Chemistry 1
West Linn High School
Unit 3 Packet and Goals
Name:_________________________________
Period:_________
Unit 3 – Electrons and the Periodic Table
Unit Goals:
As you work through this unit, you should be able to:
1.
Understand the modern theory of atomic structure from Bohr's model of the atom to the quantum model of the atom.
(5.1)
2. Understand the difference in energy of the various atomic orbitals. (5.1)
3. Understand how to write electron configurations and fill in orbital diagrams for various atoms and ions using the periodic
table. (5.2 and 7.1))
4. Define Valence electrons and know the number of valence electrons and common charges of the representative elements
and their ions based on the position on the periodic table. (7.1)
5. Understand how atomic emission spectra relate to changes in electron orbital energy. (5.3)
6. Describe the history of the development of the periodic table. (6.1)
7. Understand how the periodic table is organized using periodic law. (6.1)
8. Identify metals, non-metals and metalloids based on their location on the periodic table. (6.1)
9. Identify the position of elements in periods and families on the periodic table. (6.2)
10. Describe trends of similar properties within elements of the same family on the periodic table. (6.2)
Reading:
Chapter 5: (pp 128-150) Sections 5.1 – 5.3
Chapter 6: (pp 160-182) Sections 6.1-6.3
Key Terms:
ground state, excited state, quantum, quantum mechanical model, orbitals, Pauli's Exclusion Principle, Hund's
Rule, Aufbau Principle, electron configuration, orbital diagrams, electron configuration, periodic table, periodic
law, representative elements, period, group, metals, non-metals, alkali metals, alkaline earth metals,
transition metals, halogens, noble gases, metalloids, atomic size, ionization energy, electronegativity
Classwork and Labs
Homework:
HW 1
Description
Atomic theory and Orbitals
Worksheet
Goals
1-4
Score
Unit 3 Note Packet
Flame Test Lab
HW 2
Periodic Table WS
5-7
Periodic Trends Activity
Unit 3 Exam (goals 1-5)
HW 3
Periodic Table and Trends WS
8-9
Chapter 5: Goals 1-3
What are the chemical properties of atoms and molecules related to?
5.1. Revising the Atomic Model
Electron/Atomic History:
Dalton:
Thomson:
Rutherford
Energy Levels in Atoms: Page 128
Bohr:
• Electron’s are found in Energy levels.
• Quantum: Specific energy values that have exact values in between them.
• Pictured the Atom like a solar system with electrons “Orbiting” around the nucleus.
•
•
Each Orbit path associated with a particular energy.
Now better understand how electrons are organized, but it still isn’t quite correct
The Quantum Mechanical Model: page 130
Erwin Schrodinger (1926) Quantum Mechanical Model
Schrodinger’s equation:
• Electrons are found in regions of empty space around the nucleus. “Electron Clouds”
• Shape of the region corresponds to different energy levels.
• Shape of the region represents the probability of the electron’s location 90% of the time.
• PROBLEM? These are still called “Atomic Orbitals”
Atomic Orbitals: page 131
Table 5.1 (page 132)
Principal Energy
Level
5.2.
Number of
sublevels
Type of sublevel
Electron Arrangement in Atoms
Electron Configurations: Page 134
Aufbau Principle:
Pauli Exclusion Principle:
Hund’s Rule:
Maximum number
of electrons
1. The atomic number of phosphorus is 15. Write the electron configuration of a phosphorus atom.
2.
Give the full & noble gas (abbrev) GROUND STATE electron configuration and underline the “outermost” electrons
for the following elements:
Full
Abbrev
a. C
b. Ar
c. Ni
a. B
b. Si
c. S
3. Why do we care about electron configurations?
4. The full ground state electron configuration for Mo is:
1s22s22p63s23p64s23d104p65s14d10
5. What is “wrong” with Mo’s electron configuration?
6. What is the expected electron configuration of Mo?
7. How many unpaired e-s are in the electron configuration above?
8. Which e-s are most energetic? Least energetic?
9. How could you abbreviate the electron configurationabove?
10. Why would you abbreviate and leave off the “inner-core” electrons?
11. What other elements would you predict have electron configurations like Mo?
12. Which electrons are usually closest to the nucleus in the following:
1s22s22p63s23p64s1
13. Which electrons are furthest from the nucleus?
14. Which electrons would you think would be easiest to remove from the atom above? Why?
15. What name is given to these “outermost” electrons?
Ions and Electrons (page 194-198): Short segment of Chapter 7.
16. Give the full & noble gas (abbrev) GROUND STATE electron configuration for the following ions:
Full
Abbrev
O2Cu1+
N3-
17. What keeps an electron bound to an atom?
18. In the space below, build an atom that contains 3 protons, 4 neutrons and 3 electrons.
19. How many valence electrons are in the atom above?
20. Predict what would happen when the valance (outermost) electron absorbed some energy from an outside
source.
21. What would happen when that electron loses that energy that it absorbed?
22. How are different colors of light formed?
Electromagnetic Radiation
23. Describe how light is created.
24. List the different forms of visible light in order of lowest energy to greatest energy.
25. What color of visible light would be the most appropriate to use if you are star gazing at night and do not want
to disturb other star gazers? Why?
26. Predict the color of visible light produced in the following situations:

A 3s electron jumps to the 4th energy level as it absorbs light, then falls back down to the ground state.

A 3s electron jumps to the 7th energy level as it absorbs light, then falls back down to the ground state.

A 5p electron absorbs energy and jumps to the 6th energy level.
Chapter 6
6.1.
Organizing the Elements: Page 160 Periodic Table History
Mendeleev’s Periodic Table: Page 161
– Proposed a table for 70 elements based on mass and properties
Today’s Periodic Table: Page 162
Henry Moseley:
– Determined the atomic number of elements and arranged the table in order of atomic number
How is the Periodic Table Organized?
 Periods

Groups or Families
Periodic Law
6.2.
Classifying the Elements
Metals
Non-metals
Metalloids
• Transition Metals
• Inner-transition Metals
• Alkali Metals
• Alkaline Earth Metals
• Halogens
• Noble (inert) gases
• Representative Elements
(1A-7A)
Basis of the Law
Chemical Properties and Families
•
Chemical Properties of elements are based on their “
•
Families are groups of Elements that have similar “
•
Valence Electrons
• Outermost electrons in an atom.
6.3.
Periodic Trends
Atomic Radius (Page 174)
• In general, the more VALENCE electrons, the smaller the
size of the atoms electron clouds.
• Exception: Hydrogen and Helium
• However, adding periods increases the size of the atom.
Ion Radius (Page 179)
• - Ions (Anions) have full VALENCE electrons = LARGER than
their ATOM
• + Ions (Cations) have empty VALENCE electrons = SMALLER
than their ATOM
• BUT…. The size of one ion compared to the next is the
same pattern as ATOMIC RADIUS
1st Ionization Energy (Page 177)
• The amount of energy required to remove an electron
from an atom.
• Atoms that tend to LOSE electrons have LOW Ionization
energy.
• Lowest = Francium
• Atoms that do not lose electrons easily have HIGH
Ionization energy.
• Highest = Helium
Electronegativity (Page 181)
• The tendency for an atom to attract electrons of
another atom.
• (-) Anions that gain electrons have high electronegativity
values (they desire more electrons)
• Noble Gasses have very small electronegativities.
• Examples
• Highest = Fluorine
• Lowest = Francium (NON-Noble gas)
• Lowest group = Noble gasses
Journal and Warm-up:
Remember to record; date and question for each warm-up problem. Show all work and edits to correct.
Journal and Warm-up:
Remember to record; date and question for each warm-up problem. Show all work and edits to correct.