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SGSF-1K
PreAP ESPS Spring Final
Study Guide
KEY
Unit 1 – Safety & Matter
1) Why should you clean any dirty glassware before using it for an experiment?
Avoid contamination
2) Why do goggles or face shields provide more eye protection than personal eyeglasses?
Eyeglasses don’t protect from side splashes.
3) What should you do with long hair in the laboratory?
Tie it up.
4) What can happen if you wear loose clothing or jewelry in a laboratory?
Catch on things, get contaminated by chemicals
5) What is the best type of footwear to wear in a lab?
Closed toed shoes
6) What should you do if a piece of lab equipment is not working properly?
Tell the instructor
7) What should you do if acid is splashed on your skin? What if it’s splashed all over your clothes?
Skin: Flush with lots of water; Clothes: get them off and flush skin with lots of water
8) Why should you clean your hands after working with chemicals? What should you use to clean
them? Chemicals could be on your hands – contamination. Wash with soap and water
9) How does hot glass look compared to cold glass?
Same
10) What does it mean if a material is “flammable”?
It will easily ignite and burn.
11) What should you do if an uncontrolled fire occurs in the laboratory?
Immediately tell the instructor
12) When should you wear special eye protection devices (such as goggles) in the laboratory?
Whenever working with chemicals or glassware
13) In the lab, what should you do it you don’t understand a direction or part of a procedure?
Ask the instructor
14) What are you supposed to do with chemical wastes after you finish a lab?
Dispose of according to instructions from instructor
15) What could happen if you point the end of a test tube that you’re heating towards someone?
Material could shoot out the end and hit someone.
16) What’s the best way to pick up a piece of glassware you’ve been heating?
With tongs
17) If you get injured in the laboratory (cut, burn, etc.), what should you do?
Immediately tell the instructor
18) Why is it important to check glassware for chips or cracks?
Cracks and chips weaken glass so it more easily breaks
19) Is it appropriate to return all unused chemicals to their original containers?
No
20) Should you start working on a lab even if the instructor is not yet present?
No
21) When is okay to remove chemicals or other equipment from the laboratory?
Never (only under direct supervision of an instructor)
22) Why are unauthorized experiments prohibited?
Potential for unexpected reactions and/or consequences
23) When are students allowed to enter the chemical preparation/storage area?
Never
24) Why do people wear laboratory aprons?
To protect clothes and skin from chemical splashes/contamination
25) Is it okay to pick up broken glass with your bare hands as long as the glass is placed in the trash?
No, ask instructor (who will use tongs and broom with dust pan)
26) Why should you never leave a lit burner unattended?
Flame could go out and allow unburned gas to accumulate in room – potential for explosions
27) What instrument is used to weigh objects?
Balance
28) What piece of lab equipment is a small glass container used to view chemical reactions or to heat
small amounts of a substance?
Test Tube
29) What instrument is used to measure volume very precisely?
Graduated Cylinder
30) What instrument is a wide-mouthed container used to transport, heat, and store substances?
Beaker
31) Define:
 Matter – ANYTHING THAT HAS MASS AND TAKES UP SPACE

Substance – MATTER THAT ALWAYS HAS THE EXACT SAME COMPOSITION

Mixture – 2 OR MORE SUBSTANCES TOGETHER - COMPOSITION CHANGES

Element – SUBSTANCE THAT CAN’T BE BROKEN DOWN INTO SIMPLER
SUBSTANCES

Compound – COMPOSED OF 2 OR MORE ELEMENTS IN FIXED RATIO
32) Name two examples each of an element, and a compound
 Element:
OXYGEN, CARBON, GOLD, SILVER, ETC.

Compound:
WATER, SALT, SUGAR, CARBON DIOXIDE, SILICON DIOXIDE, ETC.
33) How are a compound and a mixture different?
MIXTURE RETAINS MANY PROPERTIES OF IT THE COMPONENTS; NO FIXED
RATIO
34) Describe and give an example of each:
 Heterogeneous mixture: SAND, RIVER WATER, SALSA, CHOCOLATE CHIP COOKIE

Homogeneous mixtures. MILK, KOOL-AID, SUGAR COOKIE, GASOLINE

Record which picture represents each type: a - HOMOGENEROUS; b - HETEROGENEOUS
35) Define:
 Solution – HOMOGENOUS MIXTURE OF SOLUTE DISSOLVED IN SOLVENT
o Solvent – DOES THE DISSOLVING
o Solute – IS DISSOLVED

Suspension – HETEROGENEOUS MIXTURE THAT SEPARATES INTO LAYERS OVER
TIME

Colloid – HOMOGENEOUS MIXTURE WITH PARTICLE SIZE BETWEEN
SUSPENSION AND SOLUTION

Alloy – HOMOGENEOUS MIXTURE OF 2 OR MORE METALS
36) Describe the following: (in terms of motion, volume & shape)
 Solid – Vibrate - Locked in place; FIXED VOLUME; FIXED SHAPE

Liquid – Glide – still attracted to each other; FIXED VOLUME; INDEFINITE SHAPE

Gas – Bounce; INDEFINITE VOLUME; INDEFINITE SHAPE
37) Describe these changes as physical or chemical. Substances are……
 Mixed
P
Burned
C
Melted
P

Rusted
C
Cut
P
Cooked
C

Spoiled
C
Dissolved
P
Digested
C

Bent
P
Frozen
P
Reacted
C
38) What does the Kinetic Theory state? ALL PARTICLES OF MATTER ARE IN CONSTANT
MOTION
39) The temperature of a substance does not change during a PHASE CHANGE
40) Complete this diagram with the 6 common phase changes:
A. DEPOSITION
B. SUBLIMATION
C. VAPORIZATION
D. CONDENSATION
E. MELTING
F. FREEZING
41) Define
a. Heat of Fusion – ENERGY THAT MUST BE ABSORBED TO CHANGE FROM
SOLID TO LIQUID
b. Heat of Vaporization – ENERGY THAT MUST BE ABSORBED TO CHANGE
FROM LIQUID TO GAS
42) Label the following phase change diagram with the states of matter and all 6 phase changes. Use
arrows to display the direction of each specific phase change.
43) What volume of silver metal will have a mass of exactly 2500.0 g. The density of silver is 10.5
g/cm3.
D = M/V  V = M/D  V = 2500.0g/10.5g/cm3 = 238.1cm3  238cm3
44) A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The block weighs 1587 g. From
this information, calculate the density of lead.
D = 1587g/(4.50cm x 5.2cm x 6.00cm) = 11.3g/cm3
45) What is the mass of the ethanol that exactly fills a 200.0 mL container? The density of ethanol is
0.789 g/mL.
M = D x V  M = (0.789 g/mL)(200.0mL) = 157.8g  158g
Unit 2 – Atomic Structure & Periodic Table
46) Subatomic particle descriptions:
A. Name the three subatomic particles: proton, neutron, electron
B. Give the location where each can be found: nucleus – proton & neutron; electrons in
electron cloud
C. Give their electric charges: proton is +; electron is - ; neutron is neutral
D. Give their relative masses: proton ~ neutron are about equal in mass (1); electron
~1/2000th as much
E. Describe the composition and characteristics of the nucleus: protons with (+) and neutrons
(neutral) for net + charge; almost all of atom’s mass in nucleus
47) Atomic number and mass number:
A. What determines the atomic number of an atom? - # of protons
B. What determines the atomic mass of an atom? - # of protons + # of neutrons
C. How can you determine how many neutrons will be in a given atom? Atomic mass # –
atomic number (i.e. # of protons) = # of neutrons
D. For an atom to be neutral, what subatomic particles have to have been present in the same
number? – protons (+) and electrons (-)
E. What number is unique for any given element? Protons (atomic #)
48) Isotopes & Ions
A. Define “isotope”: - atoms of an element with differing numbers of neutrons (and hence,
differing atomic mass)
B. What is same about all isotopes of a given element? # of protons (& # of electrons in
neutral atom)
C. What is different between isotopes of a given element? # of neutrons
D. What determines the listed atomic mass for an element with many isotopes? - weighted
average, isotopes have a greater effect than uncommon ones
E. Define “ion” – an atom of an element that carries a charge (not neutral)
F. How does a neutral atom become an ion? – gaining of losing electrons (gaining = negative
charge; losing = positive charge)
G. How do you determine the charge of an atom? - # of protons - #of electrons
49) Electron energy levels:
A. The number of energy levels filled in an atom is determined by what? What is the new #
sequence? – how many electrons the atom has (always fill lowest first); 2, 8, 18, 32
B. What causes an electron to jump to a different energy level? Atom gains or loses energy
C. Define ground state: - all the electrons in an atom have lowest possible energy
D. Define excited state; - at least one electron at a higher energy level than ground state
E. What causes the glow of a neon light? – electrons dropping from excited state to ground
state by emitting energy as light
F. In a flame test, which color corresponds with the highest excited energy level and which
with the lowest? Red is lowest; violet is highest
50) Complete the following table:
Element
Symbol
Se
C
P
Fe
Xe
Zr
Br
Ca
Mn
Na
Ag
Atomic
Number
34
6
15
26
54
40
35
20
25
11
47
Atomic
Mass
78.96
12.011
30.974
55.847
131.30
91.22
79.904
40.08
54.938
22.990
107.87
Mass
Number
79
12
31
56
131
91
80
40
55
23
108
Protons
Neutrons
Electrons
34
6
15
26
54
40
35
20
25
11
47
45
6
16
30
77
51
45
20
30
12
61
34
6
15
26
54
40
35
20
25
11
47
51) Draw a Bohr Diagram for Carbon, Sodium, and Sulfur.
C
52) Calculate the average atomic mass for the following:
A. Calculate the average atomic mass of sulfur if 95.00% of all sulfur atoms have a mass of
31.972 amu, 0.76% has a mass of 32.971amu and 4.22% have a mass of 33.967amu.
32.021amu
B. An atom has the following three isotopes: 24 amu (percent abundance = 78.99%), 25 amu
(percent abundance = 10.00%), and 26 amu (percent abundance = 11.01%). Calculate the
average atomic mass.
24.320amu
53) History of the Periodic Table
A. Who first arranged the periodic table by atomic mass?
Dmitri Mendeleev
B. What characteristics did Mendeleev use to place the elements in order when creating the
periodic table?
Rows - Increasing mass
Columns – Similar properties
C. What are valence electrons?
Electrons in the highest occupied energy level
D. What do elements that belong to the same group have in common?
They have the same number of valence electrons; similar chemical properties
E. What is the “octet rule”?
Atoms are most stable if they have filled or empty outer shell of electrons
Filled shell contains 8 electrons (octet)
Except for H and He (atomic #1 & #2)
Atoms gain, lose, or share electrons to make filled or empty outer shell
Atoms gain, lose, or share electrons based on what is easiest.
54) Structure of the Periodic Table
A. What is the definition of atomic mass?
Individual Atoms: Number of protons plus number of neutrons. (p+ + n)
Elements: Weighted average of the masses of the isotopes.
B. What units do you put on atomic mass?
amu (Atomic Mass Units)
1 amu = 1/12 mass of Carbon-12
C. Does the atomic number increase or decrease as you move from left to right? How much does
it change from one element to the next?
Increases by exactly 1 for each element as you move from left to right
D. The vertical columns on the Periodic Table are called what?
Groups or Families (memory aid: families stand up for each other)
E. The horizontal rows are called what?
Periods (memory aid: read the row like a sentence and end it with a “period”.)
F. What states of matter are represented by the metals at room temperature (solid, liquid, and/or
gas)?
Solid and liquid (all but mercury are solids)
G. Every element in the Carbon group (family) has how many valence electrons?
4
H. The majority of the Periodic Table is made up of what class of element?
metals
I. Which group is the least likely to react?
Noble Gases
J. The chemical and physical properties of Li are most similar to the chemical and physical
properties of what other elements?
K, Rb, Cs, Fr
55) Trends on the Periodic Table
A. What is the trend in reactivity within the alkali metals and the alkaline earth metals?
Reactivity increases as you go down the column (most reactive elements at the bottom)
B. What is the trend in reactivity within the halogens?
Reactivity decreases as you go down the column (most reactive elements at the top)
C. There are two groups (families) on the Periodic Table that are considered the MOST reactive,
what are they (give name and number)?
Group 1 (IA) – Alkali Metals
Groups 17 (VIIA) – Halogens
D. Using your answer from the previous question, how many valence electrons does each of those
groups (families) have?
Group 1 (IA) – Alkali Metals has 1 valence electron
Groups 17 (VIIA) – Halogens has 7 valence electrons
E. Na is a very reactive element in Period 3, what element in Period 2 would be similar in its
reactivity?
Lithium (both are in group 1 – IA)
F. Looking at Group 2 (IIA) which element would be the MOST reactive?
Radium
G. If a reaction is going to take place, which Halogen would you expect to react the fastest?
Fluorine
H. If K reacts very violently in water, what you expect Fr to do?
React even more violently
I. Define Ionization Energy. Which metal has the highest ionization energy? Why?
The amount of energy needed to remove an electron from an atom.
Li – because it’s one valence electron is in energy level 2 which is close to the nucleus
resulting in a much stronger magnetic pull on it than on the valence electrons of other
members of the group which as in higher energy level.
J. Define Electronegativity? Which element has the highest electronegativity? Why?
The ability of an atom to attract or gain an electron
F – because it only needs to gain 1 electron to fulfill the octet rule and its highest energy
level is 2 so the magnetic pull on that electron will be greater due to the energy level’s
proximity to the nucleus
K. List the following elements in order of increasing ionization energy: Li, Ti, Ba, K, Fr 
Fr, Ti, Ba, K, Li
L. List the following elements in order of increasing reactivity: C, S, I, Te, F  C, Te, S, I F
M. Label the numbers and letters in the above Periodic Table
1 – Alkali Metals
2 – Transition Metals
3 - Halogens
4 – Lanthanide Series
5 – Actinide Series
6 - Alkaline Earth Metals
7 – Noble gasses
A – Metals (includes H)
B – Non-metals
C – Metalloids
N. Use the periodic table to complete the following table:
Element
# of valence electrons
Magnesium
2
Oxygen
6
Neon
8
Aluminum
3
Lewis Dot Diagram
O. List and illustrate the 4 periodic trends we discussed in this section:
Metal Reactivity Increases:
Nonmetal Reactivity Increases:
56) Electron Configuration:
Unit 3 – Bonding
Review Ionic Bonds
57) What two types of elements will transfer electrons to form an ionic bond?
Metal & nonmetal
58) What is an ionic bond?
An electrostatic force that holds together a cation and anion
59) How can you describe the electrical charge of an ionic compound?
Neutral
60) Metals will lose all of their valence electrons to form a positively charged cation. The oxidation
number will be equal to the number of valence electrons.
61) Nonmetals will gain valence electrons to form a negatively charged anion. The oxidation number
will be equal to 8 – the # of valence electrons.
62) Which elements are least likely to undergo bonding? Noble Gases
63) Write the oxidation number for the following elements.
a. Cesium
___+1____
b. Potassium
___+1____
c. Aluminum
___+3____
d. Iron (III)
___+3____
e. Magnesium ___+2____
f. Barium
___+2____
g. Nitrogen
___-3____
h. Iodine
___-1____
i. Oxygen
___-2____
j. Selenium
___-2____
64) Which alkali metal has the highest ionization energy? Hydrogen
65) Which alkaline earth metal has the lowest ionization energy? Radium
66) What properties characterize ionic compounds?
 High melting point & boiling point
 Solid at room temperature
 Brittle
 Ionic solids can’t conduct electricity
 Molten ionic compounds can conduct electricity
Transferring of Electrons
67) Please draw the transfer of electrons for the following:
a. NaBr
b. MgO
c.
CaF2
d. K2S
68) How do the following atoms fulfill the octet rule when they form ionic bonds? What valence or
oxidation number does each of these have?
a. Lithium
Loses 1 e-; valence 1+
b. Beryllium
Loses 2 e-; valence 2+
c. Boron
Loses 3 e-; valence 3+
d. Carbon
Loses 4 e-; valence 4e. Nitrogen
Gains 3 e-; valence 3f. Oxygen
Gains 2 e-; valence 2g. Fluorine
Gains 1 e-; valence 1h. Neon
Full shell – no change
69) Define the indicators (what you look for) that identify each type of compound based on their names
and their formulas:
a. Simple Binary Ionic –
i. Name – Group 1A/2A metal or Aluminum & Nonmetal with “ide” ending
ii. Formula – Group 1A/2A metal or Aluminum & Nonmetal
b. Polyvalent Binary Ionic –
iii. Name – Roman Numerals!
iv. Formula – Transition Metal & Nonmetal
c. Polyatomic Ionic –
v. Name – Metals & anions with “ate” or “ite” endings
vi. Formula – Parentheses!
Writing Formulas for Ionic Compounds
Write the formula for the following ionic compounds:
70) Barium oxide
BaO
71) Lithium nitride
Li3N
72) Magnesium fluoride
MgF2
73) Lithium phosphate
Li3PO4
74) Calcium iodide
CaI2
75) Copper (II) nitrite
Cu(NO2)2
76) Beryllium chloride
BeCl2
77) Gold (III) carbonate
Au2(CO3)3
78) Potassium sulfate
K2SO4
79) Calcium hydroxide
Ca(OH)2
80) Barium chloride
BaCl2
81) Lithium oxide
Li2O
82) Magnesium oxide
MgO
83) Calcium fluoride
84) Iron (II) chloride
85) Silver bromide
CaF2
FeCl2
AgBr
Naming Ionic Compounds
Write the name for the following ionic compounds. Be sure to include parenthesis with Roman
numeral for the transition metals:
86) AlBr3
87) Fe2(SO3)3
88) MgCl2
89) K3PO4
90) TiCl2
91) Cr2O3
92) Ag2S
93) Cu(OH)2
94) Zn3N2
95) AgNO3
96) PbF2
97) Cu2S
98) CaF2
99) FeCl3
100) Li3P
101) Au(C2H3O2)3
aluminum bromide
iron (III) sulfite
magnesium chloride
potassium phosphate
titanium (II) chloride
chromium (III) oxide
silver (I) sulfide
copper (II) hydroxide
zinc (II) nitride
silver (I) nitrate
lead (II) fluoride
copper (I) sulfide
calcium fluoride
iron (III) chloride
lithium phosphide
Gold (III) acetate
Review Covalent Bonding
102) How many valence electrons are in the following:
a. Carbon
4
103)
104)
105)
106)
b. Silicon
4
c. Chlorine
7
d. Nitrogen
5
What is the name given to the pairs of valence electrons that do not participate in bonding?
Lone or unpaired
What is the name given to the pairs of valence electrons that do participate in bonding?
Shared pair
In a covalent bond, electrons are shared between atoms.
Which of the following is a covalent molecule (circle one)?
a. LiCl
b. BaO
c. NO2
d. MgBr2
Writing Formulas for Covalent Compounds
107)
Write the formula for the following covalent compounds:
a. Sulfur tribromide
SBr3
b. Dinitrogen tetrachloride
N2Cl4
c. Dihydrogen monoxide
H2 O
d. Phosphorus Trihydride
PH3
e. Tetracarbon octahydride
C4H8
Nomenclature for Covalent Compounds
108)
Name the following covalent compounds:
a. N4O6 tetranitrogen hexoxide
109)
b. SO3
sulfur trioxide
c. SeF4
selenium tetrafluoride
d. PCl5
phosphorous pentachloride
e. XeI4
xenon tetraiodide
How many electrons are shared in a single bond? Double bond? Triple bond?
a. Single bond: 2
b. Double bond: 4
c. Triple bond: 6
110)
Give 2 examples of a diatomic molecule.
Hydrogen, oxygen, fluorine, nitrogen, chlorine, iodine, bromine
111)
List 3 properties of covalent molecules.
a. Low melting point
b. 2 nonmetals
c. Isn’t soluble
112)
In a nonpolar covalent bond, electrons are shared equally.
113)
In a polar covalent bond, electrons are shared unequally.
114)
Label the partial positive and partial negative charge on the molecule below.
115)
Define electronegativity (use the textbook).
Ability of an atom to attract electrons
116)
Electronegativity determines the type of bond.
117)
Draw an example of a polar molecule.
118)
Draw an example of a nonpolar molecule.
119)
What is an intramolecular force? Provide 2 examples.
The forces within a molecule are called intramolecular forces. This includes ionic and
covalent bonds.
120)
What is an intermolecular force?
The forces between 2 molecules are called intermolecular forces. These forces are much
weaker than intramolecular forces.
121)
Label the intermolecular and intramolecular forces on the figure below:
122)
Which force is the strongest, intermolecular or intramolecular?
Intramolecular
Why are polar molecules “sticky”?
Because the partial charges are attracted to one another making them hard to separate.
123)
124) Fill in the chart below:
Molecule: F2
Name: fluorine
Polar or nonpolar bonds: nonpolar
Lewis Structure:
Molecule: HCl
Name: hydrogen monochloride
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: CCl4
Name: carbon tetrachloride
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: SiO2
Name: silicon dioxide
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: HI
Name: hydrogen monoiodide
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: N2H4
Name: dinitrogen tetrahydride
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: CO
Name: carbon monoxide
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: CCl4
Name: carbon tetrachloride
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: H2O
Name: dihydrogen monoxide
Polar or nonpolar bonds: polar
Lewis structure:
Molecule: CO2
Name: carbon dioxide
Polar or nonpolar bonds: polar
Lewis structure:
Unit 4 – Chemical Reaction
125)
Which number represents the coefficient? 5
126)
Which number represents the subscript? 2
127)
How many atoms of Hydrogen are present? 10
128)
What is a chemical reaction? A change that results in a substance with a different identity.
129)
A substance that enters into a chemical reaction is called what? Reactants
130)
A substance that is formed by a chemical reaction is called what? Products
131)
T/F
132)
What does the  stand for? Yield
133)
The law of conservation of mass states that mass can’t be created or destroyed.
134)
Write the general equations for the 5 main types of chemical reactions.
a.
5H2
In a true chemical reaction a new substance must form. True
Synthesis = A + B  AB
b. Decomposition = AB  A + B
c. Single Replacement = A + BC  AC + B
d. Double Replacement = AB + CD  AD + CB
e. Complete Combustion = CxHy + O2  CO2 + H2O
135) Balance the following chemical reactions, label the reactants and products, and identify the type
of reaction.
a.
b.
c.
d.
e.
f.
g.
EVERYTHING ON THE LEFT OF THE ARROW IS A REACTANT
EVERYTHING ON THE RIGHT OF THE ARROW IS A PRODUCT
2Au2O3 → 4Au + 3O2
Reaction type: Decomposition
1SiC + 2Cl2 → 1SiCl4 + 1C
Reaction type: Single
1S8 + 24F2 → 8SF6
Reaction type: Synthesis
3KOH + 1H3PO4 → 1K3PO4 + 3H2O
Reaction type: Double
2H2 + 1O2 → 2H2O
Reaction type: Synthesis
1S8 + 12O2 → 8SO3
Reaction type: Synthesis
1Zn + 2HCl → 1ZnCl2 + 1H2
Reaction type: Single
136)
Which metals would be replaced in a solution, if a piece of lead metal were added?
Cu, Hg, Ag, Au
137) You have a mixture that contains solutions of both Sodium and Aluminum compounds. You
wish to ‘recover’ the aluminum but not the sodium. What other metal should you add to this
mixture?
Li, Ca, or Mg
138)
Which of the metals could NOT be used to remove mercury from a nitrate solution?
Ag or Au
139)
Write the equation for “magnesium reacts with chlorine to produce magnesium chloride”
Mg + Cl2  MgCl2
140)
CuCl2 + Mg  Cu + MgCl2 Which is the more active metal?
Magnesium
Unit 5 - Nuclear
141)
On December 7, 1941, the Japanese attacked Pearl Harbor and the US declared war on Japan
and then joined World War II.
142)
What rare uranium isotope is useful in the fission reaction?
U-235
143)
Where did they locate the facility test site?
144)
Who was selected to run the project from conception to completion?
145)
The facility at Oak Ridge, TN was top secret. Why?
Los Alamos
Robert Oppenheimer
Creation of Y-12 (separate U isotopes) – So Germans would not know about it
146)
In 1941, what substance was created to be used to fuel the A-bomb? Plutonium
147)
What president took over for Roosevelt and had to be updated on the secret?
148)
In February 1945, the first bomb was completed which used U-235. What was it called? Little
Truman
Boy
149)
What was the name given to the bomb that used plutonium? Fat Man
150)
What code name was given to the first test of an atomic weapon? Trinity
151)
At 5:30 am, on July 16, 1945, the world entered the “Atomic Age.” How powerful was the
bomb?
~20,000 tons of TNT
152)
On August 6, Enola Gay dropped Little Boy over Hiroshima. What percentage of the city was
destroyed? 90%