Chapter (1)
1.
Classify each of the following statements as a hypothesis, a law, or a theory.
(a) An autumn leaf gravitates toward the ground because there is an attractive force between the leaf
and Earth.
(b) All matter is composed of very small particles called atoms.
2.
Classify each of the following as an element, a compound, a homogeneous mixture, or a
heterogeneous mixture:
(a) water from a well
(b) argon gas
(c) sucrose
(d) chicken noodle soup
3.
The density of methanol is 0.7918 g/mL. Calculate the mass of 89.9 mL of the liquid.
1
4. Make the following conversions:
(a) 72 °F to °C
(b) 216.7 °C to °F
(c) 233 °C to K
(d) 315 K to °F
(e) 2500 °F to K
(f) 0 K to °F
5. Express the following numbers in scientific notation:
(a) 0.000000027
(b) 356
(c) 47,764
(d) 0.096.
2
6.
What is the number of significant figures in each of the following measurements?
(a) 4867 mi
(b) 56 mL
(c) 60,104 tons
(d) 2900 g
(e) 40.2 g/cm3
(f) 0.0000003 cm
(g) 0.7 min
(h) 4.6 × 1019 atoms
7. Carry out the following operations as if they were calculations of experimental results, and express
each answer in the correct units with the correct number of significant figures:
(a) 7.310 km ÷ 5.70 km
(b) (3.26 × 10−3 mg) − (7.88 × 10−5 mg)
(c) (4.02 × 106 dm) + (7.74 × 107 dm)
(d) (7.8 m − 0.34 m)/(1.15 s + 0.82 s)
3
8.
Carry out the following conversions:
(a) 70 kg to pounds.
(b) 88.6 m3 to liters.
9. How many significant figures are there in the number 0.0203610 g?
A. 8
B. 7
C. 6
D. 5
E. 4
10. Which of the following represents a chemical change?
A. Boiling water to form steam
B. Burning a piece of coal
C. heating lead until it melts
D. mixing iron powder and sand at room temperature
E. breaking glass
11. Which of the following is a homogeneous mixture?
A. Sand and water
B. Oil and water
C. Air
D. Carbon
E. Oxygen gas
4
12. A 9.4-g sample of powdered titanium was added to a graduated cylinder originally containing 6.8 mL
of water. After addition of the titanium, the water level in the graduated cylinder was 8.9 mL. What is
the density of titanium?
A. 1.4 g/cm³
B. 4.5 g/cm³
C. 1.1 g/cm³
D. 0.22 g/cm³
E. 5.4 g/cm³
13. A laboratory technician analyzed a sample three times for percent iron and got the following results:
22.43% Fe, 24.98% Fe, and 21.02% Fe. The actual percent iron in the sample was 22.81%. The
analyst's
A. Precision was poor but the average result was accurate.
B. Accuracy was poor but the precision was good.
C. Work was only qualitative.
D. Work was precise.
E. C and D.
14. Which of the following is an extensive property?
A. Density.
B. Specific heat.
C. Mass.
D. Color.
E. Melting point.
15. The answer for the following mathematical operation is:
3.254 + 4.10984 × 2.15
5.63 − 0.5
A. 2.37
B. 2.4
C. 2.371
D. 2.3706
5
Chapter (2)
1.
Indicate the number of protons, neutrons, and electrons in each of the following species:
𝟏𝟓
𝟕𝐍
2.
𝟏𝟑𝟎
𝟔𝟑
𝟖𝟒
𝟏𝟖𝟔
𝟐𝟎𝟐
, 𝟑𝟑
𝟏𝟔𝐒 , 𝟐𝟗𝐂𝐮 , 𝟑𝟖𝐒𝐫 , 𝟓𝟔𝐁𝐚 , 𝟕𝟒𝐖 , 𝟖𝟎𝐇𝐠
Give the number of protons and electrons in each of the following common ions:
K+ , Mg2+, Fe3+ , Br− , Mn2+ , C4−, Cu2+
3.
Write the formulas for the following ionic compounds:
(a) copper bromide (containing the Cu+ion)
(b) manganese oxide (containing the Mn3+ ion)
(c) mercury iodide (containing the Hg22+ion)
(d) magnesium phosphate (containing the PO43− ion).
6
4.
Which of the following pairs of atoms are isotopes of one another?
(a) 11B, 11C
(b) 55Mn, 54Mn
120
(c) 118
50Sn, 50Sn
5.
What are the empirical formulas of the following compounds?
(a) Al2Br6
(b) Na2S2O4
(c) N2O5
(d) K2Cr2O7
7
6. Name these compounds:
(a) KClO
(b) Ag2CO3
(c) FeCl2
(d) KMnO4
(e) FeO
(f) Fe2O3
(g) TiCl4
(h) Li3N
(i) Na2O
7.
Which of the following pairs of compounds cannot be used to prove the law of multiple proportions?
A. C2H6; C2H4
B. SO2; SO3
C. CO2; SO2
D. NO; NO2
E. H2O; H2O2
8
8. Which of the following pairs of compounds have the same empirical formula?
A. H2O; H2O2
B. C4H8O2; C2H4O2
C. N2O2; N2O5
D. C2H2; C6H6
E. C6H12O6; C12H22O10
9. Cathode rays are deflected away from a negatively charged plate because:
A. They are not particles
B. They are positively charged particles
C. They are neutral particles
D. They are negatively charged particles
E. They are emitted by all matter
10. In which of the following ions the number of electrons equal to the number of neutrons?
A.
27 3+
13Al
B.
16 2−
8O
C.
35 −
17Cl
D.
14 4−
6C
E.
26
2+
12Mg
11. Which formula/name pair is incorrect?
A. FeSO4 Ferrous sulfite
B. Fe2(SO3)3 Iron(III) sulfite
C. FeS Ferrous sulfide
D. FeSO3 Iron(II) sulfite
E. Fe2(SO4)3 Ferric sulfate
9
Chapter (3)
1. Naturally occurring magnesium has the following isotopic abundances:
Isotope
Abundance (%)
Atomic mass (amu)
24
Mg
78.99
23.98504
25
Mg
10.00
24.98584
26
Mg
11.01
25.98259
What is the average atomic mass of Mg?
2.
The atomic masses of 𝟔𝟑𝐋𝐢 and 𝟕𝟑𝐋𝐢 are 6.0151 amu and 7.0160 amu, respectively. Calculate the
natural abundances of these two isotopes. The average atomic mass of Li is 6.941 amu.
3.
How many atoms are there in 5.10 moles of sulfur (S)?
4.
How many moles of cobalt (Co) atoms are there in 6.00 × 109 Co atoms?
10
5.
How many grams of gold (Au) are there in 15.3 moles of Au?
6.
Dimethyl sulfoxide [(CH3)2SO], Calculate the number of C, S, H, and O atoms in 7.14 × 103 g of
dimethyl sulfoxide.
7. Calculate the percent composition by mass of (CHCl3) compound.
11
8.
What are the empirical formulas of the compounds with the following compositions?
(a) 40.1% C, 6.6 % H, 53.3 % O
(b) 18.4 % C, 21.5 % N, 60.1 % K.
9.
The empirical formula of a compound is CH. If the molar mass of this compound is about 78 g, what
is its molecular formula?
10. Balance the following equations using the method outlined in:
(a) N2O5 → N2O4+O2
(b) P4O10 + H2O → H3PO4
(c) NH3 + CuO → Cu + N2 + H2O
12
11. Silicon tetrachloride (SiCl4) can be prepared by heating Si in chlorine gas:
Si(s) + 2Cl2(g)→SiCl4(l)
In one reaction, 0.507 mole of SiCl4 is produced. How many moles of molecular chlorine were used in
the reaction?
12. Consider the reaction
MnO2+4HCl→ MnCl2+Cl2 + 2H2O
If 0.86 mole of MnO2 and 48.2 g of HCl react, which reactant will be used up first? How many grams
of Cl2 will be produced?
13
13.
Nitroglycerin (C3H5N3O9) is a powerful explosive. Its decomposition may be represented by
4C3H5N3O9→6N2+12CO2+10H2O + O2
(a) What is the maximum amount of O2 in grams that can be obtained from 2.00 × 102 g of
nitroglycerin?
(b) Calculate the percent yield in this reaction if the amount of O2 generated is found to be 6.55 g.
14. What is the total number of sulfur (S) atoms in 1.00g (Na2S4O6)?
A. 13.4 x 1021
B. 1.48 x 1021
C. 2.23 x 1021
D. 4.46 x 1021
E. 8.91 x 1021
14
N2 + 3H2 → 2NH3
15.
In the above reaction, 2.8g nitrogen (N2) is reacted with excess hydrogen (H2), and 2.55g of ammonia
(NH3) was collected as the actual yield. What is the percent (%) yield of this reaction?
A. 67%
B. 82%
C. 75%
D. 50%
E. 25%
16. Sulfur dioxide, (SO2) reacts with H2S as follows:
2H2S + SO2 → 3S + 2H2O
When 7.50g of H2S and 12.75g of SO2 are allowed to react until reaction is complete, which of the
following statements applies?
A. 0.199 mol of sulfur (S) is formed
B. 0.330 mol of sulfur (S) is formed
C. 0.0216 mol of H2S remain
D. 0.0231 mol of H2S remain
E. SO2 is the limiting reagent
15
Chapter (4)
1.
Identify each of the following substances as a strong electrolyte, weak electrolyte, or nonelectrolyte:
(a) Ba(NO3)2
(b) Ne
(c) NH3
(d) NaOH.
2. Write ionic and net ionic equations for the following reactions.
(a) Na2S(aq)+ZnCl2(aq)→
(b) Mg(NO3)2(aq)+2NaOH(aq)→
16
3. Identify each of the following species as a Brønsted acid, base, or both:
(a) PO43−
(b) ClO2−
(c) NH4+
(d) HCO3−
4.
For the complete redox reactions given here, write the half-reactions and identify the oxidizing and
reducing agents.
(a) 4Fe+3O2→2Fe2O3
(b) Cl2+2NaBr→2NaCl+Br2
(c) H2+Cl2→2HCl
17
5. Give the oxidation number of the underlined atoms in the following molecules and ions:
(a) ClF
(b) C2H2
(c) K2CrO4
(d) NaIO3
6. Predict the outcome of the reactions represented by the following equations by using the activity
series, and balance the equations.
(a) Cu(s)+HCl(aq)→
(b) I2(s)+NaBr(aq)→
(c) Mg(s)+CuSO4(aq)→
(d) Cl2(g)+KBr(aq)→
18
7. Calculate the molarity of each of the following solutions:
(a) 6.57 g of methanol (CH3OH) in 1.50 × 102 mL of solution
(b) 10.4 g of calcium chloride (CaCl2) in 2.20 × 102 mL of solution
8.
Determine how many grams of each of sulfuric acid (H2SO4) would be needed to make 2.50 ×
102 mL of a 0.100 M solution.
9. A 46.2-mL, 0.568 M calcium nitrate [Ca(NO3)2] solution is mixed with 80.5 mL of 1.396 M calcium
nitrate solution. Calculate the concentration of the final solution.
19
10. A sample of 0.6760 g of an unknown compound containing barium ions (Ba2+) is dissolved in water
and treated with an excess of Na2SO4. If the mass of the BaSO4 precipitate formed is 0.4105 g, what
is the percent by mass of Ba in the original unknown compound?
11. What volume of a 0.500 M HCl solution is needed to neutralize each of the following?
(a) 10.0 mL of a 0.300 M NaOH solution
(b) 10.0 mL of a 0.200 M Ba(OH)2 solution
20
12. The SO2 present in air is mainly responsible for the acid rain phenomenon. Its concentration can be
determined by titrating against a standard permanganate solution as follows:
5SO2+2MnO−4+2H2O→5SO42−+2Mn2++4H+
Calculate the number of grams of SO2 in a sample of air if 7.37 mL of 0.00800 M KMnO4 solution
are required for the titration.
13. The net ionic equation for the reaction
Pb(NO3)2 (aq) + 2KCl (aq) → PbCl2 (s) + 2KNO3 (aq)
would be:
A. Pb2+ (aq) + 2Cl- (aq) → PbCl2 (s)
B. K+ (aq) + NO3- (aq) → KNO3 (aq)
C. K+ (aq) + Cl- (aq) → KCl (aq)
D. Pb2+ (aq) + 2 NO3- (aq) → Pb(NO3)2 (aq)
E. None of the above
21
14. Na2CO3 (aq) + H2SO4 (aq) → NaSO4 (aq) + H2O (l) + CO2 (g)
The spectator ions in the above reaction are:
A. Na+ only
B. SO42- only
C. Na+, CO32-, and SO42D. Na+ and SO42E. CO32- and SO42-
15. Calculate the percentage of oxalic acid, H2C2O4 (diprotic acid), in an impure solid sample, given
that 0.7984 g of the sample required 37.98 mL of 0.2283 M NaOH for complete neutralization
A. 48.89%
B. 97.78%
C. 28.59%
D. 1.086%
E. 38.98%
16. What is the total concentration in mole/litre(M) of nitrate (NO3-) anion in 500.0 mL aqueous
solution containing 0.250 mol (KNO3) and 0.125 mole (Pb(NO3)2)?
A. 1.25 M
B. 1.00 M
C. 0.500 M
D. 0.250 M
E. 0.750 M
17. What is the concentration of HNO3 in the final solution when 70.0 mL of a 6.00 M HNO3 solution
is diluted with pure water to a total volume of 0.15 L?
A. 3.57 M
B. 2.80 M
C. 12.6 M
D. 1.75 M
22
18. Consider the following reactions:
AgNO3(aq) + Zn (s) → Ag (s) + Zn(NO3)2(aq)
Zn(NO3)2(aq) + Co (s) → No reaction
AgNO3(aq) + Co (s) → Co(NO3)2(aq) Ag (s)
Which is the correct order of increasing activity for these metals?
A)
Ag < Zn < Co
B)
Co < Ag < Zn
C)
Co < Zn < Ag
D)
Ag < Co < Zn
E)
Zn < Co < Ag
23
Chapter (6)
1.
A gas expands in volume from 26.7 mL to 89.3 mL at constant temperature. Calculate the work
done (in joules) if the gas expands:
(a) against a vacuum
(b) against a constant pressure of 1.5 atm
(c) against a constant pressure of 2.8 atm.
2. Determine the amount of heat (in kJ) given off when 1.26 × 104 g of NO2 are produced according to
the equation
2NO(g)+O2(g)→2NO2(g) ΔH=−114.6kJ/mol
24
3. A quantity of 85.0 mL of 0.900 M HCl is mixed with 85.0 mL of 0.900 M KOH in a constantpressure calorimeter that has a heat capacity of 325 J/°C. If the initial temperatures of both solutions
are the same at 18.24°C, what is the final temperature of the mixed solution? The heat of
neutralization is −56.2 kJ/mol. Assume the density and specific heat of the solutions are the same as
those for water.
4. At 25°C, the standard enthalpy of formation of HF(aq) is given by −320.1 kJ/mol; of OH−(aq), it is
−229.6 kJ/mol; of F−(aq), it is −329.1 kJ/mol; and of H2O(l), it is −285.8 kJ/mol.
(a) Calculate the standard enthalpy of neutralization of HF(aq):
HF(aq)+OH−(aq)→F−(aq)+H2O(l)
(b) Using the value of −56.2 kJ as the standard enthalpy change for the reaction
H+(aq)+OH−(aq)→H2O(l)
calculate the standard enthalpy change for the reaction
HF(aq)→H+(aq)+F−(aq)
25
5. From the following data,
C(graphite)+O2(g) →CO2(g)
ΔH°rxn=−393.5kJ/mol
H2(g)+1/2O2(g) → H2O(l)
ΔH°rxn=−285.8kJ/mol
2C2H6(g)+7O2(g)→4CO2(g)+6H2O(l) ΔH°rxn=−3119.06kJ/mol
calculate the enthalpy change for the reaction
2C(graphite)+3H2(g)→C2H6(g)
6. Calculate ∆U and determine whether the process is endothermic or exothermic for the following
cases:
(c) q = 0.763 kJ and w = -840 J.
(b) A system releases 66.1 kJ of heat to its surroundings while the surroundings do 44.0 kJ of work on
the system.
26
7. Based on the equation ∆U = q + w: if the reaction is exothermic and work is done by the system
on the surroundings, then:
A. q is -ve and w is -ve
B. q is -ve and w is +ve
C. q is +ve and w is +ve
D. q is +ve and w is -ve
E. none of the above
8. Consider the following reaction:
2 Mg(s) + O2(g) →2 MgO(s) ∆H = -1204 kJ
(a) Is this reaction exothermic or endothermic?
(b)Calculate the amount of heat transferred when 3.55 g of Mg(s) reacts at constant pressure.
(c) How many grams of MgO are produced during an enthalpy change of -234 kJ?
(d) How many kilojoules of heat are absorbed when 40.3 g of MgO(s) is decomposed into Mg(s) and
O2(g) at constant pressure?
27
9. Given that:
2SO3 (g) → O2 (g) + 2SO2 (g)
∆H°rxn = 188 kJ
What is the enthalpy change for the following reaction?
SO2 (g) + ½ O2 (g) →SO3 (g)
A. -94.0 kJ
B. -188.0 kJ
C. +199.0 kJ
D. -376 kJ
E. +94.0 kJ
10. A quantity of 1.056g of an organic compound (molar mass =122.12g/mol) was burned in a
constant-volume bomb calorimeter. The temperature in the calorimeter rose from 23.32 °C to
25.91°C. If the heat capacity of the calorimeter (bomb plus water) was 10.84 kJ/°C, what is the
molar heat of combustion of the organic compound?
A. 2.808 kJ/mol
B. 1.084x103 kJ/mol
C. 3.247x103 kJ/mol
D. 2.59 kJ/mol
E. 702x102 kJ/mol
28
11. A 60.0g sample of an alloy was heated to 96.00 °C and then dropped into a beaker containing
87.0g of water at temperature of 24.10 °C. the temperature of the water rose to a final
temperature of 27.63 °C. The specific heat of water is 4.184 J/g.°C. What is the specific heat of
the alloy?
A. 6.23 J/g.°C
B. 2.16 J/g.°C
C. 0.118 J/g.°C
D. 1.72 J/g.°C
E. 0.313 J/g.°C
12. The standard heat of formation (∆H°f) for potassium chloride is the enthalpy change for the
reaction:
A. K (g) + ½ Cl2 (g) → KCl (g)
B. K+(g) + Cl- (g) → KCl (s)
C. 2K (g) + Cl2 (g) → 2KCl (s)
D. K (s) + ½ Cl2 (g) → KCl (s)
E. K+(g) + Cl- (g) → KCl (l)
29
Chapter (7)
1. Find
(a) What is the frequency of light having a wavelength of 456 nm?
(b) What is the wavelength (in nm) of radiation having a frequency of 2.45 × 109 Hz?
2. A particular form of electromagnetic radiation has a frequency of 8.11 × 1014 Hz.
(a) What is its wavelength in nanometers? In meters?
(b) To what region of the electromagnetic spectrum would you assign it?
(c) What is the energy (in joules) of one quantum of this radiation?
30
3. Calculate the frequency (Hz) and wavelength (nm) of the emitted photon when an electron drops
from the n = 4 to the n = 2 level in a hydrogen atom.
4. Protons can be accelerated to speeds near that of light in particle accelerators. Estimate the
wavelength (in nm) of such a proton moving at 2.90 × 108 m/s. (Mass of a proton = 1.673 ×
10−27kg.)
5. Which orbital in each of the following pairs is lower in energy in a many-electron atom?
(a) 2s, 2p
(b) 3p, 3d
(c) 3s, 4s
(d) 4d, 5f
31
6. The ground-state electron configurations listed here are incorrect. Explain what mistakes have
been made in each and write the correct electron configurations.
Al: 1s22s22p43s23p3
B: 1s22s22p5
F: 1s22s22p6
7. Indicate the total number of:
(a) p electrons in N (Z = 7)
(b) s electrons in Si (Z = 14)
(c) 3d electrons in S (Z = 16)
8. Which of the following species has the most unpaired electrons: S+, S, or S−? Explain how you
arrive at your answer.
32
9. Which set of quantum numbers correctly describes an electron in the outermost orbital of a sulfur
atom?
A.
B.
C.
D.
E.
n= 3, l = 2, ml= –2
n= 2, l= 1, ml= –1
n= 2, l= 0, ml= 0
n= 3, l= 1, ml= –1
n= 3, l = 3, ml= –2
10. Based on the Broglie relationship, the wavelength (𝜆) of a particle weighing 1.85 x 10-28 kg and
moving at a speed of 2.31 x 106 ms-1 is expected to be:
A. 6.45 x 1011 m
B. 1.55 x 10-12 m
C. 2.79 x 1011 m
D. 2.86 x 10-11 m
E. 3.58 x 10-12 m
11. Which of the following sets of quantum numbers would be correct for an electron in an atom?
n
l
ml
ms
A.
0
1
0
+½
B.
1
1
0
+½
C.
1
0
1
-½
D.
2
1
-2
+½
E.
2
1
-1
+½
12. Which of the following represent the electronic configuration of
A. [Kr]5s2 4d10 5p4
B. [Ar]4s2 3d10 4p4
C. [Ar]4s2 3d8 4p6
D. [Kr]5s2 4d8 5p6
E. [Kr]5s2 4d9 5p5
33
52Te?
13. Which of the following group of elements have the same number of valence electrons?
A. B and Al
B. Li and Be
C. Na and Cl
D. As and Se
E. F and Ne
14. Consider the element with the electron configuration [Xe]6s2 4F7. This element is:
A. A lanthanide element
B. A nonmetal
C. An alkali earth element
D. A noble element
15. How many electrons are in the 4p subshell (sublevel) of a vanadium (V) atom?
A. 0
B. 2
C. 4
D. 5
16. Concerning the electron configuration of sulfur (S), 1s2 2s2 2p6 3s2 3p4, which of the following
represents the valence shell electrons of sulfur?
A. 1s2
B. 1s2 2s2
C. 1s2 2s2 2p6
D. 1s2 2s2 2p6 3s2
E. 3s2 3p4
34
Chapter (8)
1. Group the following electron configurations in pairs that would represent similar chemical
properties of their atoms:
(a) 1s22s22p5
(b) 1s22s1
(c) 1s22s22p6
(d) 1s22s22p63s23p5
(e) 1s22s22p63s23p64s1
(f) 1s22s22p63s23p64s23d104p6
2. Write the ground-state electron configurations of the following transition metal ions:
(a)
Sc3+
(b)
V5+
(c)
Cr3+
3. Name the ions with +3 charges that have the following electron configurations:
(a) [Ar]3d3
(b) [Ar]
(c) [Kr]4d6
(d) [Xe]4f145d6
35
4. Group the species that are isoelectronic: Be2+, F−, Fe2+, N3−, He, S2−, Co3+, Ar.
5. List the following ions in order of increasing ionic radius: N3−, Na+, F−, Mg2+, O2−.
6. Arrange the following in order of increasing first ionization energy: F, K, P, Ca, and Ne.
7. Arrange the elements in each of the following groups in increasing order of the most positive
electron affinity:
(a) Li, Na, K
(b) F, Cl, Br, I
(c) O, Si, P, Ca, Ba
8. Which oxide is more basic, MgO or BaO? Why?
36
9. Which of the following atoms has the largest radius (largest size)?
A. Cl
B. S
C. P
D. O
E. N
10. Of the following atoms, which has the lowest first ionization energy?
A. Li
B. Na
C. Cs
D. Rb
E. O
11. Which of the following elements would have the largest second-ionization energy?
A. Sc
B. Sr
C. Ca
D. K
12. Which of the following represent an isoelectronic series?
A. B5-, Si4-, As3-, Te2B. O2-, F-, Ne, Na+
C. S, Cl, Ar, K
D. Si2-, P2-, S2-, Cl2E. F-, Cl-, Br-, I-
37
13. Which of the following statements is not correct?
A. K has higher ionization energy than Na
B. Na has larger atomic radius than Cl
C. F- has larger ionic radius than Li+
D. Cl has higher electronegativity than Si
E. Cs+ has higher ionic radius than K+
14. Select the element with the most negative electron affinity.
A. H
B. Li
C. C
D. F
E. Ne
15. Which of the following compounds is expected to have the highest lattice energy?
A. NaCl
B. NaBr
C. MgF2
D. LiCl
16. Which of the following is a basic oxide?
A. P4O10
B. MgO
C. Al2O3
D. SO2
E. Cl2O7
38
Chapter (9)
1. Specify which compound in the following pairs of ionic compounds has the higher lattice energy:
(a) KCl or MgO
(b) LiF or LiBr,
(c) Mg3N2 or NaCl.
2. Classify the following bonds as ionic, polar covalent, or covalent, and give your reasons:
(a) the SiSi bond in Cl3SiSiCl3
(b) the SiCl bond in Cl3SiSiCl3
(c) the CaF bond in CaF2
(d) the NH bond in NH3
3. Write Lewis structures for the following molecules and ions:
(a) OF2
(b) N2F2
(c) OH−
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4. Draw three resonance structures for the chlorate ion, ClO3−. Show formal charges.
5. Which of the following molecules is expected to be non-polar (has zero dipole moment)?
A. H2O
B. NH3
C. CF4
D. SF2
E. HCl
6.
The correct order of electronegativity for the following elements is:
A. N > O > F
B. S > O > F
C. Li > K > Na
D. As > P > N
E. F > Cl > Br
7.
The formal charge on the N, C, and O atoms in the following Lewis structure are:
N
C
O
A.
0
+1
0
B.
0
-1
-1
C.
-1
0
+1
D.
0
0
-1
E.
-2
0
0
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8. According to Lewis, the total number of lone pair of electrons in NCl3 is:
A. 6
B. 10
C. 9
D. 8
E. 13
9.
Which of the following molecules does not obey the octet rule?
A. SbCl5
B. PCl3
C. GeH4
D. Br2
E. SeO2
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Chapter (10)
1. Predict the geometry of the following molecules and ion using the VSEPR model:
(a) CH3I
(b) ClF3
(c) H2S
(d) SO3
(e) SO42−
2. Arrange the following compounds in order of increasing dipole moment:
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3. List the following molecules in order of increasing dipole moment: H2O, CBr4, H2S, HF, NH3,
CO2.
4. What is the hybridization state of the central N atom in the azide ion, N3−?
5. How many pi bonds and sigma bonds are there in the tetracyanoethylene molecule?
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6. The molecular geometry of (BBr3) is ……………. and the molecule is………….
A. Octahedral, nonpolar
B. Trigonal pyramidal, polar
C. Trigonal bipyramidal, polar
D. Trigonal planar, nonpolar
E. Tetrahedral, nonpolar
7. What is the hybridization of (Xe) in (XeF4)?
A. sp2
B. sp3
C. sp3d
D. sp3d2
E. sp
8. The molecular structure of PCl3 has a:
A.
B.
C.
D.
Linear with one lone pair of electrons
Tetrahedral with 3σ-bonds
Square planar with sp3 hybridized nitrogen atoms
Trigonal pyramidal with one lone pair of electrons on the central atom
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Chapter (24)
1. Name the following compounds:
CH3─C≡C─CH2─CH3
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2. Classify each of the following molecules as alcohol, aldehyde, ketone, carboxylic acid, amine, or
ether:
(a) CH3─O─CH2─CH3
(b) CH3─CH2─NH2
(c) CH3─CH2─CH2─OH
(d)
(e)
(f)
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3. Which of the following compounds is aldehyde?
A. CH3COCH3
B. CH3OCH3
C. CH3CHO
D. CH3COOH
E. HC ≡ CH
4. The IUPAC systematic name of the following compound is:
A. 2-Ethyl-3-methylpentane
B. 2,3,3-Trimethylpentane
C. 2-Ethyl-2,3-dimethylbutane
D. 2-Methyl-2-isopropylbutane
E. 3-Ethyl-2,3,3-trimethylpropane
5. Alkynes have the general formula:
A. CnH2n-4
B. CnH2n-2
C. CnH2n
D. CnH2n+2
E. CnH2n+4
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6. Which of these molecules is unsaturated?
7. How many 𝜎 and π bonds are in:
A. 19𝜎, 3π
B. 10𝜎, 3π
C. 16𝜎, 2π
D. 10𝜎, 2π
E. 18𝜎, 2π
8. Which of the following compounds is cis-geometric isomer?
A)
B)
C)
D)
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