Unit 2 – Atomic Theory and Periodicity Review

Section I: History

In each box, write the name of the scientist(s) associated with the statement. Choose from among the following:

  • Democritus
  • Dalton
  • Bohr
  • Rutherford
  • Thomson
  • Schroedinger and Heisenberg

There are small negatively charged particles inside an atom
Thomson / His discovery was made after conducting an experiment with gold foil
Rutherford
There is a small, dense, positively charged nucleus
Rutherford / Atoms are small, hard spheres
Dalton
Most of an atoms mass is in the nucleus
Rutherford / Electrons are found in electron clouds, not in defined paths
Schroedinger and Heisenberg
Electrons follow a definite path but can jump from one path to another
Bohr / Elements combine is specific proportions to make compounds
Dalton
Atoms are mostly empty space
Rutherford / Atoms are uncuttable
Democritus
Atoms of one element are all the same, but atoms of different elements are different
Dalton / His theory of atomic structure led to the “plum pudding” model of atoms
Thomson
Electron paths cannot be defined for certain
Schroedinger and Heisenberg / All substances are made of atoms
Dalton

Section II: Atomic Vocabulary (unscramble)

  1. Weighted average of all naturally occurring isotopes of the same element. (mictoasams)atomic mass
  2. The building blocks of matter (moats)atoms
  3. Positively charged particle in an atom (torpno)proton
  4. Made up of protons and neutrons (ucselun)nucleus
  5. Particle in an atom that has no charge (tronune)neutron
  6. Atoms with the same number of protons but a different number of neutrons (sootpies)isotopes
  7. Negatively charged particle in an atom (cleenrot)electron
  8. Number of protons in a nucleus (mictoabrumen)atomicnumber
  9. Regions where electrons are likely to be found (renectolscudlo)electronclouds
  10. Sum of protons and neutrons (samsbrumen)massnumber

Section III: Isotopes Practice

Complete the table below.

Complete atomic symbol / Atomic # / Mass # / # of protons / # of neutrons / # of electrons
C / 6 / 14 / 6 / 8 / 6
I / 53 / 78 / 53 / 25 / 53
Cl / 17 / 35 / 17 / 18 / 17
Fe / 26 / 54 / 26 / 28 / 26
He / 2 / 4 / 2 / 2 / 2
U / 92 / 238 / 92 / 146 / 92
  1. Name the element which has the following number of particles:
  1. 82 electrons, 125 neutrons, 82 protons Lead-207
  2. 53 protons, 53 electrons, 74 neutrons Iodine-127
  1. Naturally occurring europium consists of two isotopes with masses of 151 and 153 amu. The respective abundances are 48.03% and 51.97%. What is the atomic mass of europium?

152.04 amu (or 154.02 g/mol)

  1. Strontium consists of four isotopes. There masses and abundances are listed below. Use this data to calculate the atomic mass of strontium.

MassAbundance

840.50%

869.9%87.71 amu (or 87.71 g/mol)

877.0%

8882.6%

Section IV: Nuclear Reactions

Write the equations for the reactions described below.

  1. Decay of polonium-218 by alpha (α) emission. Po  He +Pb
  1. Decay of carbon-14 by beta (β) emission. C β +N
  1. The alpha decay of radon-198 Po  He +Po
  1. The beta decay of uranium-237 Po  β +Np

Section V: Unit Conversions (Mass-Moles-Atoms)

  1. How many moles are in 4.14x1022 atoms of boron?

0.0687mol

  1. Determine the mass in grams of 6.8 moles of iron.

380 g

  1. What is the mass of 1.62 x 1023atoms of carbon?

3.23 g

  1. How many atoms are in 2.17 grams of zinc?

2.00 x 1022 atoms

Section VI: Regions of the periodic table

  1. Name the following regions of the periodic table.

  1. Group IA alkalimetals
  2. Group IIA alkalineearthmetals
  3. Group VIIA halogens
  4. Group VIIIA noblegases
  5. Groups IB – VIIIB transitionmetals
  6. The top row of the f block lanthanides
  7. The bottom row of the f block actinides

  1. List the six metalloids (aka semimetals). Boron, silicon, germanium, arsenic, antimony, & tellurium
  2. List the seven diatomic elements. hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, iodine

Section VII: Electron configurations, Orbital diagrams, and Lewis dot diagrams

Write the electron configurations and draw the corresponding orbital diagrams and the Lewis dot diagram for the elements below.

  1. Hydrogen1s1
  2. Boron1s2 2s2 2p1
  3. Sodium1s2 2s2 2p6 3s1
  4. Krypton1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
  5. Chromium1s2 2s2 2p6 3s2 3p6 4s2 3d4
  6. Phosphorus1s2 2s2 2p6 3s2 3p3
  7. Carbon1s2 2s2 2p2
  8. Oxygen1s2 2s2 2p4
  9. Potassium1s2 2s2 2p6 3s2 3p6 4s1
  10. Cobalt1s2 2s2 2p6 3s2 3p6 4s2 3d7
  11. Platinum 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d8

Write the abbreviated electron configurations and orbital diagrams for the elements below.

  1. Platinum [Xe] 6s2 4f14 5d8
  2. Plutonium[Rn]7s2 5f6
  3. Neodymium[Xe]6s2 4f4
  4. Lead[Xe] 6s2 4f14 5d10 6p2
  5. Cesium[Xe]6s1

Describe each of the following rules for electrons filling orbitals in an electron cloud.

  1. Aufbau rule – each electron occupies the lowest energy orbital available
  2. Pauli exclusion principle – a max of two electrons may occupy one orbital but the must have opposite spins
  3. Hund’s rule – single electrons must occupy each equal-energy orbital before an opposite spin electron is added
  4. Heisenberg uncertainty principle – it is impossible to know both the velocity and position of an electron at the same time

Section VIII: Periodic Trends

  1. Rank the following elements by increasing atomic radius: carbon,aluminum, oxygen, potassium.

O, C, Al, K

  1. Rank the following elements by increasing electronegativity: sulfur,oxygen, neon, aluminum.

Ne, Al, S, O

  1. Rank the following elements by increasing first ionization energy: bromine, strontium, arsenic, calcium

Sr, Ca, As, Br

  1. Why does fluorine have a higher first ionization energy than iodine? Fluorine’s valence shell is closer to the nucleus that iodine’s (fluorine has fewer shielding electrons). Thus, the nucleus has a stronger attraction for fluorine’s valence electrons, making them harder to remove.