ESCS Fall Final - Study Guide
Name: ______KEY______________________________________________ Period: _____ Date: _________
Unit 1 – Safety & Measurement
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 Independent and Dependent Variable (include how they are used when graphing):
Independent – X-axis – What cause change
Dependent – Y-axis – What changes – what you are measuring
32. What should you do when the data in an investigation do not support the original hypothesis?
Modify the hypothesis
33. Provide the SI base units for the following:
a. Length – Meter (m)
b. Mass – Kilogram (kg)
c. Temperature – Kelvin (K)
d. Time – Seconds (s)
e. Amount of Substance – Mole (mol)
f. Electric Current – Ampere (A)
g. Luminous Intensity – Candela (cd)
34. Convert the following numbers between standard and scientific notation:
a. 2.0094 x 103 =
2,009.4
b. 0.0000945
=
9.45 x 10-5
8
c. 4.603 x 10
=
460,300,000
d. 8.0965 x 10-4 =
0.00080965
e. 90,294
=
9.0294 x 104
f. 0.000000389 =
3.89 x 10-7
g. 390
=
3.9 x 102
35. Determine the number of significant digits in the following numbers:
a. 0.00450
=
3
b. 5,795
=
4
c. 100
=
1
d. 50,603,000
=
5
e. 90.0
=
3
f. 1,000,000,500 =
8
g. 0.78210
=
5
h. 0.000800400 =
6
Unit 2 – Matter
36. Define:
 Matter – ANYTHING THAT HAS MASS AND TAKES UP SPACE
 Substance – MATTER THAT ALWAYS HAS THE EXTACT SAME COMPOSITION
 Mixture – 2 OR MORE SUBSTANCE TOGETHER - COMPOSITION CHANGES
 Element – SUBSTANCE THAT CAN’T BE BROKEN DOWN INTO SIMPLER SUBSTANCES
 Compound – COMPOSED OF 2 OR MORE ELEMENTS IN FIXED RATIO
37. 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.
38. 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
39. Define:
 Solution – HOMOGENOUS MIXTURE OF SOLUTE DISSOLVED IN SOLVENT
o Solvent – DOES THE DISSOLVING
o Solute – IS DISSOLVED
40. List the 5 basic techniques for separation and describe methods you could use to separate the following
mixtures:
1. FILTRATION
- Sand and Water: FILTER
2. MAGNET
3. EVAPORTAION
- Salt and Water: EVAPORATION
4. DISSOLVING
5. PHYSICAL
- Iron and Sulfur: MAGNET
41. Describe the following: (in terms of volume & shape)
 Solid – FIXED VOLUME; FIXED SHAPE
 Liquid – FIXED VOLUME; INDEFINITE SHAPE
 Gas – INDEFINITE VOLUME; INDEFINITE SHAPE
42. Define each of the following
(Ex: melting point: substance changes from a _______ to a _________)
 Freezing point – LIQUID  SOLID
 Boiling point – LIQUID  GAS
 Melting point - SOLID  LIQUID
43. List the most common physical and chemical properties:
 Physical – VISCOSITY, CONDUCTIVITY, MALLEABILITY, HARDNESS, MELTING POINT,
BOILING POINT, DENSITY
 Chemical – FLAMMABILITY, REACTIVITY
44. What are the 3 types of evidence that a chemical change may have occurred?
1. FORMATION OF A GAS
2. FORMATION OF A PERCIPITATE
3. CHANGE IN COLOR
45. 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
46. What happens to the pressure of an enclosed gas if:
a. The volume of the container is increased? PRESSURE DECREASES
b. The number of particles is increased? PRESSURE INCREASES
c. The temperature if lowered? PRESSURE DECREASES
DECREASE = LESS COLLISION
INCREASE = MORE COLLISION
47. Define
 Endothermic – ABSORBS ENERGY FROM SURROUNDINGS

Exothermic – RELEASES ENERGY INTO SURROUNDINGS
Unit 3 – Atomic Structure
48. 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
49. 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 #)
50. 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. Define “ion” – an atom of an element that carries a charge (not neutral)
e. How does a neutral atom become an ion? – gaining of losing electrons (gaining = negative charge;
losing = positive charge)
f. How do you determine the charge of an atom? - # of protons - #of electrons
51. 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
Unit 4 – Periodic Table
52. 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.
53. What units do you put on atomic mass?
amu (Atomic Mass Units)
1 amu = 1/12 mass of Carbon-12
54. 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
55. The vertical columns on the Periodic Table are called what?
Groups or Families (memory aid: families stand up for each other)
56. The horizontal rows are called what?
Periods (memory aid: read the row like a sentence and end it with a “period”.)
57. 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)
58. Every element in the Carbon group (family) has how many valence electrons?
4
59. List 4 properties of metals
a. good conductors (heat and electricity)
b. solids at room temperature (except mercury)
c. malleable and ductile (not brittle)
d. luster
e. high density (mass per volume)
60. The majority of the Periodic Table is made up of what class of element?
metals
61. Which group contains the maximum number of valence electrons, making it the most chemically stable?
Noble Gases
62. The chemical and physical properties of Li are most similar to the chemical and physical properties of
what other elements?
K, Rb, Cs, Fr
63. 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)
64. What is the trend in reactivity within the halogens?
Reactivity decreases as you go down the column (most reactive elements at the top)
65. 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
66. 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
67. 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)
68. Looking at Group 2 (IIA) which element would be the MOST reactive?
Radium
69. If a reaction is going to take place, which Halogen would you expect to react the fastest?
Fluorine
70. If K reacts very violently in water, what you expect Fr to do?
React even more violently
71. What do each of the four terms in the element box represent?
55
Cs
Cesium
132.91
55 – atomic #
Cs – element symbol
Cesium – element name
132.91 – atomic mass
72. List the group numbers and names of the families for the Representative elements on the periodic table:
1- Alkali Metals
2- Alkaline Earth Metals
3- Boron Family
4- Carbon Family
5- Nitrogen Family
6- Oxygen Family
7- Halogens or Halides
8- Noble Gasses
73. Use the periodic table to complete the following table:
Element
# of valence electrons
Lewis Dot Diagram
Magnesium
2
Oxygen
6
Neon
8
Aluminum
3
74. Draw a Bohr model of Sulfur and Neon.
Sulfur
Sulfur
Neon
Neon
16
Protons
10
16
Electrons
10
16
Neutrons
10
Unit 5 – Ionic Bonding
75. What two types of elements will transfer electrons to form an ionic bond?
Metal & nonmetal
76. What is an ionic bond?
An electrostatic force that holds together a cation and anion
77. How can you describe the electrical charge of an ionic compound?
Neutral
78. 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.
79. Nonmetals will gain valence electrons to form a negatively charged anion. The oxidation number will
be equal to 8 – the # of valence electrons.
80. Which elements are least likely to undergo bonding? Noble Gases
81. Write the oxidation number for the
e. Magnesium ___+2____
following elements.
f. Barium
___+2____
a. Cesium
___+1____
g. Nitrogen
___-3____
b. Potassium
___+1____
h. Iodine
___-1____
c. Aluminum
___+3____
i. Oxygen
___-2____
d. Iron (III)
___+3____
j. Selenium
___-2____
82. 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
83. Please draw the transfer of electrons for the
following:
a. NaBr
b. MgO
c.
CaF2
d. K2S
Write the formula for the following ionic compounds:
84. Barium oxide
85. Lithium nitride
86. Magnesium fluoride
87. Lithium phosphate
88. Calcium iodide
89. Copper (II) nitrite
90. Beryllium chloride
91. Gold (III) carbonate
92. Potassium sulfate
93. Calcium hydroxide
94. Barium chloride
95. Lithium oxide
96. Magnesium oxide
97. Calcium fluoride
98. Iron (II) chloride
99. Silver bromide
BaO
Li3N
MgF2
Li3PO4
CaI2
Cu(NO2)2
BeCl2
Au2(CO3)3
K2SO4
Ca(OH)2
BaCl2
Li2O
MgO
CaF2
FeCl2
AgBr
Write the name for the following ionic compounds. Be sure to include parenthesis with Roman numeral for
the transition metals:
100.
101.
102.
103.
104.
105.
106.
107.
108.
109.
110.
111.
112.
113.
114.
115.
AlBr3
Fe2(SO3)3
MgCl2
K3PO4
TiCl2
Cr2O3
Ag2S
Cu(OH)2
Zn3N2
AgNO3
PbF2
Cu2S
CaF2
FeCl3
Li3P
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