Golden Valley HS • Chemistry
Name _______________________________
Period: ______________
Final Exam Review – Semester 1
Units 1 (I&E), 2 (Atomic Structure) & 3 (Electrons & Periodicity)
1. What makes one element different from another?
2. What is an isotope?
3. How can you tell the metals from the nonmetals on the periodic table?
4. Give the group number for the
halogens: ______ alkali metals: ______ noble gases: _____ alkali earth metals: _______
5. Which group of elements has two electrons available for bonding? _____________________________
6. Elements like Na, K and Cs are arranged in the same vertical column because they have what
characteristic in common?
7. What does the atomic number represent? ___________ What does the atomic mass represent?
__________
8. Is the periodic table arranged by increasing atomic mass or by increasing atomic number? _______
9. If an element has an atomic number of 24, which element am I referring to? ________
10. What is the atomic number of potassium? ___________ What is the atomic mass of phosphorus?
_________
11. What type of elements mostly make up the periodic table (hint: “solids” is not the answer) _______
12. An element has an atomic mass of 62 and an atomic number 42. How many protons, neutrons and
electrons are there? _______________________________________
13. A new element is discovered and it has 115 protons and 135 neutrons. What is its mass?
______________
14. What has to be true about the protons and the electrons on an atom that is neutral? ____________
15. What is the charge of a proton? __________ of a neutron? ____________ of an electron? ______
16. Where are protons found in an atom? _____________ Where are the neutrons located? ____________
17. Where are electrons found in atoms? ___________________________
18. What force holds the nucleus together?
____________________________________
19. What is the smallest subatomic particle? ____________________________________
20. What is the charge of a nucleus in an atom?
____________________________________
21. Does the nucleus of an atom have a high density or a low density?
______________________________
22. Roughly, what percentage of an atom’s mass is made up of the nucleus?______________________
23. What part of an atom makes up most its volume: the nucleus or the electron cloud?
24. Name the three types of radiation: ___________________________________________________
25. Of the three types, which is the most dangerous? What do you need to stop it? _____________________
26. Which type of radiation is the heaviest? ______________________________________________
S. Cool, 2010
27. What is the charge of an alpha particle? ________ beta particle? _________ gamma particle? ____
28. What materials are needed to stop beta radiation? ____________________________________
29. Complete the following nuclear decay equations:
a.

+ ______________
b.
+

________________
30. Which of the following correctly represents the alpha decay of Thorium-234?
c.

+
d.

+
e.

+
31. What types of reactions release the greatest amount of energy: combustion, decay, fission or
synthesis?
32. Which element on the periodic table has the highest electronegativity? ______________
33. You have three trends to remember, atomic size, ionization energy and electronegativity.
Atomic Size
Ionization Energy
Electronegativity
left to right
top to bottom
34. What is the greatest source of unavoidable error?
35. True or false: Scientific theories are useful so scientists do not have to repeat experiments.
36. What is it called when ideas have been widely demonstrated over time to be consistent with known
scientific observations?
37. True or false: Scientific theories are useful because they can be used to help predict future events.
38. What is the difference between a hypothesis and a theory?
39. Which of the following is an example of a hypothesis:
a. Water the grass.
b. Grass requires light and water to survive.
c. If I cover the grass with black plastic, then it will die.
S. Cool, 2010
Golden Valley HS • Chemistry
Name _______________________________
Period: ______________
Final Exam Review – Semester 1
Units 4 (Bonding), 6 (The Mole) & 7 (Balancing Eqns)
1. ___ Mg(OH)2(aq) + ___ H3PO4
2.
(aq)
 ___ H2O
(l)
+ ___ Mg3(PO4)2(aq)
___ HNO3(aq) + ___ Ni(s)  ___ Ni(NO3)2(aq) + ___ H2(g)
3. ___ C5H12(l) + ___ O2(g)  ___ H2O
(g)
+ ___ CO2(g)
4. ___ C5H10 (l) + ___ O2(g)  ___ H2O
(g)
+ ___ CO2(g)
5.
C3H8 + O2  CO2 + H2O
This chemical equation represents the combustion of propane. When correctly balanced, the coefficient for
water is
a. 2
b. 4
c. 8
d. 16
6. Which of the following is a balanced equation for the combustion of ethanol (CH3CH2OH)?
a. CH3CH2OH + 3O2  CO2 + 2H2O
c. CH3CH2OH + O2  2CO2 + 3HO
b. CH3CH2OH + 3O2  2CO2 + 3H2O
d. CH3CH2OH + 2O2  3CO2 + 2H2O
7. How many atoms are contained in 28 g of nitrogen (N)?
a. 6.0 x 1023
b. 3.0 x 1023
c. 1.2 x 1024
d. 6.0 x 1024
8. Find the mass, in grams, of 3.0 x 1023 molecules of neon (Ne).
9. What is the mass of 4.0 mol of calcium (Ca)?
10. How many liters are in 2.50 mol CO2 at STP?
11. How many atoms are in a chromium sample with a mass of 13 grams?
a. 1.5 x 1023
c. 1.9 x 1026
23
b. 3.3 x 10
d. 2.4 x 1024
12. How many moles of chlorine gas are contained in 9.02 x 1023 molecules?
a. 1.5 moles
c. 6.02
b. 2.0 moles
d. 9.03
13. How many moles of CH4 are contained in 96.0 grams of CH4?
a. 3.00 moles
c. 12.0
b. 6.00 moles
d. 16.0
moles
moles
moles
moles
14. Find the number of moles of argon in 600 g of argon.
15. How many moles of carbon-12 are contained in exactly 6 grams of carbon-12
a. 0.5 mole
c. 3.01x1023 moles
b. 2.0 moles
d. 6.02 x 1023 moles
16. Estimate the molar mass for the following substances:
(NH4)2CO3 ___________
Fe2(SO4)3 ___________ AuCl3 ___________
17. What is the number of moles in 500 L of He gas at STP?
S. Cool, 2010
20. How many valence electrons are in Group 2A? ____ Group 6A? ____ in the halogens? ____ in the alkali
metals? _______
21. **When cations and anions join, they form what kind of chemical bond?
a. ionic
b. hydrogen
c. metallic
d. covalent
22. An ionic bond is a bond between ____________ A covalent bond is a bond between __________________
23. Create an ionic compound out of the following pairs of ions:
a. Li+ and O-2 _____________
c. Sn+4 and N3- _____________
b. Au+3 and S-2 _____________
d. Cu2+ and O2- _____________
24. What are the charges of ions in Group 1A? ________
Group 6A? _________ Group 7A ________
25. What particle is free to drift in metals?__________________________________________________
26. What is the basis of a metallic bond?_____________________________________________________
27. Circle the formulas that represent molecular compounds and underline the ionic compounds:
Kr BaI2 NaCl N2O4 CaO NH3 Li2O HF Al2S3 ZnO Xe Mg3N2 SO2 BeF2
28. How can you tell the difference between a molecular compound and an ionic compound?
29. Name the seven diatomic molecules____________________________________________________________
30. **Some of the organic molecules found in the human body are NH2CH2COOH (glycine), C6H12O6 (glucose), and
CH3(CH2)16COOH (stearic acid). The bonds they form are
a. nuclear
b. metallic
c. ionic
d. covalent
31. What is the charge of a particle having 9 protons and 10 electrons? ______
32. What force holds together ionic compounds? ____________________________________
33. How many electrons do the following atoms give up when they become an ion? Ga ____ H ____ Ca ____
34. How many electrons do the following atoms gain when they become an ion? P ______ Cl ______
S _____
35. What is the formula of the ion formed when potassium achieves the noble-gas electron configuration? ____
36. How many electrons are transferred from calcium to iodine when forming calcium iodide, (CaI 2)? _________
37. What is the charge on the cation in the ionic compound sodium sulfide (Na 2S)?
_______
38. What is the net charge of the ionic compound calcium fluoride? ___________
39. What is the name given to the electrons in the highest occupied energy level of an atom? ___________
40. Draw the Lewis dot structures for the following molecules:
H2
H2S
NH3
H2O
HF
N2
41. Which elements can form diatomic molecules joined by a single covalent bond? ______________
42. Which elements can form diatomic molecules joined by a double covalent bond? _____________
43. Which elements can form diatomic molecules joined by a triple covalent bond? _______________
S. Cool, 2008, 2010
** questions taken from www.cde.ca.gov