Honors Chemistry: Spring Exam Review Sheet
Name:
True or False: Decide whether the following statements are either true or false. If the statement is false, change the underlined word
or phrase to make the statement true.
TRUE
1) The density of a silver bar stays the same even as the mass of the bar increases..
Density is an intensive property. It does not depend on the amount present.
FALSE
2) The number 6.023 x 1023 contains 4 significant figures.
When decimal is present, start on left and count all digits starting with first nonzero number.
TRUE
3) Mass is an example of an extensive property.
It depends on the amount present.
FALSE
4) The CRITICAL POINT of a substance is the temperature above which the liquid state of a substance can longer
be formed.
The triple point is the temperature and pressure where all 3 states of matter are present in an equilibrium.
FALSE
5) The decomposition of H2O to form elemental hydrogen and oxygen is an example of a CHEMICAL change.
Two new substances are formed.
TRUE
6) A solution is a homogeneous mixture of two or more substances.
FALSE
7) The change of CO2 (s) to CO2 (l) is known as MELTING.
Deposition is when a gas turns directly into a solid.
FALSE
8) The OIL DROP EXPERIMENT was used to the discovery of the charge of an electron.
The gold foil experiment led to the discovery of the nucleus.
TRUE
9) The addition of an electron to an element results in the formation of an anion.
TRUE
10) An electron was the first subatomic particle to be discovered.
It was discovered by JJ Thomson using a cathode ray tube.
FALSE
11) An alpha particle is a high-speed electron.
A beta particle is a high speed electron.
FALSE
12) Isotopes of an atom have the same number of protons, but a different number of NEUTRON.
Ions (cation or anion) have different number of electrons
TRUE
13) Oxygen is one of the seven diatomic elements.
FALSE
14) Atoms of the same element must have the same ATOMIC number.
The atomic number gives the number of protons in an atom, which defines the identity of an atom.
TRUE
15) Sulfur is an atom that has six valence electrons.
TRUE
16) Chlorine is classified as a halogen.
FALSE
17) A cathode ray is a beam of ELECTRONS.
The ray bends towards the positive electrode when passed through charged plates.
TRUE
18) Elements in the periodic table are sequentially arranged in order of increasing atomic number.
FALSE
19) An example of THE LAW OF DEFINITE PROPORTIONS is that water always consists of 89% oxygen and
11% hydrogen by mass.
The law of conservation of mass states that mass is neither created nor destroyed during a process.
FALSE
20) The EXPERIMENTAL yield of a product is the experimental mass obtained from a specified reaction.
The theoretical yield is the maximum amount of product that is possible based on calculations.
FALSE
21) Spectator ions DO NOT APPEAR in a net-ionic equation.
Only ions involved in the chemical reaction appear in the net ionic equation.
FALSE
22) Sucrose (C12H22O11) will dissociate completely when dissolved in water.
Sugar (C12H22O11) is a molecular compound which do not dissociate when dissolved in water.
TRUE
23) Nitric acid is a strong acid.
Along with HCl, HBr, HI, HClO4, and H2SO4
FALSE
24) The explosion of a hydrogen bomb is accompanied by the release of heat; thus it is an EXOTHERMIC process.
ANY process gives off heat is considered to be exothermic.
FALSE
25) Green electromagnetic radiation with a wavelength of 500 nm has an energy that is INVERSELY proportional
to its wavelength.
The energy is directly proportional with the frequency ( E = h ).
TRUE
26) According to Heisenberg’s uncertainty principle, it is not possible to measure simultaneously and accurately the
position and momentum of an electron at the same time.
FALSE
27) The ground-state electron configuration for carbon is 1s22s22p2
Carbon has 4 valence electrons on the 2nd energy level.
FALSE
28) According to the HUND’S RULE, electrons must enter degenerate orbits successively with the same spin.
Pauli’s states that no two electrons can have the same quantum number.
FALES
29) The atomic radius of fluorine is SMALLER than that of nitrogen.
Radius of atoms decrease from left to right across a period
TRUE
30) Noble gases tend to be inert because their valence shell of electrons is filled.
They have full s and p sublevels, so they don’t need to gain, lose, or share electrons
FALSE
31) The first ionization energy of nitrogen is LARGER than that of carbon.
Ionization energies increase from left to right across a period
FALSE
32) The Lewis structure for N2 includes a TRIPLE bond.
Must draw the Lewis structure (10 total electrons to distribute).
FALSE
33) The bond that forms between atoms of Ca and O in the compound CaO is IONIC.
Bond between a metal and a nonmetal.
FALSE
34) When the activation energy of a reaction is greater than the energy released when the bonds of the products are
form, the process is said to be ENDOTHERMIC.
The amount of energy put in is greater than the amount of energy given off.
FALSE
35) The compound H2CO has an AX3E2 base structure.
PCl3 has 26 valence electrons to distribute. The 3 peripheral atoms require 24 electrons, so there
must be a lone pair.
FALSE
36) The molecular geometry of NH3 is tetrahedral.
It has an AX4E base structure.
TRUE
37) ICl is a polar molecule where I2 is not.
FALSE
38) The formula C2H4O can be considered an empirical formula.
Empirical formulas must be reduced to simplest ratio
FALSE
39) To be classified as a combustion reaction, OXYGEN must be one of the reactants.
All combustion reactions have oxygen as a reactant. Common combustion reactions involve
hydrocarbons.
TRUE
40) A solution is saturated when the maximum amount of solute is dissolved into a given amount of solvent.
FALSE
41) Standard temperature and pressure (STP) for a gas is 0˚C
Standard temperature of a gas is 0C or 273 K.
TRUE
42) According to Graham’s law, you should expect NH3 (g) to effuse faster through a tiny hole than CO2 (g) at 25˚C.
Lighter gases will effuse faster than heavier gases at the same temperature.
FALSE
43) According to the kinetic theory of gases, gases have THE SAME average kinetic energies at the same temperature.
AND 1atm.
According to the kinetic molecular theory, all gases have the same average kinetic energy at the same temperature.
FALSE
44) You would expect a gas at LOW pressures to behave like an ideal gas.
Gases act most ideally at low pressures and high temperatures.
FALSE
45) The vapor pressure of a liquid will INCREASE as the temperature of the liquid increases.
The rate of evaporation increases, so there will be more gas particles present.
TRUE
46) The boiling point of a liquid is the temperature at which the vapor pressure of water equals the atmospheric pressure.
The NORMAL boiling point of a substance occurs at 760 mm Hg or 1 atm
FALSE
47) H2O is predicted to have a HIGHER normal boiling point that H2S.
H2O exhibits hydrogen bonding while H2S will have dipole-dipole forces (H is not bonded to N, O, or F).
TRUE
48) The polarizability of CCl4 should be larger than that of CH4.
Polarizability depends on the total number of electrons present in a compound or atom.
TRUE
49) A reaction with the general format AB + C--CB +A is a single replacement reaction.
FALSE
50) omit this question
TRUE
51) When the [H+] is greater than the [OH-] in a solution, the pH of the solution will be less than seven.
TRUE
52) A substance that changes the activation energy of a reaction and is not permanently changed in the process is
called a catalyst
FALSE
53) Hydrochloric acid can be classified as a polyprotic acid. It is monprotic
FALSE
54) The freezing point of a pure substance is higher than a solution made of the substance. lower
FALSE
55) The pH of a 0.25 M acetic acid solution can be found by the calculation: pH = -log[0.25]
Multiple Choice: Choose the best answer to each question.
56) When a solid substance undergoes a physical change to a liquid, which of the following is always true?
a.
b.
c.
A new substance is formed.
Heat is given off.
A gas is given off.
d. It vaporizes.
e. It melts
57) An empty container weighs 15.230 g. When filled with water (density = 1.00 g/ml), it weighs 35.920 g. When filled
with an unknown liquid to the same mark as it was filled with water, it weighs 36.261 g. What is the density of the
unknown liquid?
a. 1.02 g/ml
b.
c.
d. 1.20 g/m3
e. none of the above.
1.02 g/m3
1.20 g/ml
58) Which of the following is not an intensive property of a compound?
a.
b.
c.
d. mass
chemical formula
density
melting point temperature
e. elemental composition.
59) A pure solid is heated and it decomposes into two substances, one a liquid and the other a gas. One can conclude with
certainty that:
a.
b.
The two products are elements.
One of the two product is an element
d. the liquid is a compound and the gas is an element.
e. Both products are compounds.
c. The original solid is not an element.
60) Which of the following was not one of Dalton’s postulates for his Atomic theory?
a.
Matter is made up of small indestructible spheres called atoms.
b. Atoms have a positive dense core that is surrounded by electrons.
c.
d.
e.
Atoms combined in small whole-number ratios to form elements.
An atom can neither be created nor destroyed during a chemical or physical process
They are all postulates of Dalton’s atomic theory.
61) The development of modern atomic theory took a leap forward with the proof that an atom is mostly empty space.
Which of the following experiments was responsible for this proof?
a.
b.
Millikan’s oil drop experiment.
Cathode-ray deflection in a magnetic field
d. Separation of gases by gas chromatography
e. The Wilson cloud chamber
c. The bombardment of gold foil with alpha particles
62) What do these elements have in common: Na, K, and Cs?
i.
ii.
iii.
iv.
a. 1 and 3
Metallic
Nonmetallic
Form cations
Form anions
b. 2 and 4
(5) Alkali metals
(6) Alkaline-earth metals
(7) Halogens
c. 1, 4, and 7
d. 1, 3, and 5
e. 2, 4, and 6
63) What is the percentage mass of Al in aluminum oxide (Al2O3)?
a.
b.
26.46%
20.93%
d. 47.08%
e. 53.21%
c. 52.92%
64) What is the mass of 4.25 x 1020 molecules of H2O?
a.
0.127 g
d. 142 g
e. 4.25 x 1020 g
b. 0.0127 g
c.
1420 g
65) What is the coefficient in front of BF3 when the following equation is balanced?
BF3 + NaBH4 → NaBF4 + B2H6
a.
b.
c.
d. 4
1
2
3
e. 5
66) Give the following information:
AgBr (s) + ½ Cl2 (g) → AgCl (s) + ½ Br2 (l) ΔH = -27.6 kJ
Which of the following statements concerning the reaction are true: (1) heat is released; (2) heat is absorbed; (3) the reaction
is exothermic; (4) the reaction is endothermic?
a. 1 and 3
b.
c.
d. 2 and 3
e. 1 and 4
2 and 4
1 only
67) 0.50 mol of ammonium nitrate is added to a 50.0 ml of water in a thermally insulated reaction vessel. Both are initially
20.0˚C. After stirring the temperature of the solution is found to be less than 10.0˚C. Which statement best explains the
temperature change? (Temp goes down, which means the process is endothermic)
a.
b.
Heat is evolved from the system to the surroundings.
Heat is absorbed by the surroundings from the system.
c. Heat is absorbed by ammonium nitrate when it dissolves and becomes hydrated.
d.
e.
Heat is evolved by ammonium nitrate when it dissolves and becomes hydrated.
The heat energy of the hydrated ions is lass than the ions in solid ammonium nitrate.
68) What is the energy of a photon that has a frequency of 9.00 x 10 11 1/s?
a.
1.66 x 10-45 J
b. 5.97 x
c.
10-22
J
d. 5.00 x 10-22 J
e. 3.32 x 10-45 J
4.99 x 10-27 J
69) According to Bohr’s model of the atom, emission of electromagnetic radiation by heated atoms in a vacuum is directly
due to which of the following?
a.
b.
c.
d.
Photons absorbed by the atom
Particle emission from the nucleus
Momentum possessed by electrons
Electrons being excited from a lower to higher energy level
e. Electrons falling back from an excited state to their ground state.
70) Which of the following quantum numbers for an electron is not permissible?
a.
n = 5, l = 2, ml = 0
b. n = 3, l = 2, ml = 3
c.
n = 4, l = 3, ml = -2
d. n = 1, l = 0, ml = 0
e. n = 2, l = 1, ml = -1
71) How many electrons are there in the outer energy level of arsenic?
a.
b.
3
4
d. 6
e. 7
c. 5
72) Which elements have the lowest ionization energies?
a.
b.
Those on the right side of the periodic table
Nonmetal
d. Halogens
e. Both (a) and (b)
c. Those on the left side of the periodic table
73) Why doesn’t sodium ordinarily occur in the 2 + ion state?
a.
low electron affinity
b. high second ionization energy
c.
d. small density
e. large electronegativity
large atomic radius
74) Based on its position on the periodic table, which atom has the largest atomic radius?
a.
b.
c.
Li
Na
K
d. Rb
e. Cs
75) Which element is considered a transition metal?
a.
b.
c.
Na
Ca
Cl
d. Cu
e. Al
76) Which of the following is an exception to the general ionization energy trend?
a.
b.
Na
Ca
d. Ar
e. P
c. B
77) Which of the following is the largest in size?
a.
b.
c.
ClCl
Na+
d. I+
e. I-
78) What is the electron configuration for Se-2?
a.
b.
c.
[He]2s22p4
[He]2s22p6
[Ar]4s23d104p4
d. [Ar]4s23d104p6
e. [Ne]3s23p2
79) What is the correct order of increasing radii for the isoelectronic series Rb +, Sr+2, Se-2, Br-?
a.
b.
c.
Rb+ < Sr+2 < Se-2 < BrBr- < Se-2 < Sr+2 < Rb+
Se-2 < Br- < Rb+ < Sr+2
d. Sr+2 < Rb+< Br- < Se-2
e. None of the above
80) Which of the following compounds contains a double bond in its Lewis dot structure?
a.
N2
b. SO2
c.
Cl2
d. NH4+
e. SO4-2
81) Which of the following possess π electrons?
a. CH2Cl2
d. (a) and (b)
b. OCNe. (b) and (c)
c. N2O
82) In which of the following would you expect the central atom to use sp 3d2 hybridization?
a.
PF5
d. SiO2
e. SO3
b. BrF5
c. CO2
83) A 15.0 L gas sample is heated from 29˚C to 67˚C at constant pressure conditions. Under these conditions, which of the
following statements is correct?
a.
b.
The number of moles of gas increases.
The number of moles of gas decreases.
d. The volume decreases.
e. None of the above is correct.
c. The volume increases.
84) A sample of CO2 (g) occupies 3.0 L at 35˚C at 1.0 atm. What is the new volume if the temperature and pressure are
changed to 48˚C and 1.5 atm?
a.
b.
c.
3.0 L
1.9 L
4.7 L
d. 4.3 L
e. 2.1 L
85) Which of the following substances should have the highest melting point?
a.
b.
Kr
SO3
d. CHCl3
e. H2O
c. Ca(NO3)2
86) Which substance should exhibit hydrogen bonding in the liquid phase?
a. SF4
d. HBr
b. CH3CH3
e. CH3CH2NH2
c. NaH
87) To an equilibrium mixture of 2 SO2 (g) + O2 (g)  2 SO3 (g), some helium, an inert gas, is added at constant volume.
The addition of He causes the total pressure to double. Which of the following is true?
a.
b.
c.
d. [SO3] increases.
The concentrations of all three gases are unchanged.
The number of moles of SO3 increases.
The number of moles of O2 increases.
e. [SO2] increases.
88) For a certain diprotic acid,H2X, Ka1 = 1 x 10−2 and Ka2 = 1 x 10−6. When the salt NaHX is dissolved in water, what is true
about the pH of the resulting solution?
a.
b.
pH < 7
c. pH = 7
d. the pH can’t be determined
pH > 7
89) In general, as temperature goes up, the rate of a chemical reaction
a.
b.
goes up if the reaction is exothermic
goes up if the reaction is endothermic
c. goes up regardless of whether the reaction is endothermic or exothermic.
d.
e.
stays the same regardless of whether the reaction is endothermic or exothermic.
stays the same if the reaction is first order
90) A Bronsted/Lowry acid is defined as a substance that
a.
b.
increases [H+] when placed in water
decreases [H+] when placed in water
c. acts as a proton acceptor
d. acts as a proton donor
91) Which of the following is the conjugate base of acetic acid?
a. C2H3O2b.
c.
H2C2H3O2+
H3O+
d. OHe. HC2H3O2+
Free Response: Answer each of the following questions completely.
1.
Determine the correct number of significant digits in the following numbers.
a.
b.
c.
2.
2
4
2
d. 49600.0 g
e. 200 rattlesnakes
6
infinite
Make the following temperature conversions.
a.
3.
0.00067 ml
210.0 m
0.040 L
-70º F to ºC
–56.7C
b. 480ºC to K
753 K
c. 15ºF to K
264 K
What is Dalton’s atomic theory? Which postulates are no longer true? Why?
I.
II.
III.
IV.
All matter is composed of tiny, indivisible particles called atoms.
All atoms of a given element have identical properties that different from those of other elements.
Atoms cannot be created, destroyed, or transformed into atoms of another element during a chemical
reaction.
Compounds are formed when atoms of different elements combined with one another in small wholenumber ratios.
Postulates I and II are no longer true. Postulate I is false, because of the existence of subatomic particles (protons,
neutrons, and electrons). Postulate II is false, because not all atoms of the same elements are identical due to isotopes
(a different number of neutrons) and ions (a different number of electrons).
4.
A 37Cl- atom of chlorine contains how many protons, electrons, and neutrons?
protons = 17, neutrons = 20, and electrons = 18
5.
Name the following compounds.
a.
b.
c.
d.
CaSO4
PF5
KBr
Na2S
Calcium sulfate
Phosphorus pentafluoride
Potassium bromide
Sodium sulfide
e. H2SO4 (aq)
f. NaClO3
g. Cu(CN)2
Sulfuric acid
Sodium chlorate
Copper (II) cyanide
Write the formula for the following compounds.
e.
f.
g.
h.
6.
Tin (IV) chloride
Chromium (III) hydroxide
Cesium cyanide
Dinitrogen trioxide
SnCl4
Cr(OH)3
CsCN
N2O3
e.
f.
g.
h.
calcium phosphate
osmium tetraoxide
sulfurous acid
Hydroselenic acid
Balance the following reactions, and then write the net ionic equation for each.
a.
AgNO3 (aq)+ CaCl2 (aq) → AgCl (s) + Ca(NO3)2 (aq)
2 AgNO3 (aq) + CaCl2 (aq)  2 AgCl (s) + Ca(NO3)2 (aq)
NET: Ag+ (aq) + Cl- (aq)  AgCl (s)
b.
VO (s) + Fe2O3 (s) → FeO (s) + V2O5 (s)
2 VO (s) + 3 Fe2O3 (s)  6 FeO (s) + V2O5 (s)
NET: Same as above.
c.
Na (s) + H2O (l) → NaOH (aq) + H2 (g)
2 Na (s) + 2 H2O (l)  2 NaOH (aq) + H2 (g)
Ca3(PO4)2
OsO4
H2SO3
H2Se
NET: 2 Na (s) + 2 H2O (l)  2 Na+ (aq) + 2 OH- (aq) + H2 (g)
d.
NH4NO3 (s) → N2O (g) + H2O (l)
NH4NO3 (s)  N2O (g) + 2 H2O (l)
NET: Same as above.
e.
MnO2 (s) + HCl (aq) → Cl2 (g) + MnCl2 (aq) + H2O (l)
MnO2 (s) + 4 HCl (aq)  Cl2 (g) + MnCl2 (aq) + 2 H2O (l)
NET: MnO2 (s) + 4 H+ (aq) + 2 Cl- (aq)  Cl2 (g) + Mn+2 (aq) + 2 H2O (l)
7.
Write the balanced chemical equations for the following reactions, and determine the type of reaction.
a.
Aqueous silver (I) nitrate reacts with aqueous copper (II) chloride to form insoluble silver (I) chloride and aqueous
copper (II) nitrate.
2 AgNO3 (aq) + CuCl2 (aq)  2 AgCl (s) + Cu(NO3)2 (aq); Double displacement.
b.
Aluminum metal reacts with oxygen gas to form solid aluminum oxide.
4 Al (s) + 3 O2 (g)  2 Al2O3 (s); Synthesis
c.
Aqueous barium chloride reacts with aqueous potassium sulfate to form solid barium sulfate and aqueous potassium
chloride.
BaCl2 (aq) + K2SO4 (aq)  BaSO4 (s) + 2 KCl (aq); Double displacement.
d.
Solid potassium chlorate decomposes to solid potassium chloride and oxygen gas.
2 KClO3 (s)  2 KCl (s) + 3 O2 (g); Decomposition.
8.
Ethylene, C2H4, is used to make plastic polyethylene. How many moles of ethylene are there in 5.50 g?
0.196 mol C2H4
9.
(molar mass = 28.06 g/mol)
How many atoms of oxygen are found in a 19.25 g sample of gallium nitrate?
4.08 x 1023
(Ga(NO3)3)
(molar mass = 255.74 g/mol)
(9 mol O / 1 mol Ga(NO3)3)
10. Cholesterol is made up of 83.87% C, 11.99% H, and 4.14% O. It has a molar mass of 386.64 g/mol. What is the molecular
formula of cholesterol?
The molecular formula for cholesterol is C27H46O
(Remember, use mole ratios to find empirical formula, then molar
mass / empirical mass to find ratio for molecular formula)
11. Given the following balanced reaction:
5 H2C2O4 + 2 KMnO4 → 4 H2O + 10 CO2 + 2 MnO2 + 2 KOH
a.
How many grams of CO2 are formed when 10.05 g of H2C2O4 and 26.72 g of KMnO4 are mixed together?
9.82 g CO2 (H2C2O is the limiting reagent) USE Stoichiometry!
b.
If the reaction produced 36.5 g of KOH, how many grams of H 2O were also formed?
23.4 g H2O is formed.
12. Determine the following for a sucrose solution.
a.
The number of moles of sucrose in 200.0 ml of a 0.345 M solution.
0.0690 moles sucrose
b.
(remember to convert volume to liters)
The molarity of a solution containing 10.0 g of sucrose in enough water to make a 500.0 ml solution.
0.0584 M
(sucrose is C12H22O11; molarity is moles solute / liter solution)
13. Using the solubility rules, determine if the following compounds would be soluble or insoluble in water to a large extent.
a.
b.
c.
d.
PbBr2
CsCl
Cu(C2H3O2)2
Mn(OH)2
IS
S
S
S
e. ZnS
f. Ag2SO4
g. Na2O
h. Ca3(PO4)2
IS
IS
S
IS
14. 10.30 ml of 0.150 M HCl is reacted with 11.25 ml of 0.1355 M NaOH. What is the limiting reactant?
NaOH is the limiting reactant. (Remember to start with a balanced equation, and then USE
Stoichiometry.)
15. What is the molarity of a solution of HC2H3O2 if 50.00 ml of the HC2H3O2 is reacted completely with 25.86 ml of 0.1201 M
NaOH?
0.0621 M HC2H3O2 (Again. Balanced equation, then stoichiometry!)
16. Calculate the energy of a photon with a wavelength of 420 nm.
The energy of the wavelength is 4.73 x 10-19 J (c = λ ν; E = h ν) Must convert nm to m
17. Calculate the wavelength of a photon of light that is given off when an electron in a hydrogen atom drops from the 4 th energy
level to the 2nd energy level.
The wavelength is 4.86 x 10-7 m or 486 nm. (E = RH(1/n2lo – 1/n2hi); c = λ ν)
18. List three properties that distinguish metals from nonmetals.
Metals are shiny, ductile, malleable, and conduct electricity. Nonmetals are dull, brittle,
and do not conduct electricity.
19. For the following compounds, I) drawing the Lewis dot structure, II) determine the molecular geometry, III) Determine the
hydridization of the central atom (sp3, ect.), and IV) Determine the polarity.
NOTE: unfortunately I can’t put the lewis dot structures on this sheet. See me if you are having problems.
a. CS2
linear;
sp;
nonpolar
e. SO3 Trigonal planar;
sp2;
nonpolar
b. ICl
linear;
N/A;
polar
f. TeCl4 See-saw;
sp3d; polar
c. ClF3
T-shaped;
sp3d; polar
g. PF5 Trigonal bipyramid
sp3d; nonpolar.
3 2
d. SF6
Octahedral;
sp d ; nonpolar
20. A mixture of 2.00 g of O2 (g) and 3.25 g SO2 (g) will exert what pressure on the inside of a 2.00 L container at 27˚C?
The pressure on the container is 1.39 atm. ( Ptot = ntot R T / V)
21. What volume of O2 (g) is given off when 3.25 g of KNO3 (s) decomposes to KNO2 (s) and O2 (g) at STP?
The volume of O2 is 360 ml. (Balance the reaction, and then stoichiometry and gas law)
22. How many times faster does N2 effuse than O2 at 25˚C
N2 effuses 1.07 times faster than O2. (Graham’s Law of Effusion)
23. Calculate the concentrations of each of the following solutions:
a.
the molality of 142 g of Na2CO3 in 2.00 kg of water.
The molality is 0.670 m ( molality = mole solute / Kg solvent)
b.
the mole fraction of 32.2 g of NaCl in 265.0 g of H 2O.
The mole fraction is 0.0361 (mole fraction = mole part / total # of moles)
24. A solution is 5.85 % acetic acid by mass. The solution is also known to have a density of 1.150 g/ml. What is the molarity
of the acetic acid solution?
5.85 grams acetic acid
__________________
100 grams solution
X
1.150 grams solution
_________________ X
0.001 ml solution
1 mol acetic acid
______________
= 1.12 Molar acetic acid
60.05 grams acetic acid
25. Caffeine is made up of 49.5% carbon, 5.2% hydrogen, 16.5 % oxygen, and 28.9% nitrogen. A solution made up of 8.25 g of
caffeine in 100.0 ml benzene (d = 0.877 g/ml) freezes at 3.03˚C. Pure benzene (k f = 5.10˚C/m) freezes at 5.50˚C. What is
the molecular formula of caffeine?
Empirical formula = C4H5ON2;
MM = 194.36 g/mol; Molecular formula = C8H10O2N4
(Mole ratio to find empirical. Use ΔTf = kf m i to find molality, then use molality and Kg
solvent to find moles solute. Divide grams solute by moles solute to find molar mass.
Molar mass / empirical mass = ratio to get molecular formula)
26. Name the following and calculate the boiling point and freezing point of the following solutions.
a.
b.
0.3 Molar Mg(OH)2
i=3
c. 0.6 Molar K3PO4
Magnesium Hydroxide
potassium phosphate
FP= -1.62 C
BP= 100.46 C
FP= -4.32 C
BP= 101.22 C
0.8 Molar NH4NO3
i=2
FP= -2.88 C
BP= 100.82 C
i=4
d. 1.3 Molar Glucose (C6H12O6)
i=1
FP= -2.34.C
BP= 100.66 C
ammonium nitrate
glucose
27. Determine the pH of the following solutions.
a.
0.38 M HNO3 (aq)
Strong acid: pH = 0.42
d. 0.26 M HIO
Weak Acid; pH = 5.61
b.
0.65 M Mg(OH)2
Strong Base; pH = 14.11
e. 0.15 M C6H5NH2
c.
0.75 M Na2CO3
Salt (weak base); Kb = 2.13 x 10-4; pH = 12.10
Weak Base; pH = 9.02