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A.P. Chemistry-Semester 1 Final
NET IONICS – TRY THESE
•
• Part A -20 scantron questions-no calculator
– Can use periodic table and pink sheet
• Part B- Short Essay – no calculator
– Can use periodic table and pink sheet
• Part C - can use calculator
Write the net ionic equation for the reactants and the products for the laboratory
situations described below. In all cases, a reaction occurs. Assume that
solutions are aqueous unless otherwise indicated. Omit formulas for any ions or
molecules that are unchanged by the reaction (net ionic). You must balance the
equations.
1) Example: A strip of magnesium is added to a solution
of silver nitrate.
Mg + 2Ag+ ---> Mg 2+ +2 Ag
(a) Solid calcium carbonate is strongly heated.
CaCO3 ---> CaO + CO2
– Can use periodic table and pink sheet
(b) A piece of nickel metal is immersed in a solution of
copper(II) sulfate.
Ni + Cu2+ ---> Ni2+ + Cu
hydrated ions acceptable with correct charge
1 point for Ni(OH)2 as product
(c) Equal volumes of equimolar solutions of sodium
hydrogen phosphate and hydrochloric acid are mixed.
HPO42¯ + H+ ---> H2PO4¯
incorrect charge on H2PO4¯ when only one product
occurs, 1 point only
1 product point for transfer if H+ from an ionic reactant
to product when a phosphate species is incorrectly but
consistently written.
(d) Chlorine gas is bubbled into a solution of
sodium bromide.
Cl2 + 2 Br- ---> 2 Cl¯ + Br2
no credit for monatomic Cl as reactant or Br as
product
(g) Drops of liquid dinitrogen trioxide are added to
distilled water.
N2O3 + H2O ---> 2 HNO2
1 product point for H+ + NO2¯
(h) Solutions of potassium permanganate and sodium
oxalate are mixed in an acidic solution.
16 H+ +2 MnO4¯ + 5C2O4 2¯ ---> 2Mn2+ + 10CO2 + 8 H2O
(e) Ammonia gas is bubbled into a solution of
ethanoic (acetic) acid.
NH3 + HC2H3O2 ---> C2H3O2¯ + NH4+
1 product point for NH4C2H3O2
1 point for NH3 + H+ ---> NH4 +
(f) Solid ammonium carbonate is added to a
saturated solution of barium hydroxide.
(NH4)2CO3+ Ba 2++ 2 OH¯ ---> 2 NH3 +BaCO3 +2H2O
1 product point for either NH3 or BaCO3
2 product points for all three species correct
*Note about barium hydroxide = marginally soluble-mostly
insoluble unless small amounts being formed or in hot water or in
an acidic environment in this case given as a solution.
1. During the gravimetric dehydration of a hydrate what factors
could cause you to have too large of a percentage of water driven
off? Too little?
This comes from the experiment we did about dehydrating a hydrate.
The percentage of water expected to be driven off comes from the
formula itself. Some factors that would cause you to have too much
are: spattering some of the original sample or excessive heating that
causes it to decompose or not waiting long enough for crucible to cool.
Too little would be caused by incomplete heating.
2.
If you assigned to only use one solvent to separate Ca2+
from Na+ and K+, which acid would you use?
Hydrochloric, Sulfuric or Nitric?
You need to choose the one that only precipitates Ca2+.
Sulfuric acid is the right choice. Know those rules!
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3.
Given a list of gases, what criteria can you use to determine
which one deviates the most from ideal gas behavior?
At very large pressures, several factors become more
important…Larger atoms deviate more, so do those atoms or
ions or molecules with very strong intermolecular forces.
4.
6. 100. mL of 0.100 M Al(NO3)3 is added to 50.0 mL of 0.35 M
Ba(NO3)2. What is the molarity of the NO31- ion in the resulting solution?
0.100L 0.100 mol Al(NO3)3 3 NO3- ions = 0.0300 mol NO3-
What volume of water must be added to dilute 140. mL of a
1.00 M NaOH solution to 0.25 M solution?
M1V1 = M2V2 (140. mL)(1.00M) = (.25M)V2 V2 = 560. mL
Remember in solutions volume is additive so (560.-140.) =
amount of water added = 420 mL
5.Equal number of moles of different gases are placed into a container with a
small hole, what criteria determines which effuses faster? How could you
determine their partial pressures when the system reaches equilibrium?
1.L
1 mol Al(NO3)3
0.0500 L 0.35 mol Ba(NO3)2 2 NO3- ions = 0.0350 mol NO31L
1 mol Ba(NO3)2
(0.0300 mol + 0.0350 mol) NO3- = 0.433 M NO3(.100 L + .0500 L)
According to Graham’s Law, the smaller the molar mass the
faster it effuses. Each gases partial pressure is directly related
to its mass/mole ratios.
7. Complete the trends, using arrows…
Period
Group
Atomic Radius
First Ionization Energy
Electronegativity
9.
What is the empirical formula for a compound that is 68.4 %
Cr and the rest is Oxygen?
68.4 g Cr 1 mol = 1.32 mol Cr 31.6 g O 1 mol =1.98 mol O
52.0 g
16.0 g
Activity of Metals
1.32 mol Cr/1.32 = 1
8.
What happens to the ionization energy of calcium when you
remove the first electron, then the second electron and then the
third electron?
Because calcium would like to lose 2 electrons to be stable we
would expect the huge jump to be between the second and the
third ionization energies.
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4
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12. If you know the molality of a solution, what would you need to
know to determine its mole fraction? Molarity? Percent by mass?
Molality = mol solute / kg solvent
You would need to
Mol fraction = mol solute / total mol know the formula for
Molarity = mol solute / L solution
% mass = g solute / total grams
both the solute and
solvent, and the density
of the solvent.
= 1.5
Empirical Formula = Cr2O3
10.
Rank the following molecules in order from least to greatest
dipole moment; NO2, H2, CH4, CHCl3.
Least = H2, CH4, NO2, CHCl3= greatest
13.
3
1
1.98 mol O/1.32
Give reasons to explain the reason that CaS has a higher
melting point than LiF.
Both are ionic, so the greater the charge, the greater the IM
forces. Since Ca2+ and S2- their IM forces are greater resulting
in a higher melting point.
14.
What types of hybridization of the carbon atom are found in
the compound CH3CHCCH2?
Number the carbons 1-4 from left to right and draw the
molecule…
1. C in CH3 = sp3
2. C in CH= sp2
3. C = sp
4. C in CH2= sp2
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15.
What mass of Calcium Chloride can be made from the reaction
of 17.9 g of Sodium Chloride with 43.0 g of Calcium Nitrate?
First write a balanced equation…
1
+ ___Ca(NO
3)2
2___NaCl
1
2
___CaCl
2 + ___NaNO3
Next Find the limiting reagent…
17.9 g NaCl 1 mol = .306 mol NaCl
16. 0.50 L of a 0.75 M aluminum sulfate solution is added to
0.10 L of a 1.20 M potassium sulfate solution. How
many moles of barium chloride are necessary to
completely precipitate all the sulfate ions?
2 Al3+ + 3 SO42-
Al2(SO4)3
(0.50 L x 0.75 mol/L) = 0.375 mol Al2(SO4)3 = 1.13 mol SO42.306 mol NaCl 1 mol Ca(NO3)2
58.5 g
43.0 g Ca(NO3)2 1 mol = .262 mol Ca(NO3)2
2 mol NaCl
= .153 mol Ca(NO3)2
K2SO4
2 K+
+ SO42-
(0.10 L x 1.20 mol/L) = 0.120 mol K2SO4 = 0.120 mol SO42-
Total moles of sulfate = (1.13 + 0.120 mol) = 1.25 mol
NaCl = L.R.
164 g
.306 mol NaCl 1 mol CaCl2 111 g CaCl2 = 17.0 g CaCl2
Ba2+ + SO42BaSO4 which means a 1:1 ratio and
therefore you will need 1.25 mol of BaCl2
2 mol NaCl 1 mol CaCl2
17.Give the set of quantum numbers for the last electron in Ag and As.
Ag = 4 d 9
n=4
As = 4 p 3
n=4
l=2
l=1
ml = +1
ms =± ½
m1 = +1
ms = + or – ½
18. What is the molecular mass of a substance where 1.33 g of the gas
at 120oC and 780. mm Hg has a volume of 950. mL?
19. A 5.00 L flask contains 4.00 mol of ammonia gas and 6.00
mol of hydrogen chloride gas. There is also enough helium to
have a partial pressure of 1.50 atm. The temperature in the
flask is 125oC.
a. What is the total pressure in the flask?
nHe = (1.50 atm)(5.00L)= 0.230 mol= 1.50 atm
(0.0821) (398 K)
n = PV/RT = (1.03 atm) (0.950 L) = 0.0303 mol
Pt = (.230 mol + 4.00 mol + 6.00 mol) 1.50 atm = 66.7 atm
.230 mol
(0.0821)( 393 K)
1.33 g / 0.0303 mol = 43.9 g/mol
b. Calculate the density in g/L of the gas mixture in the flask
0.230 mol 4.00 g = 0.920 g He
c. What is the mole fraction of each remaining
substance if 4.00 moles of gaseous ammonium
chloride is made?
Write the balanced equation (no Helium reacts Duh!)
1 mol
4.00 mol 17.0 g
.
= 68.0 g NH3
1 mol
6.00 mol 36.5 g = 219 g HCl
1 mol
(0.920 + 68.0 + 219 g ) g/ 5.00 L = 57.6 g/L
NH3(g) +HCl (g)
NH4Cl(g)
Which is the limiting reactant? 4 mol + 6 mol
4 mol??
NH3 = L.R. Now calculate the moles present AFTER the reaction
for Each Species!!!
Total moles after = 6.23 mol
Mol NH3 = 0 mol
Mol HCl = ( 6.00 – 4.00) mol = 2.00 mol
Mol NH4Cl = 4.00 mol
Mol He = 0.230 mol
XHCl = 2.00/6.23 =.321
XNH4Cl = 4.00/6.23 = .642
XHe = 0.230/6.23 = .0369
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20.
A compound contains only C, H and N. It is 58.51% C and 7.37% H by
mass. Helium effuses through a porous frit 3.20 times as fast as the compound does.
Determine the empirical and molecular formulas of this compound.
4.88 mol C/ 2.44 = 2
7.30 mol H / 2.44 = 3
Empirical formula = C2H3N efm = 41.03 g/mol
3.20 = √X
1 = √4.00
21.
Draw Lewis dot diagrams for the following species:
carbonate ion, carbon dioxide and carbon monoxide.
CO2
CO32-
CO
2.44 mol N/2.44 =1
X = 40.96 g/mol
Molecular formula = C2H3N
b. Describe the molecular shape expected for each species. Include any resonance structures.
CO32- will be trigonal planar with 3 resonance structures
CO2 is linear
CO is linear
c.
Compare the bond strength and bond length of the carbon-oxygen bonds in the three
species.
Carbon monoxide would have the strongest bond with the
shortest length due to the fact that it is a triple bond.
Carbon dioxide would be next, with a double bond.
The carbonate ion has resonance which means each bond is like
a 1 1/3 strength and a length shorter than a single bond.
22.
a.
Draw the complete Lewis dot structures for the molecules CF4 and SF4.
23. Oxygen is found in the atmosphere as diatomic gas, O2, and as ozone, O3.
a.
Draw the Lewis structures for both molecules.
b.
b.
In terms of molecular geometry, account for the fact that the CF4
molecule is nonpolar, whereas the SF4 molecule is polar.
CF4 is symmetrical with equal pull in all directions, and this is
nonpolar.
SF4 has a lone pair on the central atom and is not symmetrical,
therefore it is a polar molecule.
Use the principles of bonding and molecular structure to account for the fact
that ozone has a higher boiling point than diatomic oxygen.
Ozone is nonpolar, but does have a lone pair on the central atom,
this creates a small IM force like a dipole, whereas oxygen is
symmetrical and has less IM forces.
Ozone has a greater molar mass.
c.
Use the principles of bonding and molecular structure to account for the fact
that ozone is more soluble than diatomic oxygen in water.
d.
Explain why the two bonds in O3 are the same length and are longer than the
bond length of the bond in diatomic oxygen.
Same as letter “b”
Ozone has resonance which means each bond is like a
1 ½ strength and a length shorter than a single bond
but longer than a double bond.
24.
When NH3 gas is introduced at one end of a long tube, while HCl gas is
introduced simultaneously at the other end, a ring of white ammonium chloride is
observed to form in the tube after a few minutes. This ring is closer to the HCl end of
the tube than the NH3 end.
The molecules of gas are in constant motion so the HCl and NH3 diffuse along the
tube. Where they meet, NH4Cl(s) is formed. Since HCl has a higher molar mass, its
velocity (average) is lower, therefore, it doesn’t diffuse as fast as the NH3.
25.
Explain each of the following using your knowledge of intermolecular forces and molecular
structure.
a.
Br2 has a higher boiling point than Cl2.
Both have a dipole moment of zero, but bromine has a larger molar mass.
b.
F2 has a greater bond length than O2.
Oxygen has a double bond and fluorine has a single bond. Single bonds are
longer than double bonds.
c.
LiF has a higher melting point than NaCl.
The Coulomb’s force in LIF is larger due to the smaller atom size, therefore
a higher melting point.
d.
C4H10 has a higher boiling point than CH4.
Both are nonpolar, CH4 is symmetrical and has a smaller molar mass. Butane
has greater IM forces due to its larger mass.
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