Chemistry L3 Final Exam Review
1. How many valence electrons are present in the following groups?
Group 1
1 valence electron
Group 2
2 valence electrons
Group 16
6 valence electrons
Group 17
7 valence electrons
Group 18
8 valence electrons
2. Which electrons are involved in chemical bonding?
Valence electrons found in the outer energy level
3. Describe the process, and particles involved in the formation of:
A) ionic bonds
Metal atoms lose electrons to nonmetal atoms. This gives the metal atoms a positive and the
nonmetal atoms a negative charge. The oppositely charged ions attract forming the bond.
B) covalent bonds
2 nonmetal atoms share electrons to complete their outer octets. This draws the atoms
together so their outer energy levels overlap.
C) metallic bonds
Metals have weak holds on their outer electrons. The atoms in a metal freely exchange these
electrons creating a sea of electrons shared by the metal atoms.
4. Describe diatomic molecules including bond type.
Diatomic molecules are made up of 2 atoms of the same element held together by a covalent
bond.
5. What elements can form diatomic molecules?
Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine Iodine
6. How do electronegativity and ionization energy affect the formation of:
A) ionic bonds
Metals have low Ionization energies and low electronegativities which allows them to lose
electrons easily. Nonmetals have high Ionization energies and high electronegativities
allowing them to pick up electrons. If the difference in electronegativity is greater than 1.7
there is a large enough difference for the metal to lose and the nonmetal to gain electrons.
B) covalent bonds
If the electronegativity difference is less than 1.7 there is not enough of a difference for a full
electron transfer causing the electrons to be shared by the atoms.
7. Define:
Cations
Positive ions
Anions
Negative ions
8. Describe the difference between polar and non-polar bonds.
If the electronegativity difference is less than 0.3 there is no significant difference in
attraction and the electrons are shared equally resulting in nonpolar bonds.
If the electronegativity difference is 0.3 to 1.7 than the electrons are shared unevenly resulting
in electrons orbiting the higher electronegative atom more often giving that atom a partial
negative charge and the other atom a partial positive charge.
9. What determines where electrons will orbit most frequently in a polar bond?
Electrons will orbit the atom with the higher electronegativity more frequently.
10. State the octet rule.
Atoms will gain, lose or share electrons to achieve 8 electrons in their outer energy level.
11. What determines the amount (degree) of polarity for a:
A) bond
The difference in electronegativity; the higher the difference the greater the polarity.
B) molecule
The presence and strength of polar bonds along with the shape of the molecule.
12. What is the assumption/principle that the VSEPR is based upon?
In a covalent bond electrons occur in pairs. These pairs of electrons will repel each other and
spread out to maximize the distance between pairs.
13. Describe the following intermolecular forces, (including the strength of attraction for each).
A) dipole-dipole forces
Forces of attraction between opposite charges on adjacent polar molecules. (Middle strength)
B) Hydrogen bonds
Strong dipole-dipole force between Hydrogen and a highly electronegative atom to which it
is not directly bonded (Strongest)
C) London Dispersion Forces
Weak short-lived attraction between momentary dipoles that are the result of random electron
shifts. (Weakest)
14. For the following compounds:
a) draw a Lewis Structure
b) for the central atom determine the number of:
1) electron pairs
2) shared pairs
3) unshared pairs
c) VSEPR shape
d) bond polarity
e) molecular polarity
CH3F
4 pairs
4 shared
0 unshared
Tetrahedral
SiO2
2 pairs
2 shared
0 unshared
Linear
No Molecular Polarity
NH3
4 pairs
3 shared
1 unshared
Pyramidal
SO324 pairs
3 shared
1 unshared
Pyramidal
NH41+
4 pairs
4 shared
0 unshared
Tetrahedral
No Molecular Polarity
15. What type of intermolecular forces act between water molecules?
Hydrogen Bonds
Explain how/why these forces occur in water.
Water is highly polar and Oxygen is a strongly electronegative atom so it attracts the
Hydrogen atoms on nearby molecules.
Describe how the intermolecular forces influence the properties of water.
The hydrogen bonds in water result in a higher the Melting and Boiling Point, makes the
water a good solvent and forms crystals in a solid state.
16. Draw and identify the major VSEPR shapes.
Linear
Pyramidal
Bent
Trigonal Planar
Tetrahedral
17. Molar mass is a general term. Define this term and list the specific types of molar mass.
Molar Mass – mass of 1 mole of a substance, includes formula mass, atomic mass, empirical
mass and molecular mass
18. What is the affect of changing the:
A) subscripts in a formula
This changes the formula describing a different substance, which may or may not exist.
B) coefficient preceding a formula
This changes the number of moles of the substance, but not its identity.
C) How many atoms of each element are in each compound?
NH3
1 Nitrogen, 3 Hydrogen
Zn(NO3)2
1 Zinc, 2 Nitrogen, 6 Oxygen
2 CCl4
2 Carbon, 8 Chlorine
3 AgNO3
3 Silver, 1 Nitrogen, 9 Oxygen
2 Fe2(SO4)3 4 Iron, 6 Sulfur, 24 Oxygen
19. Define:
Empirical Formula
Smallest whole number ratio for the atoms in a compound
Molecular Formula
Actual number of atoms, of each element, in one mole of a substance
20. Why can the empirical formula and molecular formula of a molecular compound differ?
The empirical formula gives the smallest ratio while molecular formula gives the actual
numbers. The molecular formula is a whole number multiple of the empirical formula.
21. How do we determine the molecular formula, given the empirical formula, molecular
mass and formula mass?
Find the multiple between the empirical and molecular mass by using 𝑿 =
Use X to multiply the Empirical formula to find the Molecular Formula.
[Molecular Formula = X(Empirical formula)]
𝑴𝒐𝒍𝒆𝒄𝒖𝒍𝒂𝒓 𝒎𝒂𝒔𝒔
𝑬𝒎𝒑𝒊𝒓𝒊𝒄𝒂𝒍 𝒎𝒂𝒔𝒔
23. A compound contains 40.00 % C, 6.71 % H and the remaining mass is oxygen. Based on
this information, solve the following problems:
A) What is the percentage of oxygen present? 100 – 40.00 – 6.71 = 53.29% Oxygen
B) Calculate the empirical formula of this compound.
𝟏 𝐦𝐨𝐥
𝐂 = 𝟒𝟎. 𝟎𝟎𝐠 𝐱
= 𝟑. 𝟑𝟑 𝐦𝐨𝐥
𝟏𝟐. 𝟎𝟏𝟎𝟕𝐠
𝐇 = 𝟔. 𝟕𝟏𝐠 𝐱
𝟏 𝐦𝐨𝐥
= 𝟔. 𝟔𝟔 𝐦𝐨𝐥
𝟏. 𝟎𝟎𝟕𝟕𝟗𝐠
𝐎 = 𝟓𝟑. 𝟐𝟗𝐠 𝐱
𝟏 𝐦𝐨𝐥
= 𝟑. 𝟑𝟑 𝐦𝐨𝐥
𝟏𝟓. 𝟗𝟗𝟗𝟒𝐠
C:H:O
𝟑. 𝟑𝟑 𝟔. 𝟔𝟔 𝟑. 𝟑𝟑
∶
∶
𝟑. 𝟑𝟑 𝟑. 𝟑𝟑 𝟑. 𝟑𝟑
Empirical Formula = CH2O
C) Calculate the empirical mass for this compound.
CH2O = 12.0107 + 2(1.0079) + 15.9994 = 30.0259 g/mol
D) If the molecular mass is 120.1048, what is the molecular formula?
𝐠
𝟏𝟐𝟎. 𝟏𝟎𝟒𝟖
𝐦𝐨𝐥
𝐗=
𝟒(𝐂𝐇𝟐 𝐎) = 𝐂𝟒 𝐇𝟖 𝐎𝟒
𝐠 =𝟒
𝟑𝟎. 𝟎𝟐𝟓𝟗
𝐦𝐨𝐥
21. Write formulas for the following compounds.
Cadmium chloride
CdCl2
tin (IV) oxide
SnO2
Cobalt (II) bromide
CoBr2
Nickel (III) selenide
Ni2Se3
Silver nitrite
AgNO2
Lead (IV) bisulfate
Pb(HSO4)4
Ammonium carbonate
(NH4)2CO3
Carbon monoxide
CO
Pentachloride hexafluoride
Cl5F6
tribromine heptachloride Br3Cl7
diNitrogen trioxide
N2O3
pentaNitrogen hept-oxide N5O7
22. Name the following compounds (use roman numerals or prefixes where appropriate).
CaO
Li3N
Fe2O3
CoO
N2O4
CO
Calcium oxide
Lithium nitride
Iron (III) oxide
Cobalt (II) oxide
diNitrogen tetraoxide
Carbon monoxide
AlCl3
PtS2
Tb3P4
Co2O3
CO2
P4O10
Aluminum chloride
Platinum (IV) sulfide
Terbium (IV) phosphide
Cobalt (III) oxide
Carbon dioxide
tetraPhosphorus deca-oxide
23. Answer the following questions about the compound Ni2(Cr2O7)3
A) How many chromium atoms are represented by this formula?
6
B) What type of bond holds the Cr2O7 -2 group together?
Covalent
C) What type of Ion is the Cr2O7 -2 group?
Polyatomic ion
D) What type of bond attaches the Nickel to the dichromate group? Ionic
E) What is the formula mass of this compound?
2(58.6934) + 6(51.9961) + 21(15.9994) = 765.3508 g/mol
F) How many moles of Nickel would be in a 5 mol sample of this compound?
G) How many atoms of Nickel would be in a 5 mole sample of this compound?
10 mol (6.02x 1023atoms/mol) = 6.02 x 1022 atoms
H) Determine the percentage composition for this compound.
𝐍𝐢 =
𝐂𝐫 =
𝐎=
𝐠
)
𝐦𝐨𝐥 𝐱 𝟏𝟎𝟎 = 𝟏𝟓. 𝟑𝟑𝟕𝟔 %
𝐠
𝟕𝟔𝟓. 𝟓𝟎𝟖
𝐦𝐨𝐥
𝟐(𝟓𝟖. 𝟔𝟗𝟑𝟒
𝐠
)
𝐦𝐨𝐥 𝐱 𝟏𝟎𝟎 = 𝟒𝟎. 𝟕𝟔𝟐𝟔 %
𝐠
𝟕𝟔𝟓. 𝟓𝟎𝟖
𝐦𝐨𝐥
𝟔(𝟓𝟏. 𝟗𝟗𝟔𝟏
𝐠
)
𝐦𝐨𝐥 𝐱 𝟏𝟎𝟎 = 𝟒𝟑. 𝟖𝟗𝟗𝟖 %
𝐠
𝟕𝟔𝟓. 𝟓𝟎𝟖
𝐦𝐨𝐥
𝟐𝟏(𝟏𝟓. 𝟗𝟗𝟗𝟒
10
24. Which of the following are polyatomic ions?
SO4-2
Cl-1
C2O4-2
NO3-1
25. What is represented by a chemical equation?
Ca+2
A chemical reaction
26. What is the correct order in which to write a chemical equation?
(use , products and reactants)
Reactant  Product
27. Why must a chemical equation be balanced?
So that it will obey the Law of Conservation of Mass.
28. What are the only numbers you are allowed to change when balancing a chemical equation?
Coefficient
29. A 250.0 g sample of a carbon hydrogen compound is found to contain 224.8 g of Carbon.
A) Find the amount of hydrogen in this sample.
250.0 g – 224.8 g = 25.3 g Hydrogen
B) Determine the percentage composition of this compound.
𝐂=
𝟐𝟐𝟒. 𝟖 𝐠
𝐱 𝟏𝟎𝟎 = 𝟖𝟗. 𝟗𝟐 % 𝐂
𝟐𝟓𝟎. 𝟎 𝐠
𝐇=
𝟐𝟓. 𝟐 𝐠
𝐱 𝟏𝟎𝟎 = 𝟏𝟎. 𝟎𝟖 % 𝐇
𝟐𝟓𝟎. 𝟎 𝐠
C) How many moles of each element are present in the sample?
𝐂 = 𝟐𝟐𝟒. 𝟖 𝐠 𝐱
𝐇 = 𝟐𝟓. 𝟐 𝐠 𝐱
𝟏 𝐦𝐨𝐥
= 𝟏𝟖. 𝟕𝟐 𝐦𝐨𝐥 𝐂
𝟏𝟐. 𝟎𝟏𝟎𝟕 𝐠
𝟏 𝐦𝐨𝐥
= 𝟐𝟓. 𝟎 𝐦𝐨𝐥 𝐇
𝟏. 𝟎𝟎𝟕𝟗 𝐠
D) What is this compound’s empirical formula?
C:H
𝟏𝟖.𝟕𝟐
𝟏𝟖.𝟕𝟐
∶
𝟐𝟓.𝟎
𝟏𝟖.𝟕𝟐
= 𝟏 ∶ 𝟏. 𝟑𝟑 = 𝟑 ∶ 𝟒
C3H4
E) Calculate the formula mass of this compound.
3(12.0107) + 4(1.0079) = 40.0637 g/mol
F) The molecular mass of the compound is 120.21 g/mol, what is its molecular formula?
𝐗=
𝐠
𝐦𝐨𝐥
𝐠
𝟒𝟎.𝟎𝟔𝟑𝟕
𝐦𝐨𝐥
𝟏𝟐𝟎.𝟐𝟏
=𝟑
3(C3H4) = C9H12
30. Balance and classify the following reactions:
Ca
+
2HCl

CaCl2

2 Fe2O3
4 Fe
+
3 O2
3 Fe
+
4 H2O

FeS
+
2 HCl

H2S
5 O2

3 CO2
C3H8
2 KI
+
+
2 HgO
2 KF

2 Hg
+
H2
+
+
+
Single Replacement
Synthesis
Fe3O4

F2
+
+
4 H2
Single Replacement
FeCl2
Double Replacement
4 H2O
Combustion
I2
O2
Single Replacement
Decomposition
31. For the following pairs of elements determine which would be able to replace the other in a
single replacement reaction.
A) Fluorine and Bromine
B) Silver and Aluminum
C) Hydrogen and Copper
D) Calcium and Zinc
32. Indicate which of the following would precipitate (PPT) or would remain dissolved (aq) if
produced in a reaction. For the soluble compounds show the dissociation of its ions.
A) Sc(C2H3O2)3
(aq)
B) PbI2
PPT
Sc3+ + 3 C2H3O21-
C) Cr(CO3)2
PPT
D) H2SO4
(aq)
33
2 C8H18
+
2 H1+ + SO42-
25 O2

16 CO2
+
18 H2O
+
Energy
A) Balance the chemical equation
B) If 2.51 moles of O2 react completely how many moles of CO2 will be produced?
𝟐. 𝟓𝟏 𝐦𝐨𝐥 𝐎𝟐 𝐱
𝟏𝟔 𝐦𝐨𝐥 𝐂𝐎𝟐
= 𝟏. 𝟔𝟏 𝐦𝐨𝐥 𝐂𝐎𝟐
𝟐𝟓 𝐦𝐨𝐥 𝐎𝟐
C) How many moles of C8H18 will be required to react completely with 2.4 moles of O2?
𝟐 𝐦𝐨𝐥 𝐂𝟐 𝐇𝟏𝟖
𝟐. 𝟒 𝐦𝐨𝐥 𝐎𝟐 𝐱
= 𝟎. 𝟏𝟗 𝐦𝐨𝐥 𝐂𝟖 𝐇𝟏𝟖
𝟐𝟓 𝐦𝐨𝐥 𝐎𝟐
D) Would this be an endothermic or exothermic reaction? How do you know?
Exothemic. Energy is being produced (given off).
34.
2 KClO3 + Energy  2 KCl
+ 3 O2
A) Balance the equation
B) Would this be an endothermic or exothermic reaction? How do you know?
Endothermic. Energy is being taken in
35.
Ca
+
2 H2O

Ca(OH)2
+
H2
A) Balance the chemical equation
B) If 2.5 moles of H2O react, how many moles of calcium hydroxide will be formed?
𝟏𝐦𝐨𝐥 𝐂𝐚(𝐎𝐇)𝟐
𝟐. 𝟓 𝐦𝐨𝐥 𝐇𝟐 𝐎 𝐱
= 𝟏. 𝟐 𝐦𝐨𝐥 𝐂𝐚(𝐎𝐇)𝟐
𝟐 𝐦𝐨𝐥 𝐇𝟐 𝐎
C) If 1.5 grams of calcium react, how many grams of hydrogen gas will be formed?
𝟏 𝐦𝐨𝐥 𝐂𝐚
𝟏 𝐦𝐨𝐥 𝐇𝟐 𝟐. 𝟎𝟏𝟓𝟖 𝐠 𝐇𝟐
𝟏. 𝟓 𝐠 𝐂𝐚 𝐱
𝐱
𝐱
= 𝟎. 𝟎𝟕𝟓 𝐠 𝐇𝟐
𝟒𝟎. 𝟎𝟕𝟖 𝐠 𝐂𝐚 𝟏 𝐦𝐨𝐥 𝐂𝐚
𝟏 𝐦𝐨𝐥 𝐇𝟐
36.
3 Mg(NO3)2(aq)
+
2 Rb3PO4(aq) 
Mg3(PO4)2(s) +
6 RbNO3(aq)
A) Balance this equation
B) Determine the states of the products
C) Write a complete ionic equation for this reaction.
3 Mg2+(ag) + 6 NO31-(aq) + 6 Rb1+(aq) + 2 PO43-(aq)  Mg3(PO4)2(s) + 6 Rb1+(aq) + 6 NO31-(aq)
D) Write the net ionic equation for this reaction.
3 Mg2+(ag) + 2 PO43-(aq)  Mg3(PO4)2(s)
E) Identify the spectator ions in the reaction.
NO31-(aq) and Rb1+(aq)
37. How many grams of HCl equal 12 moles?
𝟑𝟔. 𝟒𝟔𝟎𝟗 𝐠 𝐇𝐂𝐥
𝟏𝟐 𝐦𝐨𝐥 𝐇𝐂𝐥 𝐱
= 𝟒𝟒𝟎 𝐠 𝐇𝐂𝐥
𝟏 𝐦𝐨𝐥 𝐇𝐂𝐥
38. How many grams of Pb(NO3)2 equal 4.2 moles?
𝟑𝟑𝟏. 𝟐𝟎𝟗𝟖 𝐠 𝐏𝐛𝐍𝐎𝟑
𝟒. 𝟐 𝐦𝐨𝐥 𝐏𝐛𝐍𝐎𝟑 𝐱
= 𝟏𝟒𝟎𝟎 𝐠 𝐏𝐛𝐍𝐎𝟑
𝟏 𝐦𝐨𝐥 𝐏𝐛𝐍𝐎𝟑
39. How many moles of C8H18 equal 24.5 grams?
𝟏 𝐦𝐨𝐥 𝐂𝟖 𝐇𝟏𝟖
𝟐𝟒. 𝟓 𝐠 𝐂𝟖 𝐇𝟏𝟖 𝐱
= 𝟎. 𝟐𝟏𝟒 𝐦𝐨𝐥
𝟏𝟏𝟒. 𝟐𝟐𝟕𝟖 𝐠 𝐂𝟖 𝐇𝟏𝟖
40. How many moles of Ca(OH)2 equal 345.0 grams?
𝟏 𝐦𝐨𝐥 𝐂𝐚(𝐎𝐇)𝟐
𝟑𝟒𝟓. 𝟎 𝒈 𝐂𝐚(𝐎𝐇)𝟐 𝐱
= 𝟒. 𝟔𝟓𝟔 𝐦𝐨𝐥 𝐂𝐚(𝐎𝐇)𝟐
𝟕𝟒. 𝟎𝟗𝟐𝟔 𝐠 𝐂𝐚(𝐎𝐇)𝟐
41. Define/describe:
A) Endothermic reaction – A reaction that takes in energy from the environment
B) Exothermic reaction - A reaction that gives off energy to the environment
42. When red hot iron reacts with steam it makes solid magnetite (Fe3O4) and hydrogen gas.
A) Write a balanced equation for the reaction.
3 Fe(s) +
4 H2O(g)
 x
Fe3O4(s)
+
4 H2(g)
B) If 36 g of water and 167 g of iron are combined, based on the production of Fe3O4
1) Determine which reactant is limiting.
1 mol H2 O
1 mol Fe3 O4
231.5326 g Fe3 O4
For water: 36 gH2 O x
x
x
= 115.7 g Fe3 O4
18.0152 g H2 O
4 mol H2 O
1mol Fe3 O4
For Iron: 167 g Fe x
1 mol Fe
1 mol Fe3 O4
231.5326 g Fe3 O4
x
x
= 231 g Fe3 O4
55.845 g Fe
3 mol Fe
1mol Fe3 O4
Water is the limiting reactant.
2) Determine the amount of excess reactant consumed.
1 mol Fe3 O4
3 mol Fe
55.845 g Fe
116 g Fe3 O4 x
x
x
= 84 g Fe Consumed
231.5326 g Fe3 O4
1 mol Fe3 O4
1mol Fe
3) Determine the amount of excess reactant remaining.
167 g – 84 g = 83 g remaining
4) What is the percent yield if 96.5 g of Magnetite were actually produced by
the reaction?
𝐀𝐜𝐭𝐮𝐚𝐥 𝐲𝐢𝐞𝐥𝐝
𝟗𝟔. 𝟓 𝐠
% 𝐘𝐢𝐞𝐥𝐝 =
𝐱 𝟏𝟎𝟎 =
𝐱 𝟏𝟎𝟎 = 𝟖𝟑 %
𝐄𝐱𝐩𝐞𝐜𝐭𝐞𝐝 𝐲𝐢𝐞𝐥𝐝
𝟏𝟏𝟓. 𝟕 𝐠
43. Calculate the molar concentration of a 3.00 g sample of NaCl dissolved in 8.0 L of water.
𝟏𝐦𝐨𝐥
𝟑. 𝟎𝟎 𝐠 𝐱
𝐦𝐨𝐥 𝐬𝐨𝐥𝐮𝐭𝐞
𝟓𝟖. 𝟒𝟒𝟐𝟖 𝐠
𝐌=
=
= 𝟎. 𝟎𝟎𝟔𝟒𝟐 𝐌
𝐋 𝐬𝐨𝐥𝐮𝐭𝐢𝐨𝐧
𝟖. 𝟎 𝐋
44. How many grams of NaCl are needed to make 250 mL of a solution with the same
concentration as the one in the last problem?
𝟓𝟖. 𝟒𝟒𝟐𝟖 𝐠
𝐌 𝐱 𝐕 = (𝐨. 𝐨𝐨𝟔𝟒𝟐 𝐌)(𝟎. 𝟐𝟓𝟎 𝐋) 𝐗
= 𝟎. 𝟎𝟗𝟑𝟕 𝐠
𝟏 𝐦𝐨𝐥
45. How would an increase in the temperature affect the solubility of a:
A) solid – The solid would increase in solubility (more can be dissolved).
B) gas – The gas would be less soluble driving the gas out of the solution.
46. How would breaking a solid into smaller pieces affect:
A) Surface Area - Increases
B) Collisions - Increases
C) Reaction Rate - Increases
D) Solubility - Increases
47. Why does stirring or shaking cause a solid to dissolve faster?
These spread out the solid (if in powder form) increasing surface area. They also move the
particles increasing the frequency of collisions.
48. Water can be used to dissolve ionic substances, polar covalent substances and some
nonpolar covalent substances. What is special about water that allows this to be true?
Water has highly polar covalent bonds. The high polarity allows it to solvate both ions and
polar molecules. Some nonpolar compounds have polar parts that cancel out due to their
VSEPR shapes. Water is small enough to attach to these parts causing the molecule to
dissolve.
49. Define Colligative Properties.
Properties of a solution that depend on the number (concentration) of solute particles.
50. What affect will a large number of solute particles (high concentration) have on:
A) Melting Point Depression - The more particles the lower the melting point
B) Boiling Point Elevation - The more particles the higher the boiling point.
C) Vapor Pressure - The more particles the lower the vapor pressure.
51. Arrange the following compounds in order from smallest to largest affect on colligative
properties. Explain why you placed them in this order.
BaCl2, ZrBr4, Sc(NO3)3, Alcohol, MgSO4
Alcohol (1 particle)
Sc(NO3)3 (4 particles)
MgSO4 (2 particles)
ZrBr4 (5 particles
BaCl2 (3 particles)
The more particles the larger the effect.
52. Why doesn’t Silver Chloride have any effect on colligative properties?
Silver chloride is not soluble in water and therefore does not contribute any particles.
53. As a gas is heated, what happens to:
A)
B)
C)
D)
E)
Average Kinetic Energy - Increases
Particle Motion - Increases
Space between particles - Increases
Volume (if Pressure is constant) - Increases
Pressure (if Volume is constant) - Increases
54. Based on the combined gas law, what would you do to a gas if you needed to increase
the pressure?
Increases temperature and decreases volume
55. Why don’t real gases match the assumptions for Ideal Gases?
Real particles have volume and lose energy in collisions.
56. Which real gas comes closest in properties to an Ideal gas?
Helium
Use the ideal gas law to solve the following problems.
57. Find the pressure of a 4.00 g sample of nitrogen gas that has a volume of 6.25 L at a
temperature of 43⁰C. (R = 0.0821 L ATM/mol K)
𝒏 = 𝟒. 𝟎𝟎 𝒈 𝑵𝟐 𝒙
𝟏 𝒎𝒐𝒍 𝑵𝟐
𝟐𝟖.𝟎𝟏𝟑𝟒 𝒈 𝑵𝟐
= 𝟎. 𝟏𝟒𝟑 𝒎𝒐𝒍 𝑵𝟐
𝑽 = 𝟔. 𝟐𝟓 𝑳
𝑻 = 𝟒𝟑𝒐 C + 273 = 316 K
𝐋 𝐀𝐭𝐦
(𝟎. 𝟏𝟒𝟑 𝐦𝐨𝐥) (𝟖. 𝟎𝟖𝟐𝟏
) (𝟑𝟏𝟔 𝐊)
𝐧𝐑𝐓
𝐦𝐨𝐥 𝐊
𝐏=
=
= 𝟓𝟖. 𝟒 𝐀𝐭𝐦
𝐕
𝟔. 𝟐𝟓 𝐋
58. What is the Volume of a gas sample containing 1.25 x 1024 particles with a pressure of 921 torr
and a temperature of 8.43oC. (R = 0.0821 L ATM/mol K)
𝟏 𝐦𝐨𝐥
𝐧 = 𝟏. 𝟐𝟓 𝐱 𝟏𝟎𝟐𝟒 𝐩𝐚𝐫𝐭𝐢𝐜𝐥𝐞𝐬 𝐱
= 𝟐. 𝟎𝟖 𝐦𝐨𝐥
𝟔. 𝟐𝟓 𝐱 𝟏𝟎𝟐𝟑 𝐩𝐚𝐫𝐭𝐢𝐜𝐥𝐞𝐬
𝑷 = 𝟗𝟐𝟏 𝒕𝒐𝒓𝒓 𝒙
𝟏 𝑨𝒕𝒎
= 𝟏. 𝟐𝟏 𝑨𝒕𝒎
𝟕𝟔𝟎 𝒕𝒐𝒓𝒓
𝑻 = 𝟖. 𝟒𝟑 𝒐𝑪 + 𝟐𝟕𝟑 = 𝟐𝟖𝟏. 𝟒𝟑 𝑲
𝐋 𝐀𝐭𝐦
(𝟐. 𝟎𝟖 𝒎𝒐𝒍) (𝟖. 𝟎𝟖𝟐𝟏
) (𝟐𝟖𝟏. 𝟒𝟑 𝑲)
𝒏𝑹𝑻
𝐦𝐨𝐥 𝐊
𝑽=
=
= 𝟑𝟗. 𝟕 𝑳
𝑷
𝟏. 𝟐𝟏 𝑨𝒕𝒎