allowed to react with bromine to form aluminum bromide.
How many moles of aluminum were used? (density of aluminum=2.699 g/mL)
How many grams of aluminum bromide form, assuming the aluminum reacts completely?
Practice Problems
3.76) Aluminum hydroxide reacts with sulfuric acid as follows:
2Al(OH)3(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 6H2O(l)
Which is the limiting reactant when 0.500 mol Al(OH)3 and 0.500 mol H2SO4 are allowed to react? How many moles of Al2(SO4)3 can form under these conditions? How many moles of the excess reactant remain after the completion of the reaction?
Reactions
Know the solubility rules and the exceptions for precipitation reactions
Oxidation-Reduction (RedOx)
123FHO7654
LEO goes GER / OIL RIG
Oxidizing agents/Reducing agents
Reactions
Acid/Base/Neutralization
Titrations and Dilutions: M1V1=M2V2 (molarity= mol/L)(volume=L)
Net Ionic Equations
Strong Acids (ionize completely)
H2SO4 , HNO3, HCl, HBr, HI, HClO4
Strong Bases (dissociate completely)
Group 1-OH, Group 2-OH from Ca down
pH
pH=-log[H+]
pOH=-log[OH-]
pH+pOH=14
[H+]+[OH-]=1x1014
pH<7 acidic
pH=7 neutral
ph>7 basic
Practice Problems
4.83) Some sulfuric acid is spilled on a lab bench. You can neutralize the acid by sprinkling sodium bicarbonate on it and then mopping up the resultant solution. The sodium bicarbonate reacts with sulfuric acid as follows:
2NaHCO3(s) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) + 2 CO2(g)
Sodium bicarbonate is added until the fizzing due to the formation of CO2(g) stops. If 27 mL of 6.0 M H2SO4 was spilled, what is the minimum mass of NaHCO3 that must be added to the spill to neutralize the acid?
Practice Problems
4.40) Write the balanced molecular and net ionic equations for each of the following neutralization reactions:
Aqueous acetic acid in neutralized by aqueous barium hydroxide
Solid chromium(III) hydroxide reacts with nitrous acid
Aqueous nitric acid and aqueous ammonia react
Practice Problems
4.51) Which element is oxidized and which is reduced in the following reactions?
N2(g) + 3H2(g) 2NH3(g)
3Fe(NO3)2(aq) + 2Al(s) 3Fe(s) + 2 Al(NO3)3(aq)
Cl2(aq) + 2NaI(aq) I2(aq) + 2NaCl(aq)
PbS(s) + 4H2O2(aq) PbSO4(s) + 4H2O(l)
Thermochemistry
Ek = ½(mv2)
1 Cal = 1000 cal
1 cal = 4.184 J (specific heat of water)
∆E=q+w
H=E+PV
Bond enthalpies: reactants – products aka bonds broken – bonds formed
∆Hfo: products – reactants (diatomics in natural state = 0)
q=mCs∆t
Hess’s Law
Practice Problems
5.43) Consider the following reaction:
2Mg(s) + O2(g) 2MgO(s) ∆H = -1204 kJ
Is this reaction exothermic or endothermic?
Calculate the amount of heat transferred when 3.55g of Mg(s) reacts at constant pressure
How many grams of MgO are produced during an enthalpy change of -234 kJ?
How many kilojoules of heat are absorbed when 40.3g of MgO(s) is decomposed into Mg(s) and O2(g) at constant pressure?
Practice Problems
5.56) When a 4.25g sample of solid ammonium nitrate dissolves in 60.0g of water in a coffee-cup calorimeter, the temperature drops from 22.0 C to 16.9 C. Calculate ∆H (in kJ/mol NH4NO3) for the solution process
NH4NO3(s) NH4+(aq) + NO3-(aq)
Assume that the specific heat of the solution is the same as that of pure water.
Is this process endothermic or exothermic?
Practice Problems
5.65) From the enthalpies of reaction
H2(g) + F2(g) 2HF(g) ∆H = -537 kJ
C(s) + 2F2(g) CF4(g) ∆H = -680 kJ
2C(s) + 2H2(g) C2H4(g) ∆H = +52.3 kJ
Calculate ∆H for the reaction of ethylene with F2:
C2H4(g) + 6 F2(g) 2CF4(g) + 4HF(g)
Electrochemistry
E=hv=hc/λ
1/λ=R(1/n12-1/n2
2)
E=-Rhc(1/n2)
λ=h/mv
Bohr’s Model
Practice Problem
6.37) Calculate the energy of an electron in the hydrogen atom when n=2 and when n=6. Calculate the wavelength of the radiation released when an electron moves from n=6 to n=2.
Is this line in the visible region of the electromagnetic spectrum? If so, what color is it?
Orbitals & Nodes
s=spherical
Radial nodes starting at 2s
p=peanut
1 planar node and radial nodes starting at 3p
d=dlover leaf?
2 planar nodes and radial nodes starting at 4d
f
Quantum Numbers
Pauli Exclusion Principle
n=shell (1, 2, 3….)
l=subshell (n-1 to 0) (0=s, 1=p, 2=d, 3=f…)
ml=orientation (-l to l)
ms=spin (-1/2 or +1/2)
Practice Problem
6.56) Which orbital goes with the following quantum numbers? Which are not allowed?
2, 1, -1
1, 0, 0
3, -3, 2
3, 2, -2
2, 0, -1
0, 0, 0
4, 2, 1
5, 3, 0
Electron Configuration
Expanded
Condensed using Noble Gas configuration
Cu and Cr exceptions
Why? Hund’s Rule
Practice Problem
6.61) For a given value of the principal quantum number, n, how do the energies of the s, p, d, and f subshells vary for
Hydrogen?
A many-electron atom?
Periodic Trends
Electronegativity
Size
Ionization Energy
Electron Affinity
Effective Nuclear Charge
Family Names
Cation and Anion Size
Lewis Dot Structures
Count total valence electrons
Least electronegative atom in the middle
Fill octet, create multiple bonds if too many electrons
Resonance structures: none actually what the molecule looks like, it’s a hybrid of all of them
Formal charges
Bond strengths
Molecular Orbitals
Bond Order
MO Diagrams
Paramagnetic vs. Diamagnetic
VSPER
Molecular shapes and Angles
Gases
PV=nRT
22.4 L = 1 mole gas @ STP
Ideal gas characteristics Small, high temp, low pressure
Partial pressures Pa=XaPt
Pt=P1+P2+P3…
Effusion and Diffusion Rates Urms=√(3RT/M)
r1/r2= √(M2/M1)
J=kgm2/s2
Practice Problems
10.54) Calculate the density of sulfur hexfluoride gas at 707 torr and 21 C
Calculate the molar mass of a vapor that has a density of 7.135 g/L at 12 C and 743 torr
Practice Problem
10.69) A piece of dry ice (solid carbon dioxide) with a mass of 5.50 g is placed in a 10.0 L vessel that already contains air at 705 torr and 24 C. After the carbon dioxide has totally vaporized, what is the partial pressure of carbon dioxide and the total pressure in the container at 24 C?
Intermolecular Forces
London Dispersion
Induced Dipole (Polarizability)
Dipole-Dipole
Hydrogen Bonding
Ion-Dipole
Liquids
Intermolecular force effect on
Viscosity
Surface Tension
Phase Changes
Phase Diagrams
Specific heats to heat the substance to melting or boiling point
Heats of vaporization or fusion to melt or evaporate substance
Practice Problems
11.43) For many years drinking water has been cooled in hot climates by evaporating it from the surfaces of canvas bags or porous clay pots. How many grams of water can be cooled from 35 C to 20 C by the evaporation of 60 g of water?
(The heat of vaporization of water in this temperature range is 2.4 kJ/g. The specific heat of water is 4.18 J/gK.)
Colligative Properties
Depends on number of solute particles present, not identity
Don’t forget ions dissociate!
Vapor Pressure ↓
Pa=XaP°
Boiling Point ↑
Freezing Point ↓
Osmotic Pressure ↑
π=iMRT
Solids
Unit Cells
Simple/Primitive Cubic
1 atom
Face-Centered Cubic
4 atoms
Body-Centered Cubic
2 atoms
Concentrations
[ ]=M=moles solute/L solution=molarity
m=moles solute/kg solvent=molality
X=moles substance/total moles=mole fraction
ppm=(mass substance/total mass)x106
mass %=(mass substance/total mass)x100
Equilibrium
kc=[products]/[reactants]
kp=Pproducts/Preactants
aA + bB cC + dD
kc=[C]c[D]d/[A]a[B]b
Le Chatelier’s Principle
Noble gases being pumped in to increase the pressure have no effect (don’t change individual partial pressures)
Catalysts have no effect
Solids and liquids have no effect