Chemistry Name __________________________

Hour _______

Chemistry

Approximate Timeline

Students are expected to keep up with class work when absent.

CHAPTER 11 – MODERN ATOMIC THEORY

Day Plans for the day Assignment(s) for the day

1 Begin Chapter 11

11.1 – Atoms & Energy

o Rutherford’s Atom

o Energy & Light

o Emission of Energy by Atoms

o Flame Tests (demo)

11.2 – The Hydrogen Atom

o Energy Levels of Hydrogen

o The Bohr Model

o The Wave Mechanical Model

Assignment 11.0

Assignment 11.1

Read section(s) 11.3

2 Section 11.3 – Atomic Orbitals

o The four quantum numbers

Assignment 11.2

Assignment 11.3

Read section(s) 11.4

3 11.4 – Electron Configurations & Atomic

Properties

o Electron Configurations (full)

Assignment 11.4a

4 11.4 – Electron Configurations & Atomic

Properties

Electron Configurations (Noble Gas

shortcut)

Assignment 11.4b

5 11.4 – Electron Configurations & Atomic

Properties

o Atomic Size

o Ionization Energy

Assignment 11.4c

6 Work on Assignment 11.5 Assignment 11.5

7 Grade & discuss Assignment 11.5

Review for Chapter 11 Test

8 Chapter 11 Test Read section(s) 12.1

Chemistry Name __________________________

Hour _______

Study Guides

Chapter 11 Quizzes

Quiz 11.1 Atoms & Energy 1. Be able to describe Rutherford’s model of the atom.

2. Know the definition of electromagnetic radiation.

3. What is the relationship between the wavelength of light and the energy carried by the

light wave?

4. Think of a rainbow and answer the following.

a. Which color of light has the longest wavelength?

b. Which color of light has the highest energy?

c. Which color of light has the shortest wavelength?

d. Which color of light has the lowest energy?

Quiz 11.2 The Hydrogen Atom 5. Combine knowledge from sections 11.1 and 11.2 to match the color of light with the

amount of energy carried by specific colors.

6. Define the term “quantized”.

Quiz 11.3 Atomic Orbitals 7. Know the number of sublevels in each of the first four energy levels.

8. Define the term “orbital”.

9. Know the number of orbitals within each sublevel.

10. Know the maximum number of electrons that a single orbital can contain.

11. State the “Pauli Exclusion Prinicple”.

Quiz 11.4 Electron Configurations & Atomic Properties 12. Define the term “electron configuration”.

13. Understand each “part” of an electron configuration.

a. What does the coefficient represent?

b. What does the letter represent?

c. What does the exponent represent?

14. Define the term “valence electron”.

15. Define the term “core electron”.

16. Positive ions are formed when an atom ___ electrons.

17. Negative ions are formed when an atom ___ electrons.

18. Know the periodic trends regarding atomic size and the periodic trend regarding

ionization energy.

Study Guide

Chapter 11 Test

At the completion of chapter 11 you should…

1. Know the definitions of the following terms.

a. Electromagnetic Radiation

b. Photons

c. Quantized

d. Wave Mechanical Model

e. Orbital

f. Principle Energy Level

g. Sublevel

h. Electron Configuration

i. Valance Electrons

j. Core Electrons

k. Ionization Energy

2. Be able to describe Rutherford’s model of the atom.

3. Understand how the light emitted by atoms can be used to identify atoms.

4. Be able to describe Bohr’s model of the atom

5. Understand how the electron’s position is represented in the wave mechanical model

6. Know the four quantum numbers

a. the symbol (the letter)

b. what it represents

c. the theoretical range of values

d. the actual range of values

7. Be able to write electron configuration

a. The full electron configuration

b. Use the noble gas shortcut

8. Be able to describe and explain periodic trends in properties

a. Atomic size

b. Ionization energy

Chemistry Name __________________________

Hour _______

Assignment 11.0 – Vocabulary

Define each of the following terms.

1. Electromagnetic Radiation

2. Photons

3. Quantized

4. Wave Mechanical Model

5. Orbital

6. Principle Energy Levels

7. Sublevels

8. Electron Configuration

9. Valance Electrons

10. Core Electrons

11. Ionization Energy

Chemistry Name __________________________

Hour _______

Assignment 11.1 – Atoms & Energy

1) What is wrong with Rutherford’s model of the atom? Why did it need to be modified?

2) What is electromagnetic radiation? Provide three examples.

3) How is the frequency of a wave different from its speed?

4) What is a photon?

5) What is the relationship between wavelength of light and the energy of its photons?

Chemistry Name __________________________

Hour _______

Assignment 11.2 – The Hydrogen Atom

1) In the hydrogen spectrum (Figure 11.8, page 367) there are four different colored lines.

You already know that hydrogen has only one electron. How can we get four lines from

one electron?

2) Explain the terms ground state and excited state. When a photon of energy is absorbed by

an atom, does the electron go from the ground state to an excited state or from an excited

state to the ground state?

3) What is wrong with the Bohr model of the atom?

4) How does the wave mechanical model of the atom differ from Bohr’s model?

Chemistry Name __________________________

Hour _______

Assignment 11.3 – Atomic Orbitals

1) What is the difference between an orbit and an orbital in atomic theory?

2) Draw figure 11.14 (page 371) and fill in the different types of sublevels (s, p, d, and f) for

each principle energy level.

3) Tell how many orbitals are found in each type of sublevel: s, p, d, f.

4) What is the Pauli Exclusion Principle and how does it help us determine where an

electron is within the atom?

Chemistry Name __________________________

Hour _______

Assignment 11.4a – Electron Configurations (1)

Write the full electron configurations for the following elements.

1) nitrogen

2) silicon

3) potassium

4) iron

5) silver

6) barium

7) europium

8) platinum

9) plutonium

10) einsteinium

Give the name of the element that has the following electron configuration.

11) 1s2

2s2

2p6

3s2

3p4

12) 1s2

2s2

2p6

3s2

3p6

4s2

3d10

4p6

5s2

4d10

5p6

6s2

4f14

5d10

13) 1s2

2s2

2p6

3s2

3p6

4s2

3d10

4p6

5s2

4d2

14) 1s2

2s2

2p6

3s2

3p6

4s2

3d10

4p6

5s2

4d10

5p6

6s2

4f14

5d10

6p6

7s1

Chemistry Name __________________________

Hour _______

Assignment 11.4b – Electron Configurations (2)

Write the electron configuration for each of the following elements. You MUST use the noble

gas shortcut.

1) Magnesium [

2) Titanium

3) Germanium

4) Rubidium

5) Rhenium

6) Indium

7) Osmium

8) Dysprosium

9) Uranium

10) Nobelium

Give the name of the element that has the following electron configuration.

11) [Ne] 3s23p

5

12) [Ar] 4s23d

8

13) [Kr] 5s24d

105p

6

14) [Xe] 6s24f

5

15) [Xe] 6s24f

145d

2

16) [Rn] 7s25f

7

Chemistry Name __________________________

Hour _______

Assignment 11.4c – Electron Configurations & Atomic Properties (2 pages)

1. A) What is the difference between a valence electron and a core electron? B) Write the

electron configuration for Fe. Label both the valance electrons and the core electron

electrons.

2. Elements in groups (vertical columns) show similar chemical behavior. Why?

3. What chemical properties distinguish metals from nonmetals?

4. Explain the general trend in atomic size across rows and down columns of the periodic

table.

5. Circle the larger atom in each pair.

a. Na K

b. Na Mg

c. Fe Zn

d. O Se

6. Explain the general trend in ionization energy across rows and down columns of the

periodic table.

7. Circle the atom with the larger ionization energy in each pair.

a. Na K

b. Na Mg

c. Fe Zn

d. O Se

Chemistry Name __________________________

Hour _______

Assignment 11.5 – Chapter Review (4 pages)

1) A) How are the different types of electromagnetic radiation similar? B) How do they

differ?

2) What is a “packet” of electromagnetic energy called?

3) Do the colors of flame tests result from taking in energy or releasing energy?

4) What does it mean to say that an atom is in an “excited state”?

5) When an atom in an excited state returns to the ground state, what happens to the excess

energy of the atom?

6) What is meant by the ground state of an atom?

7) According to Bohr, what types of motions do electrons have in an atom, and what

happens when energy is applied to the atom?

8) Discuss briefly the difference between an orbit (as described by Bohr for hydrogen) and

an orbital (as described by the more modern, wave mechanical picture of an atom.)

9) A) Which orbital is the first to be filled in any atom? B) Why?

10) A) Which electrons of an atom are the valence electrons? B) Why are these electrons

especially important?

11) A) How are electron arrangements in a given group (vertical column) of the periodic

table related? B) How is this relationship shown in the properties of the elements in a

given group?

12) Write the full electron configuration (1s2 2s

2 etc.) for each of the following elements.

How many valence electrons does each atom possess?

A) strontium

B) zinc

C) helium

D) bromine

E) calcium

F) potassium

G) fluorine

H) krypton

Chemistry Name __________________________

Hour _______

13) Why do we believe that the valence electrons of calcium and potassium reside in the 4s

sublevel rather than the 3d sublevel?

14) Using the noble gas shortcut, write the electron configuration for each of the following

elements. How many valence electrons does each atom possess?

A) calcium

B) francium

C) yttrium

D) cerium

E) phosphorus

F) chlorine

G) magnesium

H) zinc

15) A) What types of ions do the metals and nonmetals form? B) Do the metals lose or gain

electrons in doing this? C) Do the nonmetals gain or lose electrons in doing this?

16) Why do the metallic elements of a given period (horizontal row) typically have much

lower ionization energies that do the nonmetallic elements of the same period?

17) Explain why the atoms of the elements at the bottom of a give group (vertical column) of

the periodic table are larger than atoms at the top of the same group.

18) In each of the following sets of elements, which element has the lowest ionization

energy? Which element has the highest ionization energy?

Lowest Highest

A) Cs, Rb, Na ____________ ____________

B) Ba, Ca, Be ____________ ____________

C) F, Cl, Br ____________ ____________

D) O, S, Te ____________ ____________

19) Arrange the following sets of elements in order of increasing atomic size.

A) Sn, Xe, Rb, Sr

B) Rn, He, Xe, Kr

C) Pb, Ba, Cs, At