CHEM 20 FINAL EXAM:
STUDY HEADINGS Jan 2011
A.
Introduction to Chemistry (Ch. 1,2,3)
physical and chemical change; physical and chemical properties;
accuracy and precision in measurements; properties of mixtures, elements, compounds
calculations using significant digits and unit analysis method
B.
Atomic Structure and Theory (Ch 4,5)
early theories of atomic structure; properties of subatomic particles: proton, neutron, electron
atomic number and average atomic mass; isotope symbols;
Dalton, Thomson and Rutherford atomic models
The Bohr atomic model; energy levels; emission and absorption spectra
Quantum mechanics model: Schrödinger equation, Pauli exclusion principle; Heisenberg uncertainty
principle quantum numbers; orbital types and shapes; electron configurations; stability
C.
Nuclear Chemistry (Ch. 28)
types of nuclear radiation and radioactive decay: alpha, beta and gamma
balance nuclear equations: transmutations, fissions, fusions, decays,
decay series, half life equations and problems
Einstein’s equation and mass changes in nuclear reactions
uses of nuclear energy
D.
The Periodic Table/ Periodic Trends (Ch. 6,10,11)
classification of elements; metals, semi-metals, non-metals; liquids, gases, solids;
main groups 1 ,2, 13 -18: properties of; valence electron configurations
transition metals: properties of; Lewis diagrams for atoms
periodic trends: ionization energy, metallic character, atomic radius, electron affinity, electronegativity;
characteristic properties of the chemical families
predict type of ion formed, number of valence electrons, ionization energy, by position in periodic table
E.
Nomenclature/Chemical Formulae (Ch. 7)
names of elements, symbols of elements; ionic compounds, molecular compounds
ionic charge and oxidation states; naming ionic compounds (formula units)
predicting formula units from ion charges; naming molecular compounds (molecules);
formulas of hydrated salts
F.
Chemical Reactions and Stoichiometry (Ch. 8,9)
finding the molar mass of elements and compounds
using Avogadro’s number: converting moles to number of particles, and vice versa
finding moles of a substance from a given mass and vice versa
determining empirical formula and molecular formula from experimental data
percentage composition and formulas of hydrated crystals
writing balanced equations; using subscripts to indicate phases
energy changes in reactions: endothermic and exothermic reactions
naming types of chemical reactions: synthesis, decomposition, combustion, dissolving( ionizations and
dissociations), single displacements, double displacements
predicting the products of single displacements using the activity series for metals and halogens
predicting the products of double displacements using solubility tables
predicting the products of combustion reactions
mass stoichiometry questions: moles to moles, mass to mass, mass to moles, moles to mass
energy questions using stoichiometry
limiting reagent problems – theoretical yield, excess reagents
G.
Chemical Bonding (Ch.12, 13, 14.1,17.2
use difference in electronegativity to identify types of bonds: covalent, polar covalent, ionic
types of bonds: ionic, molecular, polar covalent; bond dipole moments
molecular radii ( bond lengths ); covalent radius, vanderwaals radius
drawing Lewis dot structures for elements and molecules; expanded valence
predicting bond angles, molecular geometry based on lewis dot diagrams
drawing correct shapes for simple molecules
hybridization orbitals of the central atom: sp, sp2, sp3, sp3d and sp3d2
molecular bonds: sigma and pi bonds; delocalized pi bonds in benzene, C6H6
predicting the polarity of molecules from dipole moments and molecular geometry
intermolecular forces: van der waals, hydrogen bonding, dispersion forces, ionic forces,
covalent network solids
H.
Physical Chemistry: Solutions and Gases (Ch 14, 20, 21, 15, 18, 19)
qualitative properties of solutions: solvent, solute, solubility, effect of temperature and pressure on
solubilities of solids and gases in water;
Henry’s law: solubility of a gas in a liquid increases with increased partial pressure of the gas
Colligative properties of solutions: boiling point elevation, freezing point depression, vapor pressure
lowering and osmotic pressure of solutions
effects of molecular structure, intermolecular forces on solubility like dissolves like…principle,
hydrogen – bonding, polarity of molecules
determining the concentration of solutions using: Molarity (M, mol/L)
using c = n/v to solve for mass of solute, volume of solution, or Molarity of solution
preparation of solutions; diluting solutions; using the dilution formula C1V1 = C2V2
writing dissolving equations for ionic compounds, use mole ratios to determine molar concentrations of
ions in solution;
writing net ionic equations; equations to show precipitate formations
Temperature and Pressure unit conversions
Avogadro’s hypothesis; molar volumes, STP and SATP, gas stoichiometry questions
applying the Gas Laws in calculations: Boyle’s, Charles’, Guy-Lussac’s, Combined, and Ideal
I.
Organic Chemistry (Ch. 29)
naming organic compounds; draw isomers for given organic compounds;
identifying the basic hydrocarbons: alkanes, alkenes, alkynes, and cyclic hydrocarbons;
use prefixes to identify substituted groups; alkyl groups, halogens, and a few others
saturated and unsaturated hydrocarbons: general formulas for the main groups;
identify the functional groups by structure and name: alcohols, ethers, aldehydes, ketones,
carboxylic acids, esters and amines
name, draw, structural formula, indicate bond angles for organic molecules;
identify, name and draw structural formulas for aromatic compounds: benzene and its derivatives
outline some of the use of hydrocarbons in modern society
FINAL EXAM REVIEW: 2011 January
The first part of this review are typical multiple-choice style questions. There is only one correct answer per
question:
1.
Which of the following electron dot diagrams is not likely to exist in nature?
a) H – C ≡ C – H
b) He
c) H – F
d) H – O – O – H
2.
Which of the following is the correct molecular formula for the compound, tetraphosphorus
heptasulfide?
a) K4S7
b) P4S7
c) P4S2
d) P4(SO4)7
3.
The IUPAC name for CH3 – C ≡ C – CH3 is:
a) butyne b) butene c) 2,3 – butyne d) 2 – butyne
4.
The IUPAC name for the hydrocarbon to the right is:
a) 2,4,4 – trimethylhexane
b) 2,3,3 – trimethylhexane
c) 2,3,5 – trimethyl-4-isopropylhexane
d) 2,3,5 – triethylhexane
5.
Which is the correct Lewis dot structure for AlF3 ?
a) F – Al – F
b) F – Al – F
c)
F – Al – F
CH3
|
CH3 – CH – CH – CH2 – CH – CH3
|
|
CH3
CH3
d)
F – Al – F
|
|
|
|
F
F
F
F
6.
The correct formula unit for the hydrated salt, potassium carbonate dihydrate, would be:
a) K2CO3.2H2O
b) P2CO3.2H2O
c) K2CO3.2H
d) K2 CO3. 2OH
7.
In an experiment an aqueous solution of Ca(NO3)2 reacts with an aqueous solution of Na3PO4 in a
replacement reaction. Indicate the formula for the resulting precipitate:
a) Ca2(PO4)3
b) Ca3(PO4)2
c) NaNO3
d) Na2NO3
8.
Which of the following is a network covalent solid?
a) iron
b) sodium chloride
c) graphite
d) ice
9.
In which of the following compounds can NO geometric isomers occur?
a) 1 – pentene b) 2 – pentene c) 2 – methyl – 2 – propene d) 2 – pentyne
10.
Which of the following compounds is not classified as an aromatic compound?
a) benzene b) propene
c) toluene d) 3 – bromophenol
11.
A student dissolves 6.84 g of aluminum sulfate in enough water to form 250 ml of solution.
What is the molar concentration ( M ) of the solution?
a) 0.020 M
b) 0.080 M
c) 50.0 M d) 200 M
12.
Dinitrogen oxide can be prepared by the thermal decomposition of ammonium nitrate according to
the following equation: NH4NO3(s)
→ N2O(g) + 2 H2O(g)
When 1.000 g of ammonium nitrate is decomposed in this way, 0.550 g of dinitrogen oxide is
produced. What mass of water would also be produced in this reaction?
a) 1.10 g
b) 0.450 g
c) 0.900 g
d) 0.550 g
13.
Two organic compounds have the following structural formulas:
H H H H H
H H H H
|
|
|
|
|
H–C–C–C–C–C–H
|
|
|
|
H OH H H
|
H
|
|
|
|
H – C – C – C – C – OH
|
|
|
H CH3 H
|
H
Which one of the following statements is incorrect?
a) they are isomers
b) they have different melting points
c) they are both alcohols
d) they both have the same IUPAC name
14.
Which of the following compounds has the smallest molar mass?
a) NH3
b) H2O2
c) CH4
d) SO3
15.
Which of the following molecules is most likely to contain a triple bond?
a) O3
b) HCN
c) CH4
d) SCl6
16.
The molar mass of acetone, CH3COCH3, is:
a) 58.0 g/mol
b) 31.0 g/mol
c) 74.0 g/mol
d) none of these are correct
17.
A certain compound has a molecular mass of 42.0 g/mol. The simplest formula for the compound is
most likely:
a) NH3
b) CH4
c) C2H2O
d) BF
18.
What is the concentration of chloride ions, in mol/L, if 28.5 g of magnesium chloride is dissolved in
enough water to form 900 ml of solution?
a) 0.67M b) 0.60M
c) 0.33M
d) 0.033m
19.
Identify the correct IUPAC name for the compound shown to the right:
a) 1,2 – dimethylcyclohexane
b) 1,2 – dimethylbenzene
c) 1,2 – diethylbenzene
d) 1,3 – dimethylcyclohexane
CH3
CH3
20.
In which of the following molecules is the central atom NOT expanding its valence?
a) SF6
b) PF5
c) AsBr5
d) PCl3
21.
Which one of the following bonds would be expected to have the highest ionic character?
a) CsF
b) MgBr
c) AlO
d) HI
22.
Which of the following sets of molecules, are all the shapes linear?
a) CO2 , HCN and N2 b) HC≡CH, XeF2 and O3 c) H2S, CO and CO2 d) H2O, Cl2 and NO2
23.
The molecular shape of the CCl3+ ion would best be described as:
a) linear
b) trigonal planar c) pyramidal d) square planar
24.
Substances which, added to a reaction vessel, to make reactions happen at a faster rate are known as:
a) solutions b) hydrocarbons c) catalysts d) reagents
25.
Which of the following is an example of combustion?
a) sodium metal + chlorine gas produces sodium chloride solid
b) propane + oxygen produces carbon dioxide and water
c) sulfuric acid decomposes into hydrogen gas, oxygen gas and sulfur
d) solid lead metal reacts with a silver nitrate solution to produce lead II nitrate and silver metal
26.
What is the correct formula unit for Lead IV oxalate?
a) Pb4C2O4
b) Pb(C2O4)2
c) PbC2O4
d) Pb(C2O4)4
27.
The empirical formula of a compound which contains 75.0 g of carbon, 12.5 g of hydrogen and
100.0 g of oxygen is:
a) CH3O
b) C2H4O2
c) C2H3O2
d) CH2O
28.
Which of the following chemical equations is an example of a decomposition reaction?
a) 4 P4 + 5 S8 → 4 P4S10
b) 2 C2H6 + 7 O2 → 2 CO2 + 3 H2O
c) KClO3 → KCl + 3 O2
d) P4 + 5 O2 → P4S10
29.
How many moles of acetate ions, CH3COO-, are in 46.0 ml of a 0.250M solution of Ca(CH3COO)2?
a) 11.5 moles b) 5.43 moles c) 1.15 moles d) 0.0115 moles e) 0.023 moles
30.
Which of the following bond pairs has the highest ionic character?
a) Li and O b) Se and F c) Sr and Br d) Ca and S
31.
Which of the following is not a property of metals?
a) electrical conductors in all phases
b) lustrous – reflects light
c) low melting points
d) delocalized electrons
32.
Which of the following processes would strengthen very soft metals?
a) add gases
b) temper the metal with high temperatures
c) add elements from the middle of the periodic table such as Carbon
d) allow the metals to oxidize
33.
Which of the following molecules would have an octahedral shape?
a) H2CO b) BF3 c) SeO2 d) SF6 e) S2Cl2 f) SF4
34.
Which of the following ions would have a tetrahedral shape?
a) NO3- b) SiF62- c) SO32- d) NH3+ e) PO43-
35.
Which of the following bonds is(are) polar covalent?
a) B – F
b) C – O
c) Si – F
d) N – H
e) O – H
36.
Vials of the following gases are opened. Which of the gases would you expect to smell across the
room first?
a) chlorine
b) sulfur dioxide c) hydrogen sulfide d) ammonia
37.
Which of the following would have the highest melting point?
a) NaCl b) GeO2
c) H2
d) ZnCu
e) C2H5OH
38.
52.0 g of K2CO3 are dissolved in 518 g of distilled water, what is the concentration of the resulting
solution, in mol/L, M ?
a) 0.100M b) 7.27 x 10-4M c) 0.727M
d) 9.96M
39.
A sample of a gas occupies 1.40 × 103 mL at 25°C and 760 mmHg. What volume will it occupy at the
same temperature and 380 mmHg?
a) 2,800 mL b) 2,100 mL c) 1,400 mL d) 1,050 mL e) 700 mL
40.
A sample of nitrogen gas has a volume of 32.4 L at 20°C. The gas is heated to 220ºC at constant
pressure. What is the final volume of nitrogen?
a) 2.94 L
b) 19.3 L c) 31.4 L d) 54.5 L
e) 356 L
41.
A small bubble rises from the bottom of a lake, where the temperature and pressure are 4°C and 3.0 atm,
to the water's surface, where the temperature is 25°C and the pressure is 0.95 atm. Calculate the final
volume of the bubble if its initial volume was 2.1 mL.
a) 0.72 ml b) 6.2 ml c) 41.4 ml d) 22.4 ml e) 7.1 ml
42.
Calculate the volume occupied by 35.2 g of methane gas (CH4) at 25°C and 101.3 kPa
a) 0.0186 L b) 4.5 L c) 11.2 L d) 49.2 L e) 53.7 L
43.
A gas evolved during the fermentation of sugar was collected at 22.5°C and 702 mmHg. After
purification its volume was found to be 25.0 L. How many moles of gas were collected?
a) 0.95 mol b) 1.05 mol c) 12.5 mol d) 22.4 mol e) 724 mol
44.
Determine the molar mass of chloroform gas if a sample weighing 0.389 g is collected in a flask with a
volume of 102 cm3 at 97°C. The pressure of the chloroform is 728 mmHg.
a) 187 g/mol b) 121 g/mol c) 112 g/mol d) 31.6 g/mol e) 8.28 × 10-3 g/mol
45.
A mixture of three gases has a total pressure of 1,380 mmHg at 298 K. The mixture is analyzed and is
found to contain 1.27 mol CO2, 3.04 mol CO, and 1.50 mol Ar. What is the partial pressure of Ar?
a) 0.258 atm b) 301 mmHg c) 356 mmHg d) 5 345 mmHg e) 8 020 mmHg
46.
Which of the following elements would have the lowest electron affinity?
a) Magnesium
b) Sulfur
c) Silicon
d) Phosphorus
47.
Considering the periodic table trends, which of the following statements is FALSE?
a) Sodium has a higher metallic character than Lithium.
b) Astatine is the halogen with the largest atomic radius.
c) Oxygen has a lower first ionization energy than sulfur.
d) Noble gases have very stable electron configurations.
48.
Given the isotope symbol, Br – 81, which of the following statements is FALSE?
a) this isotope represents an atom of the element bromine
b) one atom of this isotope contains 81 protons
c) one atom of this isotope contains 46 neutrons
d) one atom of this isotope contains 35 electrons
49.
Which of the following 4th row ions would have the smallest ionic radius?
a) V2+
b) K+
c) Sc3+
d) Se2e) Br-
50.
Which of the following compounds is ionic ?
a) CuSn(brass) b) NH3
c) H2O
d) KCl
51.
The electron configuration: 1s22s22p63s23p64s23d104p5 represents which element in its ground state?
a) Zinc
b) Bromine c) Chlorine d) Manganese
52.
Which of the following group 13 ( 3A) elements would have the most metallic character?
a) Thallium b) Indium c) Gallium d) Aluminum e) Boron
53.
The most common ion formed by the nitrogen atom would have a charge of:
a) +5
b) +3 c) -5 d) -3
54.
Which of the following sets would represent isotopes of the same element?
a)
c)
55.
28
A,
14
12
Q
7
29
B,
14
30
C
14
b)
16
R
8
20
S
11
d)
1
2
X,
3
Y,
Z
1
2
3
40
T
20
40
V
19
40
W
18
Use the following characteristics to identify the unknown element:
Properties:
I) it is a colorless gas at room temperatures
II) it has a very low density
III) it is very explosive
IV it has a bonding capacity of one
a) Helium
b) Oxygen
c) Hydrogen d) Nitrogen
56.
Which of the following statements about subatomic particles is False?
a) A proton and a neutron have approximately the same mass.
b) A neutron has no electrical charge.
c) The mass of the nucleus is determined primarily by the protons and electrons present.
d) The proton and electron carry opposite electrical charges.
57.
When 2.00 moles of Uranium-234 undergoes alpha decay to form thorium-230, the mass converted to
energy is 0.001156 g. The energy released by this nuclear reaction is:
a) 5.20 x 1010 J b) 1.04 X 1011 J c) 5.20 X 1013 J d) 1.02 X 1014 J
58.
Tritium nuclei are a radioactive form of hydrogen. The isotope symbol for tritium is H – 3. How many
neutrons are present in one atom of tritium?
a) 0 b) 1
c) 2
d) 3
e) 4
59.
The theoretical value for the mass of one mole of a substance is 39.23 g. Irene obtained the
following experimental data: 32.24 g, 32.26 g, and 32.25 g. Her results could be described as:
a) precise b) accurate c) both precise and accurate d) neither precise nor accurate
60.
When Cesium metal is gently heated, it quickly melts. This would be an example of a(n):
a) physical change b) chemical change c) acid-base reaction d) precision
The second part of this review are short answer questions:
1.
Balance the following equations:
a) MnO2 + KOH + O2 + Cl2
→
b) C3H7OH(l) + O2(g)
c) Al2Cl3(s)
→
→ KMnO4 + KCl + H2O
CO2(g) + H2O(g)
Al(s) + Cl2(g)
d) SO3(g) + HNO3(aq)
→ H2SO4(aq) + N2O5(g)
2.
Draw and name two isomers for:
a) C5H12
b) C4H9OH
3.
Draw the Lewis dot structures for, and predict the shape of the following:
a) phosphorous pentachloride , PCl5
b) sulfate ion, SO42c) Ozone, O3
4.
In a reaction between solid sulfur, S8, and oxygen gas, 160.0 g of sulfur dioxide is formed:
a) Write a balanced equation for this reaction
b) how many grams of sulfur were burned?
5.
Consider the reactants propene and oxygen: 2 C3H6(g) + 9 O2(g) → 3 CO2(g) + 3 H2O(g)
How many grams of oxygen are needed to react with 12.0 moles of propene?
6.
Consider the following reaction: C16H32(g) + 2 H2(g) → 2 C8H18(g)
How many grams of C8H18 can be made using 11.4 g of C16H32 ?
7.
If 35.0 g of Mg3(PO4)2 is present in a sample, determine the number of atoms of oxygen present
in the sample.
8.
Name the following organic molecule:
9.
Show the structural or condensed formula for the following:
a) 2,3 dimethyl-2-pentene b) 4-ethyl-2-octyne
10.
Given the balanced equation which represents the reaction between butane and oxygen:
2 C4H10(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(g) + 2 888 kJ
a) is this reaction endothermic or exothermic?
b) a student has 16.0 g of oxygen available to react with 40.0 g of butane. Determine the limiting
reagent and the theoretical yield of carbon dioxide gas. If the student obtains 11.0 g of carbon
dioxide, determine the percent yield of the experiment.
11.
For each single displacement below, predict whether it occurs or not, and if it does, complete and
balance it:
CH3CH2CH = C(CH3)CH2C(F)(CH3)CH3
a) Mg(s) + Pb(NO3)2(aq) →
b) Cl2(g) + NaBr(aq)
12.
→
Write a balanced equation for the reaction and identify the precipitate formed:
Ca(NO3)2(aq) + Li2SO4(aq) →
13.
What is the difference between a bond dipole and a molecular dipole? Give an example of each.
Give an example of a molecule which is non-polar, but has bonds with a measurable dipole moment.
14.
Determine the number of moles present in the following samples:
a) 4.44 g of carbon monoxide
b) 3.65 x 1027 molecules of carbon monoxide
15.
Determine the mass of the following:
a) 4.60 moles of F2O
b) 4.58 x 1024 molecules of CO2
16.
A sample contains, by mass, 40.0 g of sulfur and 60.0 g oxygen. Determine its empirical formula.
17.
A compound contains 25.8 g of potassium 42.4 g of sulfur, and 31.7 g of oxygen. Its molar mass is
known to be 453.9 g/mol. Determine its empirical and molecular formulas.
18.
Write equations to represent the dissolving of the following ionic solids in water:
a) PbCl4(s) →
b) Mg(NO3)2(s) →
19.
Complete and balance the equation for the following double displacement reactions. In each case
indicate whether a precipitate will be formed, a gas will be produced, or neutralization will occur.
a) Mg(OH)2 + H2SO4(aq) →
b) Na2CO3(aq) + H3PO4(aq) →
20.
Complete and balance the following combustion reactions:
a) C2H5OH(l) + O2(g)
→
b) CH3COCH3(l) + O2(g) →
21.
Propane burns in the presence of oxygen to form carbon dioxide and water.
a) write a balanced equation for the reaction
b) how many grams of water can be produced from 10.0 grams of propane?
c) how many grams of carbon dioxide can be produced from 20.0 grams of propane?
d) how many moles of water can be formed from 16.2 moles of oxygen gas?
e) how many grams of carbon dioxide can be produced from 2.00 moles of propane?
22.
If a student has 15.0 g of nitrogen gas and combines it with 4.00 g of hydrogen gas, how much
ammonia can be produced? N2 + 3 H2 → 2 NH3
23.
A student prepares a solution by dissolving 2.40 moles of sodium hydroxide, NaOH, in enough
distilled water to produce 250 ml of solution. Calculate the concentration of this solution in
moles/L (M)
24.
A student is required to prepare 3.00 liters of a 4.00 x 10-2M solution of hydrochloric acid, HCl.
a) how many moles of HCl are required to prepare this solution?
b) what mass of HCl is required to prepare this solution?
25.
What volume of oxygen gas at 310 K and 800 mmHg will react completely with 4.00 L of NO gas at the
same temperature and pressure?
2NO(g) + O2(g) → 2NO2(g)
.
26.
Complete the following table.
Atomic
Number
Symbol
Protons
Electrons
Neutrons
Mass
Number
____
43
_____
43
56
_____
Sn – 119
_____
50
_____
____
_____
Rb+
37
37
_____
_____
86
Se2-
_____
34
_____
45
_____
_____
_____
______
115
_____
193
Pt
78
27.
Indicate the number of significant figures for the following measurement: 7.020 litres
28.
Perform the following calculation to the correct number of significant figures:
(4.30 mole)(1.34 L) / (5.963 kPa)
29.
Name the 4th row element that has the valence electron configuration of: s2p2
30.
The ion S2- would have the same electron configuration as which noble gas?
31.
An element has a ground state electron configuration of 1s22s22p3.
a) what type of ion would this element form to become stable? _______
b) what group (family) does this element belong to? ___________
32.
Explain why a chloride ion, Cl-, (181 pm) is so much larger than a chlorine atom (99 pm).
33.
Element Y has two isotopes, Y – 101 with an abundance of 59.25% in nature, and Y – 103 with an
abundance of 41.75%. Calculate the average atomic mass of element Y.
34.
Complete the following nuclear reaction:
249
Cf
98
11
+
B

5
1
5 n
0
________________
+
__________________
35.
Write the nuclear reaction to represent the alpha decay of Gold – 199
36.
Construct Lewis diagrams and structural formulas for the following covalently bonded molecules:
identify the correct shapes.
a) H2S
37.
b) CO
c) NO3-
d) PF3
e) TeCl4
f) XeF2
g) SO2
Iodine – 132 is used in the treatment of thyroid conditions. It has a half-life of 2.33
hours. How much of an initial 69.0 gram sample would remain after 5.8245 hours?
38.
Explain the differences/similarities in the Rutherford, Bohr, and Quantum models of the atom.
39.
Draw the structural and condensed formula for the following organic compounds:
a) 2 –methylpentane
b) 2,4-dimethyl-3-ethyl-1-hexene
c) 3-ethyloctane
d) bromocyclopropane
e) 4-methyl-1-pentanol
f) 2,3-dichloro-3-bromo-1-butene
g) 5-fluoro-2-hexyne
h) 3,4,4-trimethyl-2-decene-5-yne
i) 2-chlorophenol
j) 1,2,3-trichlorobenzene
k) 1,3,5-tribromobenzene
l) 2-pentanone
m) 3-methylbutanal
n) methoxypropane
o) ethylbutanoate
p) 2-phenyl-4-chlorooctane
40.
Match the name in the left colomn with the correct structure to the right
_________ cyclohexane
A. CH3 – CH – CH2 – CH2 –CH3
|
OH
_________ 3 – pentanol
B.
_________ 1, 4 – dimethylbenzene
C. CH3 – CH2 – CH – CH2 – CH3
|
OH
CH3
_________ 2 – pentanol
D.
CH3
________ 1,2 – dimethylchyclohexane
E.
CH3
CH3
F.
CH3
CH3
41.
Give the correct IUPAC name for the following hydrocarbons:
CH3
a) CH ≡ C – CH3
b)
|
CH2
|
CH3 – CH2 – CH – CH – CH2 – CH2 – CH2 – OH
|
CH3
____________________________
________________________________________
c)
d)
Cl
CH2 – CH3
Cl
Cl
CH2 – CH3
_______________________________
___________________________________
42.
Describe the proper procedure for:
a) preparing 200.0 ml of a 0.064 M Na2CO3(aq) solution
b) preparing 460.0 ml of a 0.030 M NaI(aq) solution from a 1.0M stock solution of NaI(aq)
43.
A student dissolves 15.00 g of sodium sulfate in 500.0 ml of distilled water.
a) write an equation to represent the dissolving process
b) determine the concentration of the solution in mol/L
c) determine [Na+] in the solution
44.
A student has a 1.40 mole sample of carbon monoxide gas which occupies 11.2 L at 1.2 atm pressure.
a) What is the temperature of the gas sample in 0C ?
b) What is the density of the gas sample in g/ml ?
45.
Write valence electron configurations for the following atoms in their ground state:
a) Cu
b) Mo
c) Br
46.
Predict the most likely oxidation states for the elements in question 45 based on electron configurations
47.
An electron has the following quantum numbers:
n=3
l=1
m= 0
and
s = +1/2
what atomic orbital is the electron locataed in?
48.
A hydrogen electron is excited from the 1s orbital to the 4p orbital. What will happen next?
49.
Explain the following periodic trends: atomic radius, metallic character, and electron affinity.
50.
Draw lewis structures, structural formulae for the following molecules:
indicate the number of pi bonds present in each molecule:
CO2 , SO3, and C2HF
Good Luck!!