Chemistry: Unit 2 Exam Study Guide Answers

1. What is the term for a row on the periodic table?_____period______

2. What is the term for a column on the periodic table?_____group (or family)______

3. State the “Periodic Law” and describe how it relates to the properties of elements in the same period compared to the properties of elements in the same group on the periodic table. The Periodic Law states that the properties of the elements will vary going across a period in a pattern that repeats in the following periods. This means that properties change going left to right across a period, but elements in the same group will have similar properties.

  1. Which IUPAC group numbers coincide with the following group names?

a. alkali metals1

b. halogens17

c. alkaline earth metals2

  1. What is the name of the group to which the following elements belong?

a. Strontium (alkaline earth metals)e. Lithium (alkali metals)

b. Bromine (halogens)f. Calcium (alkaline earth metals)

c. Krypton (noble gases)g. Fluorine (halogens)

d. Silver (transition metals)h. titanium (transition metals)

  1. Are each of the following considered metals, nonmetals or metalloids?

  1. Pb (metal)
  2. Si (metalloid)
  3. Os (metal)
  4. H (metal)
  5. Rb (metal)
  6. Se (non-metal)
  7. He (non-metal)
  8. K (metal)
  9. Te (metalloid)
  10. Ar (non-metal)

  1. For the following elements, list whether they are a metal, nonmetal or metalloid. Then write the name(not number) of the group to which they belong.

ElementMetal, Nonmetal, or MetalloidGroup Name

a. Calcium (Ca)______metal______alkaline earth metal

b. Argon (Ar) ___non-metal______noble gas______

c. Bromine (Br)_____non-metal______halogen_____

d. Cesium (Cs)______metal______alkali metal

  1. Which scientist reasoned that electrons must be part of all elements? How did he reach this conclusion?
  2. JJ Thomson reached this conclusion because he saw the same results (the same charge to mass ratio and deflection) regardless of the gas or metals used.
  1. Draw and label JJ Thomson’s experimental setup for the cathode ray tube experiments. (See diagram to the right)
  1. What were Thomson’s 2 conclusions? What observations supported each of these conclusions?
  2. Electrons are negatively charged
  3. The cathode ray is attracted to the + plate of the electric field or + pole of a magnet and it was deflected by the negative plate or pole of the field or magnet.
  4. Electrons are part of all atoms
  5. The same observations were made regardless of the gas or metals used in the experiment.
  1. What are the locations, charges, mass #s and symbols of each subatomic particle?

Subatomic particle / Location / Charge / Mass # / Symbols
Proton / nucleus / +1 / 1 / p+, 11H
Neutron / nucleus / Neutral / 1 / n0, 11H
Electon / Outside nucleus / -1 / 0 / e-, 0-1e
  1. Draw and label the experimental set-up for Rutherford’s Gold Foil Experiment. Be sure to include particle paths.

  1. What were the conclusions Rutherford drew from the Gold Foil? What experimental observations supported each conclusion?
  2. Atoms are mostly empty space
  3. Most particles went straight through the foil
  4. Atoms have a small, dense, positive nucleus
  5. A few positively charged alpha particles were reflected by the small positive nucleus
  6. Atoms of the same element have equal amounts of which subatomic particle?
  7. protons
  1. Atoms with equal numbers of protons, but different numbers of neutrons are called____isotopes______.
  1. If an atom is neutral, the number of protons equals the number of ______electrons______.
  2. The mass number is equal to the number of ___protons______plus the number of ____neutrons______.
  3. How do you determine the atomic number of an element? It equals the # of protons
  1. Circle the elements that have the same number of neutrons as boron-11.

***Boron-11 has 6 neutrons

f.

  1. Carbon-12
  2. Nitrogen-13
  1. Are and isotopes of the same element? How do you know?

They are the same element because the number of protons is equal. (same atomic #)

  1. Define “atomic mass units.” When do we use this unit?

An atomic mass unit = 1/12 the mass of . We use amu for atomic mass or average atomic weight.

  1. What is the average atomic weight (atomic mass) of an element? How do you calculate it?

The atomic weight of an element is a weighted average of the masses of each isotope of the element. The atomic weight takes into account the mass and relative abundance of each isotope of an element.

Avg atomic wt = (% as decimal of isotope 1)(mass of isotope 1) + (% as decimal of isotope 2)(mass of isotope 2) + etc.

  1. You have a sample of two isotopes, 70.2% Cl-35 and 29.8% Cl-37, what is the atomic weight of this sample?

35.60 amu

  1. Boron has two naturally occurring isotopes, B-10 and B-11, the atomic weight of Boron is 10.81amu, which of these isotopes is more abundant?

Boron-11 is more abundant (Note that the average atomic weight is closer to 11 than 10.)

  1. A sample of Carbon is found to contain 91% Carbon-12, 7% Carbon-13, and 2% Carbon-14.
  2. What would you predict to be the average atomic mass? _____carbon-12______
  3. Calculate the average atomic mass (include units on your answer)

12.11 amu

  1. Was your prediction correct?

Yes!!!

  1. Fill in the following chart (assume all atoms are neutral)

Isotope Symbol / Isotope Name / Mass # / # of electrons / Atomic Number / # of neutrons / # of protons
/ carbon-14 / 14 / 6 / 6 / 8 / 6
/ phosphorous-32 / 32 / 15 / 15 / 17 / 15
/ argon-39 / 39 / 18 / 18 / 21 / 39
/ gallium-74 / 74 / 31 / 31 / 43 / 31

Part II: Quantum Mechanics and the Bohr Model

1. Where are electrons located according to the Bohr model? Describe their movement.

Electrons are located outside the nucleus and travel in circular orbits around the nucleus according to the Bohr model. An electron in a specific orbit has a specific amount of energy. The energy of an electron increases as it goes from the inner orbit (ground state) to outer orbits.

2. Describe what happens to the energy of an electron as it moves closer to and farther away from the nucleus of an atom.

Energy increases (or is gained) as an electron moves farther away from the nucleus to higher energy levels. When an electron moves to a lower energy level closer to the nucleus it loses energy as a photon of electromagnetic radiation.

3. Why do we no longer use the Bohr model to describe the behavior of atoms?

The Bohr model only worked for Hydrogen. It did not take into account the fact that electrons do not behave (move) in the same fashion as macroscopic objects. Electrons do not move in circular orbits in the atom.

4. The Heisenberg uncertainty principle tells us it is impossible to determine both the position and the momentum of a very small particle simultaneously.

5. How many electrons are allowed in the same orbital?2

6. What must be true about the electrons in the same orbital?Opposite spins

7. How would you place 4 electrons into the orbitals of the 2p sublevel?

Place one electron into each of the three orbitals before adding the fourth electron to the first of the three 2p orbitals.

8. The principal quantum number is equal to the number of __sublevels in the principal energy level____.

9. Describe the shapes of both s and p orbitals.

s-orbitals are spherical while p-obitals are dumbbell shaped

10. What is the maximum number of electrons allowed in the p orbitals in a principal energy level?6

11. What is the maximum number of orbitals in the d sublevel?5

12. What is the maximum amount of electrons allowed in the first principal energy level?2

13. Write the orbital diagrams (with arrows) for the following elements:

a. Chlorine

b. Copper

14. Write full electron configurations for the following elements:

a. Sc1s22s22p63s23p64s23d1(1 unpaired electron)

b. P1s22s22p63s23p3(3 unpaired electrons)

c. Be1s22s2(0 unpaired electrons)

d. C1s22s22p2(2 unpaired electrons)

e. Y1s22s22p63s23p64s23d105s24d1(1 unpaired electron)

f. H1s1(1 unpaired electron)

15. Write short hand (noble gas) electron configurations for the following elements:

a. Ni[Ar]4s23d8

b. Cs[Xe]6s1

c. I[Kr]5s24d105p5

d. Ag[Kr]5s24d9

e. Ba[Xe]6s2

f. S[Ne]3s23p4

Part IV: Periodic Trends

5. Fill in the blanks with either increases or decreases.

a. Atomic radiusincreases down a group

b. Atomic radius decreases across a period.

c. Electronegativity decreases down a group.

d. Electronegativity increases across a period.

6. List each of the following in order of increasing atomic size.

a. Aluminum, Argon, Barium (smallest) Ar, Al, Ba (largest)

b. Neon, Carbon, Radium  (smallest) Ne, C, Ra (largest)

7. List each of the following in order of increasing electronegativity.

a. Copper, Zinc, Fluorine  (smallest) Cu, Zn, F (largest)

b. Silver, Gold, Copper  (smallest) Au, Ag, Cu (largest)