Name:____________________________
Date:_____________
Period:____
Review: Chemistry I End of Course Exam
Unit 1: Measurement and Lab Skills
Unit 2: Properties of Matter
Unit 3: Atomic Theory and Structure
Unit 4: Electron Configuration and Modern Atomic Theory
Unit 5: The Periodic Table
Unit 6: Ionic Compounds – Bonding and Nomenclature
Unit 7: Covalent Compounds – Bonding and Nomenclature
Unit 8: Percent Composition and Empirical Formulas
Unit 9: Chemical Reactions
Unit 10: Stoichiometry
Unit 11: Energy Changes and Reaction Rates
Unit 12: States of Matter
Unit 13: Kinetic-Molecular Theory and Gases
Unit 14: Solutions and Equilibrium
Unit 15: Acids and Bases
Matter and its Properties
1. Place a check mark next to each item that is matter which is made of atoms
(Remember that matter has mass and takes up space).
____a. apple ______b. air
_____c. light ____d. energy ____f. heat ____g. water
2. Which type of substance (element, compound, or mixture) is represented in each diagram below.
3. T or F
T or F
T or F
T or F
________
________
________
________
1.
_______________________
2.
_______________________
3.
_______________________
4.
_______________________
a. Elements can be broken down by both physical and chemical means.
b. Compounds can only be broken down by chemical means.
c. Mixtures can be separated into different components using physical means.
d. Air is a homogeneous mixture of O2, CO2, N2 and water vapor.
4. Classify the following materials as:
element
compound
homogeneous mixture
Can be separated by _____ means.
heterogeneous mixture
(circle one)
a. cheeseburger
___________________
physical
chemical
NO
b. salt water
___________________
physical
chemical
NO
c. air
___________________
physical
chemical
NO
d. hydrogen gas (H2)
___________________
physical
chemical
NO
e. silver (Ag)
___________________
physical
chemical
NO
f. steam (H2O)
___________________
physical
chemical
NO
g. table salt (NaCl)
___________________
physical
chemical
NO
h. carbon dioxide (CO2)
___________________
physical
chemical
NO
i. ice (H2O)
___________________
physical
chemical
NO
5. a. Which type of matter contains only ONE TYPE of atom? (element, compound, or mixture)
b. Which type of matter contains two or more different atoms, chemically bonded together?
(element, compound, or mixture)
Match each item with the correct statement below.
A. element
B. compound C. heterogeneous mixture D. homogeneous mixture
___ 6. describes mixture with a non-uniform composition (has different parts)
___ 7. pure substance that cannot be changed into simpler substances by physical means
___ 8. pure substance that cannot be broken down by chemical means
___9. A uniform mixture (such as air or sugar water) in which the parts are not easy to distinguish
10. What is one difference between a mixture and a compound?
A.
B.
C.
D.
A compound has the same properties as the elements that form it.
A compound can only be separated physically.
A mixture can be separated physically.
A mixture must be uniform in composition.
Energy and its Interactions
11. Label each of the following processes as endothermic or exothermic:
______________ solid to liquid to gas
______________ gas to liquid to solid
______________ solid ice melting into liquid
______________ liquid water evaporating into gas
______________ water vapor condensing into liquid
______________ liquid water freezing into solid ice
12. Sketch a potential energy diagram for an exothermic reaction.
Label the reactants and products.
potential
energy
reaction progress
Circle the correct statement.
13.
In an exothermic reaction, heat is ( absorbed from / released to ) the surroundings,
and the surroundings ( cool down / warm up).
Touching a beaker containing this reaction would feel
(cool / warm )
In an exothermic reaction, the energy of the reactants is (higher/ lower) than the energy of the
products.
14. Sketch a potential energy diagram for an endothermic reaction.
Label the reactants and products.
potential
energy
reaction progress
Circle the correct statement.
In an endothermic reaction, heat is
( absorbed from / released to)
the surroundings,
and the surroundings ( cool down / warm up)
Touching a beaker containing this reaction would feel
(cool / warm ).
In an endothermic reaction, the PE of the reactants is (higher/ lower) than the PE of the products.
Measurement and Density
15. These values were recorded as the mass of products when the same chemical reaction was
carried out three separate times: 3.20 g; 2.87 g; 3.89 g.
a. Find the average mass. Average mass = _______________________________
b. The theoretical mass of products from that reaction was 3.45 g.
Calculate the % error.
% error = _________________________
c. Was the experiment accurate (i.e was the % error < 10%)?
d. Is the set of data precise?
e. What is the difference between accuracy and precision?:
16. Convert: 65.5 m = ?? cm __________________
17. Convert 2.3 mL into L. _____________________
18. The mass of a soft lump of metal is 214 g, and the volume is 11.3 cm3.
a. What is the formula for density?
3
b. What is the density of the metal in g/cm ?
c. Using the table of metal densities,
Density Table
Metal
Density
(in g/cm3)
Aluminum
Iron
Gold
Platinum
2.7
7.9
19.3
21.4
identify the metal: _______________
The soft metal lump is then smushed together with another lump of the exact same soft metal.
The density of the new larger lump is ( greater than / equal to / less than) the original lump.
Atomic Theory and Structure
Recall that there are 3 types of subatomic particles- protons are (+), are found in the nucleus, and
have nearly the same mass as a neutron (1 a.m.u). Electrons are (-), are found in the electron cloud,
and have nearly no mass when compared to protons or neutrons. Neutrons are like protons, only
neutral.
Atomic number = # of protons
Mass number = # proton + # neutrons
To find neutrons, subtract!
Mass number – atomic number = # of neutons
Since atoms are neutral, the # of protons = # of electrons
19. Which of the following is FALSE about subatomic particles?
A.
B.
C.
D.
Electrons are negatively charged and are the lightest subatomic particle.
Protons are positively charged and have nearly the same mass as neutrons.
Neutrons have no charge and have no mass.
The mass of a neutron nearly equals the mass of a proton.
20. All atoms of the same element have the same ____.
A.
B.
C.
D.
number of neutrons
number of protons
mass numbers
mass
21. What are atoms of the same element with different numbers of neutrons?
A.
B.
C.
D.
ions
atoms
numbers of electrons
isotopes
22. Define isotope. Give an example of a pair of isotopes.
___________________________________________________________________
___________________________________________________________________
75
35X
23. Element X has an atomic number of 35 and a mass number of 75.
How many of each subatomic particle are in a neutral atom of the element?
A.
B.
C.
D.
35 protons, 35 neutrons, and 70 electrons
35 protons, 75 neutrons, and 35 electrons
75 protons, 35 neutrons, and 40 electrons
35 protons, 40 neutrons, and 35 electrons
24. 15 moles of sodium and 15 moles of carbon have the same number of _________.
A.
B.
C.
D.
atoms
grams
A and B
NONE of the above
25. What is the molar mass of Fe3(PO4)2 ?
of CaCl2?
Define mole:___________________________________________________________________
What is Avogadro’s number? ___________________________________________________
26. Convert 4.50 moles of Fe to atoms of Fe.
27. Convert 4.03 x 1022 molecules of H2O to moles of H2O.
Electron Configuration and Modern Atomic Theory
Figure 1:
Bohr Model
28. Examine the Bohr Model in Figure 1.
Place a “G” in the box that points to the
“Ground State” electron orbit.
Place an “E” in the box that points to the
“Excited State” electron orbit.
Nucleus
How is the “excited state” different from the
“ground state?” ____________________________________________________________
eelevel 3
Figure A:
electron move from
high to low orbit
Einitial
photon
Figure B:
level 5
electron move from
low to high orbit
E)
photon
E)
e-
level 2
level 1
e-
Efinal
Efinal
Einitial
E
initial
29. Study Figures A and B above.
In which figure is energy absorbed? A or B
How do you know? _________________________________________________________
In which figure is energy emitted? A or B
How do you know? _________________________________________________________
Which figure shows a photon of the greatest energy? A or B
How do you know? _________________________________________________________
30. Write the number of valence electrons for the elements in each group in the boxes.
31. Complete the table.
Element
Symbol
# of
VALENCE
Electrons
Electron Configuration
Bohr Model
1s2 2s2 2p6 3s2 3p6
Carbon
C
[Ar] 4s2
The Periodic Table
32. A chemist needs calcium to perform an experiment in the lab and discovers that she does not have
any calcium. List the two elements that she could use for this experiment to best replace calcium.
_______________________
_____________________
Explain your choices:
______________________________________________________________
Why do elements that are in the same group behave in a similar way?
__________________________________________________________________
What is the radius trend moving across a period?__________________
What is the radius trend moving down a group?______________________
What is the largest atom? __________________ The smallest atom? _______________
33. For the elements listed above, which element has the largest atomic radius? ____
Which element has the smallest atomic radius? ____
These atoms change size from left to right. Why do their sizes follow this trend from left to right?
______________________________________________________________________________
Ionization energy is the energy needed to remove the outermost electron from an atom.
It is difficult to remove electrons from small atoms, and easier to remove them from large atoms, so…
I.E is inversely proportional to radius.
34. Arrange the elements above from lowest to highest ionization energy.
_____ < _____ < _____ < _____
35. Fill in the missing information for the element fluorine.
a. F has the ___________ atomic radius in its group.
b. F has the ___________ ionization energy in its group.
c. F has the ___________ electronegativity in its group.
d. F is the __________ reactive element in its group.
Ionic Compounds – Bonding and Nomenclature
36. In an ionic bond, electrons are
( shared / transferred / connected )
between atoms. (circle one)
37. An ionic bond is a chemical bond caused by electrostatic _________________________ between
______________ and _____________ that is formed by _________________________ electrons
between atoms. (Word bank: cation, transfer, anion, attraction)
Ionic bonds always from between ___________ and ___________. (metals, nonmetals, metalloids)
38. For the pairs of elements listed below, circle pairs that would likely form ionic bonds. Remember, an
ionic bond forms between a METAL and a NONMETAL.
C and H
Na and F
Hg and Ag
Mg and S
N and C
K and O
39. How did you know which elements in Question 38 above would form ionic bonds?
40. Name the following compounds.
(Some compounds may need a Roman numeral**.)
Fe2O3
_______________________**
NaBr
_______________________
CuCl2
_______________________ **
AlCl3
_______________________
41. Write a formula for the following chemical compounds.
copper(II) oxide
______________
calcium sulfide
______________
magnesium iodide
______________
nickel(II) bromide
______________
Covalent Compounds – Bonding and Nomenclature
42. Why do atoms share electrons in covalent bonds?
________________________________________________________________
43. Which of the following bonds is the most polar (or has the greatest difference in electronegativity)?
A.
B.
C.
D.
C—C
H—N
O—H
H—Cl
44. Which is TRUE of a nonpolar covalent bond?
A.
B.
C.
D.
electrons are shared unequally between atoms .
a cation is bonded to an anion
electrons are transferred between atoms
electrons are shared equally in atoms
water
molecule
water
molecule
45. There are strong attractions between polar water molecules which cause water to have all of the
following properties EXCEPT ____.
A.
B.
C.
D.
surface tension
liquid of greater density than solid (ice)
attraction to nonpolar molecules
higher boiling point
46. Hydrogen sulfide (H2S) boils at –60oC. Even though water is a smaller molecule that should
become a gas easier than H2S, water doesn’t boil until it reaches 100oC.
Why do water molecules require a much higher temperature to become a gas?
___________________________________________________________________
For #47-49, write the NAME or FORMULA for the following molecular compounds:
Covalent compounds are named using prefixes- mono-, di-, tri-, tetra-, penta-, etc.
47. CCI4
______________________________
S2O3
______________________________
CS2
_______________________________
48. tetraphosphorus pentoxide _______
carbon dioxide
_______
carbon monoxide
_______
49. What is the ending (suffix) used when naming binary compounds?
For #50-52, DRAW the Lewis dot structures for the following compounds.
Step 1: Determine the total # of valence electrons.
Step 2: Arrange the electrons so that each atom is surrounded by 8 electrons.
Remember that elements in Group 17 can make only one bond, while Group 16 can make two
bonds, Group 15 can make 3 bonds, and Group 14 (carbon) can make up to 4 bonds.
50. HBr
51. N2
52. a. CO2
b. O2
c. NH3
d. H2O
e. HCN
Percent Composition and Empirical Formulas
𝒑𝒂𝒓𝒕
% composition = 𝒘𝒉𝒐𝒍𝒆 x 100
53. What is the percent composition of carbon in Na2CO3 ?
The empirical formula is the simplest, or reduced version of a chemical formula.
Ex. Sugar has a true (molecular) formula of C6H12O6. The empirical (reduced) formula is CH2O.
Empirical Formula- Problem Solving Strategy
Step 1 : Start with grams of each element
Step 2: Convert to moles
Step 3: Divided by the smallest # of moles
Step 4: If needed, use a multiplier
54. What is the empirical formula of a compound that is 74.8% C and 25.2% H ?
55. What is the empirical formula of a compound that is 36.84% N and 63.16% O by mass?
The actual or true formula for a compound is called the molecular formula.
Molecular Formula- Problem Solving Strategy
Step 1: Find the molar mass of the empirical formula
Step 2: Divide the molecular weight of the true formula by the molecular weight of the empirical
formula.
Step 3: Multiply the empirical formula up by the answer from step 2.
56. A compound has an empirical formula of CH2 and a molecular weight of about 56 g/mol.
What is the molecular formula of the compound?
57. A certain compound has an empirical formula of C2H4O and a molecular weight of about 44 g/mol.
What is the molecular formula of this molecule?
Chemical Reactions
58. Label each of the following: product, reactants, subscript, coefficient, yields
H2 + Cl2
2 HCl
59. After a chemical reaction, the mass of products is _____ equal to the original mass of reactants.
A. never
B. sometimes
C. always
60. Circle ALL of the following that are TRUE about what happens in ALL chemical reactions.
A.
B.
C.
D.
Atoms are rearranged.
More energy is released than absorbed
Energy is absorbed to break the bonds of the reactants.
Energy is released when the bonds of the products are formed.
For questions #61-63,
BALANCE the reaction using coefficients when necessary.
CLASSIFY the reaction as one of the five types.
61.
__ Zn + __ HNO3  __ Zn(NO3)2 + __ H2
A. decomposition B. Synthesis
62.
D. single replacement E. double replacement
__ CaCO3 + __ HCl  __ CaCl2 + __ H2CO3
A. decomposition B. Synthesis
63.
C. combustion
C. combustion
D. single replacement E. double replacement
__ K2SO3  __ K2S + __ O2
A. decomposition B. Synthesis
C. combustion
D. single replacement E. double replacement
Reaction Rates
64. Give one reason why increasing the concentration of reactants increases the reaction rate:
more ____________________________ between reactant particles
65. List two reasons why a reaction rate increases with an increase in temperature.
more ____________________________ between _______________________
more ____________________________ of greater _____________________________.
66. The _______________(smaller, larger) the particle size of a reactant the greater surface area
available to react.
This results in ________(more, fewer) collisions between reactant particles and a
______________ (faster, slower) reaction rate.
reaction progress
67. In the diagram above of reactions B and C. Label each of the parts A, B, C, and D using the
following terms…… product, reactant, catalyzed, uncatalyzed
A: _____________________
B: _____________________________ reaction
C: _____________________________ reaction
D: _____________________
Which pathway has the lower activation energy, B or C? _____________
Which pathway corresponds to a faster reaction, B or C? ____________
Stoichiometry
Ca(OH)2 + FeCl2  CaCl2 + Fe(OH)2
74 g
127 g
111 g
?g
68. According to the reaction above, how many grams of Fe(OH)2 should be formed? ______ g
T or F ______ Because of the law of conservation of matter, the total mass of the reactants must
equal the total mass of the products.
Stoichiometry:
Stoichiometry involves changing from one chemical to another and will always involve a
balanced chemical equation (for the mole-mole step).
Problem types:
Mole- mole (1 step)
Mole-mass or mass-mole (2 steps)
Mass- mass (3 steps)
Remember to always put the given unit in the bottom of step 1 and the desired unit in the top of the
last step! Start this way and you really cannot go wrong.
69.
4 Na + O2  2 Na2O
a) How many moles of sodium will react completely with 3.82 moles of oxygen (O2)?
b) How many moles of Na2O can be produced from 13.5 mol Na?
70.
C2H4 + 3 O2  2 CO2 + 2 H2O
a) How many grams of C2H4 (28.06 g/mol) are needed to produce 66.7 grams of water?
b) How many grams of O2 are needed to react with 2.56 g C2H4 (28.06 g/mol)?
Kinetic-Molecular Theory and Gases
71. What happens to gas pressure if its volume is decreased?
What happens to the volume of a gas if the pressure is increased?
increase or decrease
increase or decrease
72. What happens to the volume of a gas if the temperature is increased? increase or decrease
What happens to the temperature of a gas if the volume is increased? increase or decrease
Gas variables include P, V, T, and the amount of gas.
73. List 3 variables and how you would change them to increase the pressure of a gas.
a)
_________________________________________
b)
_________________________________________
c)
_________________________________________
74. What does STP stand for? __________________________________________________
What are the values of STP? ___________ and ___________
For #75-78, you may refer to the following formulas and information:
P1 V1 = P2 V2
T1
T2
K = oC + 273
1 atm = 760 mmHg = 101.3 kPa
REMEMBER TO CONVERT TEMPERATURES TO KELVIN.
75. 2.00 L of a gas at 2.58 atm is compressed to a volume of 1.20 L.
What is the pressure if the temperature is constant?
76. A gas at 25.0oC and 760 mmHg is cooled to 0.00oC. What is the pressure if the volume is held
constant?
77. 5.00 L of a gas at –15.8oC and 745 mmHg is stored in a flexible container.
What is the volume at STP?
Solutions
78. Label each solution as saturated, unsaturated, or supersaturated based on the addition of solute:
_______________
_______________
_______________
79. Complete the phrase that describes the types of substances that will dissolve in each other:
_______ dissolves _______.
Therefore, polar solvents (like water) can dissolve ___________ solutes (like alcohols, sugars,
ionic compounds, etc.).
But nonpolar solutes (like fats, oils, hydrocarbons, etc.) will dissolve in _______________ solvents.
For #80-81, you may use the following formulas:
mol of solute
M = Liter of solution
.
M 1V 1 = M 2V 2
80. What is the molarity of a solution containing 1.98 moles NaCl of solute in 775 mL of solution?
81. How many mL of a 0.150 M NaBr solution are needed to make 0.100 L of 0.0500 M NaBr?
Acids and Bases
82.
List 3 properties of acids:
taste: ____________
litmus color: _______
______________________________
List 3 properties of bases:
taste: ____________
litmus color: _______
______________________________
What property to acids and bases both have?
_____________________________________________
For #83-85, you may use the following formulas:
pH = –log[H+]
pH + pOH = 14
MaVa = MbVb
83. Label each of the following as acidic (A), basic (B), or neutral (N) by writing the letter A, B, or N.
a)
hydrogen ion concentration of 1 x10-3 M _____
pH = ________
b)
[H+] = 1 x10-9 M _____
pH = __________
c)
[OH–] = 1 x10-8 M _____
pH = ___________
84. Determine the pH of the following:
85.
a)
hydrogen-ion concentration of 1 x10–4 M
______
(no calculator needed)
b)
[H3O+] = 1.0 x10–7 M
______
(no calculator needed)
c)
0.150 M hydronium ion
______
(use calculator)
10.0 mL of NaOH of unknown concentration is titrated by adding exactly 15.8 mL of 0.150 M
HCl to completely neutralize the base. What was the concentration of NaOH?
Before setting up the titration problem, write a balanced chemical equation for the neutralization
reaction.
KEY
1.
a, b, g
2.
1. compound
3.
a. F
4.
a. heterogeneous mixture
b. homogeneous mixture
c. homogeneous mixture
d. element (H2)
e. element (Ag)
f. compound (H2O)
2. mixture
b. T
c. T
Mixtures- Physical
3. element
4. element
d. T
g. compound (NaCl)
h. compound (CO2)
i. compound (H2O)
Compounds/Chemical
Elements/No. Elements cannot be broken down
5.
a. element
b. compound
6.
C
7.
10.
C
11.
endo-
12.
An exothermic graph should start with high energy reactants and end with low energy products
13.
released to
exo-
B
8.
endo-
A
endo-
warm up
9.
exo-
warm
D
exo-
higher
14.
An endothermic graph should start with low energy reactants and end with high energy
products.
absorbed from
15.
cool down
cool
lower
average= 3.32 g % error = (T-E)/T x 100 = (3.45 – 3.32)/ 3.45 x 100 = 3.7% error
The result is accurate because the average (3.32 g) is close to the theoretical (3.45 g) with less
than 5% error.
The results are not precise because the measurements are not consistent and close to each
other.
Accuracy refers to the closeness of a result to the “true” value. Precision has to do with being able to
repeat a measurement or result with consistency.
16.
6,550 cm
17.
0.0023 L
18.
Gold (closest to measured density of 18.9 g/cm3)
“equal to” b/c density does not change with the amount of a substance
19.
C
20.
B
21.
D
22.
Isotopes are different forms of the same element. The atomic number for isotopes is the
same, but the isotopes will have different mass numbers (and thus each isotope will have a
different number of neutrons).
23.
D
24.
A
14
C and 6C are examples of isotopes.
6
12
25.
357.49 g/mol
110.98 g/mole
A mole represents the amount of matter in 6.02 x 1023 atoms, molecules or formula units. The mass of
a mole , in grams, is called its molar mass.
Avogadro’s number….. It is of course 6.02 x 1023 (every good chemistry student knows this!)
26.
2.71 x 1024 atoms
27.
0.0669 mol
28.
An electron in an excited state is located at a higher energy level and must emit energy to fall
back down to the ground state at a lower energy level.
29.
Figure B b/c the electron went to a higher energy level (further from the nucleus) which requires
more energy.
Figure A b/c the electron went to a lower energy level (closer to the nucleus) which requires the
release of energy.
Figure B b/c the electron moves between the most energy levels.
30.
1
2
3
4
31.
Argon (Ar)
32.
Mg and Sr b/c they have the same number of valence electrons (2) and will react similarly
33.
Atoms get SMALLER moving across a period, but the atomic radius INCREASES going down a
group.
Nitrogen (N)
5
6
7
Carbon (C)
8
Calcium (Ca)
From this trend, we predict that atoms in the upper right hand corner of the periodic table (He
and F) will be small. Fr, found in the lower left corner of the table, is the giant of the element
world.
largest (Li)
smallest (Ne)
Atoms get smaller left to right across a period as more protons are added pulling electrons
closer without adding more electron energy levels.
34.
K < Na < Ca < Mg (Remember, the big atoms have the low I. E values. Small atoms have
high I.E.)
35.
a. smallest
b. highest
c. highest
d. most
36.
transferred
37.
An ionic bond is a chemical bond caused by electrostatic _attraction_ between
__cations___ and __anions__ that is formed by __transferring__ electrons
between atoms.
Ionic bonds always from between __metals__ and __nonmetals_.
38.
Na and F
Mg and S
K and O
39.
each has a metal and a nonmetal element
40.
Fe2O3
NaBr
CuCl2
AlCl3
iron(III) oxide
sodium bromide
copper(II) chloride
aluminum chloride
41.
copper(II) oxide
calcium sulfide
magnesium iodide
nickel(II) bromide
42.
Atoms share electrons in covalent bonds to obtain a complete valence electron level (octet).
43.
C
44.
D
45.
46.
C
It takes more energy to overcome water’s especially strong intermolecular attractive forces
CuO
CaS
MgI2
NiBr2
(hydrogen bonding) to separate the liquid molecules to the vapor phase.
47.
carbon tetrachloride, disulfur trioxide, carbon disulfide
48.
P4O5, CO2, CO
49.
-ide
50.
51.
52.
O2 has a double bond, NH3 will have 3 single bonds and one unshared e- pair, H2O has two single
bonds and 2 unshared electron pairs, HCN contains a single bond and a triple bond.
53.
11.3 % C
54.
CH4
55.
N2O3
56.
C4H8
57.
C2H4O
58.
reactants
yields
subscript
products
coefficient
59.
C
60.
A
61.
1, 2, 1, 1
D. single replacement
62.
1, 2, 1, 1
E. double replacement
63.
2, 2, 3
A. decomposition
64.
collisions
65.
more collisions between reactant particles
C
D
more collisions of greater energy
66.
smaller
67.
A:
more
higher (faster)
reactants
B: uncatalyzed reaction
C: catalyzed reaction
D: products
Reaction C has a lower activation energy caused by catalyst which allows the reaction to occur
at a faster rate.
68.
90 g
True
69.
a) How many moles of sodium will react completely with 3.82 moles of oxygen (O2)?
3.82 mol O2 x (4 mol Na) = 15.3 mol Na
(1 mol O2)
b) How many moles of Na2O can be produced from 13.5 mol Na?
13.5 mol Na x (2 mol Na2O) = 6.75 mol Na2O
(4 mol Na)
70.
a) How many grams of C2H4 (28.06 g/mol) are needed to produce 66.7 grams of water?
66.7 g H2O x (1 mol H2O) x (1 mol C2H4) x (28.06 g C2H4) = 51.9 g C2H4
(18.02 g H2O) (2 mol H2O)
(1 mol C2H4)
b) How many grams of O2 are needed to react with 2.56 g C2H4 (28.06 g/mol)?
2.56 g C2H4 x (1 mol g C2H4) x (3 mol O2) x (32.00 g O2) = 8.76 g O2
(28.06 g C2H4) (1 mol C2H4)
(1 mol O2)
71.
increase
decrease
72.
increase
increase
73.
List 3 variables you could change to cause the pressure of a gas to increase.
74.
a)
increase the number of particles (add more gas)
b)
decrease the volume (compress it)
c)
increase the temperature
Standard Temperature and Pressure
273 K (0oC)
and
75.
4.30 atm
76.
696 mmHg (or 0.916 atm)
77.
5.20 L
78.
unsaturated
79.
like dissolves like.
1.00 atm (760 mmHg)
supersaturated
polar solvents can dissolve polar solutes
saturated
nonpolar solvents can dissolve nonpolar solutes
80.
2.55 M
81.
33.3 mL
82.
taste: __sour___
litmus color: _red_
_single replacement reaction with metals to produce H2 gas__
taste: __bitter__
litmus color: _blue_
_feel slippery________________________________________
both are electrolytes (conduct electricity in solution)
83.
a) A (pH = 3)
b) B (pH = 9)
c) A (pH = 6)
84.
a) pH = 4
b) pH = 7
c) pH = 0.824
85.
0.237 M
NaOH + HCl → NaCl + H2O ( NaOH and HCl have a 1:1 mole ratio.)