Download CHEMISTRY FINAL EXAM REVIEW SHEET

CHEMISTRY FINAL EXAM REVIEW SHEET
General Information:
ÆThe Final Examination will cover Chapters 5-6 and 11-19
ÆYou will be provided with a periodic table.
ÆAll problems are simple to do without a calculator, but calculators for the free
response section only.
ÆFree-response questions and problems (12 questions = 65 points) will be followed by
a section of
multiple-choice questions on Scantron (75 questions = 75 points).
Chapter 11: Gases
Terms:
Volume (L)
Pressure (atm, kPa, mmHg)
Partial Pressure/Atmospheric Pressure
Temperature (K = C + 273)
Ideal Gas
Diffusion/Effusion
Formulas:
Boyle’s Law P1V1 = P2V2
Charles’ Law V1 = V2
T1 T2
Gay-Lussac’s Law P1 = P2
T1 T2
Avogadro’s Law V1 = V2
n1 n1
The Combined Gas Law P1V1 = P2V2
T1
T2
The Ideal Gas Law P1V1 = P2V2
T1n1 T2n1
The Universal Gas Law PV = nRT
Examples:
An ideal gas has a volume of 20 L at a certain pressure. By what factor would the volume need to change
to halve the pressure?
1.00 mole of oxygen gas occupies 25 L at 300 K. What is its pressure in atmospheres? (R =
0.0821 L·atm/K·mol)
A sample of hydrogen gas occupies 50 mL at 20°C and 100 Torr. If the volume of the gas is 300 mL at 30
Torr,
what is the Celsius temperature?
List some properties of gases.
If the partial pressures of oxygen and carbon dioxide are 245 Torr and 35 Torr, respectively, and the
atmospheric pressure is 760 Torr, what is the partial pressure of nitrogen gas if it makes up the remainder
of the atmosphere?
Chapter 5: Models of the Atom
Terms:
Proton: +1 charge, located in nucleus, relative mass = 1 amu
Neutron: No charge, located in nucleus, relative mass = 1 amu
Electron: -1 charge, located outside of nucleus, relative mass = 1/1840 amu
Atomic Number = number of protons in an element.
Mass Number = number of protons + neutrons in an element.
In an uncharged element: electrons = protons
Chemical Symbols: 20 Mass number = 20
10
Ne Atomic number = 10
Neon, therefore, has 10 protons, 10 neutrons and 10 electrons.
Isotopes
J.J. Thompson’s Cathode Ray Tube experiment
Ernest Rutherford’s Gold Foil experiment
Niels Bohr’s Model of the atom
Dalton’s Atomic Theory
Orbitals
Energy Levels
Quantum Leap
EXAMPLES:
6. Complete the following table:
Symbol
Atomic Number
Mass Number
Protons
Neutrons
Electrons
Al
14
75
33
9
19
7. Write the atomic symbol of an element which has 6 protons, 6 electrons and 7 neutrons.
8. Give the maximum number of electrons for the following orbitals:
s
p
d
f
9. Give the maximum number of electrons for the following energy levels:
1st
2nd
3rd
4th
10. Write the electron configurations for:
Br
N3Pt
Cu+
Chapter 6: The Periodic Table
Terms:
Dmitri Mendeleev
Periods
Groups or Families
Atomic radius
Ionic radius
Ionization energy
Electronegativity
EXAMPLES:
11. How many electrons are in the valence shells of these atoms and ions?
Cl
Mg
Cr
Se2-
Zn2+
12. Draw Lewis electron-dot structures of the atoms and ions given above.
13. Put these atoms and ions in order from lowest to highest atomic radius.
Ca
Ag
Se2-
Ba2+
O
14. Choose the atom or ion with the highest ionization energy in each pair.
Zn2+ or Zn
Si or S
F or F –
15. Put these atoms and ions in order from highest to lowest electronegativity.
I
Fr
Cu+
N
P3-
16. Name the group to which each of these atoms belongs.
F ____________________________
Na ____________________________
Mg ____________________________
Fe ____________________________
Pb ____________________________
Chapter 12: Chemical Bonding
Terms:
Valence electrons
Lewis (electron dot) structures
Octet rule
Ionic bond
Metallic bond
Covalent bond:
Single bond, Double bond, Triple bond
Coordinate Covalent bond
Lone pair or Unshared pair
Polar vs. Non-polar
VSEPR theory: geometry and shape
Linear, Bent, Trigonal Planar, Tetrahedral, Trigonal Pyramidal.
Examples:
17. Draw Lewis structures for these atoms and ions:
P
Br-
C
Ag+
O2-
18. Which of the following exhibit ionic bonding?
NaCl
H2O
SiO2
Fe
KNO3
N2
19. Draw Lewis structures for the following molecules and ions:
H2O
SO42-
N2
CO2
AlCl3
20. Name the geometry and shape for each of the molecules you just drew.
21. Circle any polar molecules:
H2O
CO
O2
NH4+
PCl5
Au
Chapter 13: Liquids and Solids
Terms:
Intermolecular forces:
London (dispersion) forces
Dipole forces
Hydrogen bonding
Chemical Reactions:
Reactions that produce Water:
C8H18(l) + O2(g)→ CO2(g) + H2O(g)
Combustion of hydrocarbons produces carbon dioxide and water.
HCl(aq) + NaOH(aq) →NaCl(aq) + H2O(l)
An acid and a base react to make a salt and water.
22. List some properties of liquids.
23. List some properties of solids.
24. Name the most important type of intermolecular force experienced by each:
HCl
H2
Br2
H2O
C6H6
CO
Chapter 14: Solutions
Terms:
Solubility
Unsaturated, Saturated, Supersaturated
Miscible vs. Immiscible "Like dissolves like"
Molarity (M)
Formulas:
M = mols % solution = mass of solute x 100
L mass of solution
S1 = S2
P1 P2
Examples:
25. Which compounds would you expect to be soluble in water?
CH3CH2CH2CH3
CH3CH2OH
NaCl
CH3CH2CH2CH2NH2
26. If the solubility of Cl2 gas is 0.50 g/L at 2 atm, what is its solubility at 4 atm?
What is the molarity of a solution that contains 34 g of NH3 in 2.0 L of solution?
Chapter 15: Acids and Bases
Terms:
Arrhenius acid: donates hydrogen ions in solution.
Arrhenius base: donates hydroxide ions in solution.
Brønsted-Lowry acid: donates hydrogen ions in solution.
Brønsted-Lowry base: accepts hydrogen ions in solution.
The pH scale: 0---------------------7---------------------14
Monoprotic/diprotic/triprotic acids
Amphoteric/amphiprotic
Strong acid vs. weak acid
Strong base vs. weak base
Strong electrolyte vs. weak electrolyte
27. What is the pH of solution that has a hydrogen ion concentration of 0.0001?
28. For the neutralization of sulfuric acid:
__ H2SO4(aq) + __ KOH(aq) →__ K2SO4(aq) + __ H2O(l)
What volume of 0.400 M KOH is required to neutralize 5000 mL of a 0.100 M sulfuric acid
solution?
29. Write a net ionic equation for:
2AgNO3(aq) + MgCl2(aq) →2AgCl(s) + Mg(NO3)2(aq)
Chapter 16: Equilibrium
Terms:
Equilibrium
Collision Theory
Exothermic vs. Endothermic reaction profiles
Heat of Reaction
Activation energy
Catalyst
Formulas:
Keq = [products]
[reactants]
Examples:
30. Which direction will the equilibrium shift if the following stresses are applied to the reaction:
2C(s) + H2(g) + O2(g) ↔2CO(g) + 2H+(aq)
Increased pressure in the reaction vessel.
CO removed from the reaction vessel.
Water added to the reaction vessel.
If the reaction is endothermic, what will happen if the reaction vessel is heated?
pH decreased.
NaOH added.
Solid C added.
31. Write the equilibrium constant expression for the reaction, then solve for Keq:
Mg(s) + 2HCl(aq)→ H2(g) + MgCl2(aq)
4.0 M
6.0 M
5.0 M
Chapter 17: Oxidation and Reduction
Terms:
Oxidation vs. Reduction (LEO says GER or OIL RIG)
Oxidizing agent vs. Reducing agent
Oxidation Numbers:
Elements in the free state have an oxidation number of 0.
A monatomic ion has an oxidation number equal to its charge.
Hydrogen is usually +1.
Oxygen is usually –2.
In a compound, the more electronegative element is given an oxidation number equal to its usual ionic
charge.
The sum of the oxidation numbers must equal the overall charge on the compound or ion.
Balancing Redox using half-reactions
Assign oxidation numbers to all elements.
Write half-reactions for oxidation and reduction.
Balance the half-reactions.
Balance all elements apart from O and H.
Balance O by adding H2O.
+
+
-
Balance H by adding H . (If in basic solution, then neutralize the H with an equal amount of OH
on both sides to produce H2O.)
Balance the electrons by adding e-.
Multiply each half-reaction by a coefficient to make the electrons lost in oxidation equal to the electrons
gained in reduction.
Add the two half-reactions together, then cancel species which appear on both sides.
Examples:
32. Assign oxidation numbers to all elements:
MnO4-
NO
H3PO4
Zn2+
Cl2
Al2(SO4)3
33. Balance the following:
Br2 + S2O32- →Br- + SO42- in acidic solution.
CN- + MnO4- →CNO- + MnO2 in basic solution.
Pb
Chapter 18: Nuclear Chemistry
Terms:
Natural radioactivity: alpha
beta
gamma
positron emission
electron capture
Half-life
Artificial radioactivity: transmutations, fission, fusion.
Examples:
34. Write balanced nuclear equations for:
Neptunium-240 experiences alpha decay.
A neutron strikes plutonium-239, producing 2 neutrons, tin-51 and a nuclide.
Mercury-198 captures an electron.
35. If the half-life of cesium-133 is 30 years, how much of a 600g sample of cesium-133 will be
left after 150 years?