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Chemistry 40S Atomic Structure

(This unit has been adapted from https://bblearn.merlin.mb.ca)

Name:

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Lesson 1: The Nature of Light Goals:

Describe light in terms of electromagnetic energy.

Describe the electromagnetic spectrum.

Describe the relationship between frequency, wavelength and energy of light.

Identify an element based on its flame test.

Describe line and continuous spectra.

Wavelength

Frequency

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What relationship do you notice between wavelength and frequency?

Frequency, Wavelength, Energy, and Electromagnetic Spectrum

Max Planck

If we shine sunlight through a prism we get a rainbow of colours, known as a spectrum. Each colour in the rainbow represents light of a different frequency or wavelength.

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The visible spectrum is actually a very small portion of the entire electromagnetic spectrum.

Line Spectra

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Practice: The Nature of Light

1. Put the following in order of increasing energy: green light, x-rays, radio, red light, ultraviolet, microwaves, blue light, gamma rays.

2. Describe the relationship between frequency, wavelength and energy.

Element Colour

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3. Why is light called electromagnetic radiation?

4. Compare and contrast line and continuous spectra.

5. Give the colour for the flame test of each of the following:

a) Copper

b) Strontium

c) Lithium

d) Potassium

e) Barium

f) Calcium

g) Sodium

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Lesson 2: The Quantum Mechanical Model of the Atom Goals:

Explain the development of the Quantum Mechanical Model of the atom.

Explain the formation of line spectra.

The Model of the Atom

John Dalton:

Joseph John Thompson:

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Earnest Rutherford

Explanation of Line Spectra

Neil Bohr

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Quantum Theory

Albert Einstein

Louis de Broglie

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The Quantum Mechanical Model of the Atom

Werner Heisenberg

Erwin Schrödinger

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Orbitals

s-orbitals

p-orbitals

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d-orbitals

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Practice: The Quantum Mechanical Model of the Atom

1. Identify the contribution made by each of the following to atomic theory.

a) Rutherford

b) Heisenberg

c) Planck

d) Bohr

e) de Broglie

f) Schrödinger

2. Describe s, p and d-orbitals. What do they represent? How are they different from each other?

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Lesson 3: Electron Configurations Goals:

Write the electron configuration of atoms and ions.

Write the valence electron configuration of and atom or ion.

The Orbitals

Principle Quantum Number:

Main Energy Level (n)

Number of Orbitals (n2)

Type of Orbitals

The Pauli Exclusion Principle

Wolfgang Pauli

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Principle Quantum Number:

Main Energy Level (n)

Number of Orbitals (n2)

Type of Orbitals Number of Electrons

(2n2)

The Electron Configuration

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Aufbau Principle

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Hund’s Rule

Friedich Hund

Element Filling of Orbitals Electron Configuration 1s 2s 2px 2py 2pz

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Zig-Zag Rule

7s 7p

6s 6p 6d

5s 5p 5d 5f

4s 4p 4d 4f

3s 3p 3d

2s 2p

1s

Example 1: Write the complete electron configuration for magnesium. Use the Zig-Zag Rule to assist you.

Example 2: Write the complete electron configuration for germanium. Use the Zig-Zag Rule to assist you.

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Exceptions to the Rules

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Electron Configurations of Ions

Example 3: Write the complete electron configuration for the chloride ion, Cl-. Example 4: Write the complete electron configuration for the calcium ion, Ca2+.

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Example 5: Write the complete electron configuration for the iron (II) ion, Fe2+.

Example 6: Write the complete electron configuration for the iron (III) ion, Fe3+.

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Valance Configurations

Example 7: Write the valence electron configuration for fluorine. Example 8: Write the valence electron configuration for germanium.

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Practice: Electron Configurations

1. How many electrons in an atom can have the designation?

a) 1s b) 2p

c) 3px d) 6f

e) 3dxy f) n = 2

g) 4p h) n = 5 2. Write complete electronic configurations for the following atoms and

ions:

a) P b) Ca c) Cu d) Rh

e) Sb3+ f) Ni2+ g) Fe2+ h) Ni4+

i) Zn2+ j) Br–

k) Sn2+ l) Co3+

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3. How many unpaired electrons are there in each of the following:

a) Mn b) As

c) Sr d) Tl+ e) Cu2+ f) V3+

g) Sn h) Lu 4. Write the electronic configurations for the valence electrons of each of

the following:

a) Mg b) P c) Se d) Pb2+

e) Br f) S2- g) Ni h) Ag+ i) N3- j) Fe3+

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Lesson 4: Electronegativity Goals:

Use electronegativity values to predict the type of bond between two atoms.

Describe a bond in terms of polar covalent, non-polar covalent and ionic.

Electronegativity

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Polar, Non-Polar, and Ionic Bonds

Electronegativity Difference

Character of Bond Percent Ionic Character

Example 1: What type of bond forms between sodium and chlorine in NaCl? Example 2: What type of bond forms between sulphur and oxygen in SO3?

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Example 3: What type of bond forms between aluminum and chlorine in AlCl3?

Practice: Electronegativity

Answer the following questions. Refer to the Electronegativity Table when needed.

1. Describe the trend of electronegativity on the periodic table.

2. Describe each of the following in terms of electronegativity.

a) non-polar covalent bond

b) polar covalent bond

c) ionic bond

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3. Describe the following as non-polar covalent, polar covalent, or ionic.

a) N and H b) F and F

c) Ca and O d) Al and Br e) H and I f) K and Cl

4. List the following in increasing ionic character.

a) Mg-F, Ca-I, Ca-Cl, Mg-Cl

b) Al-Cl, H-Cl, K-Cl, Cu-Cl

c) C-O, C-H, C-F, C-Br

d) S-Cl, H-C, H-F, H-H, H-Cl, H-O

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Lesson 5: Lewis Dot Structure Goals:

Draw dot diagrams for covalent compounds.

Dot Structures for Atoms and Octet Rule

Gilbert Lewis

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Drawing Lewis Dot Structure

Example 1: Draw the Lewis Dot structure for HF.

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Example 2: Draw the Lewis structure for CO2. Example 3: Draw the Lewis structure for the sulphate ion, SO42-.

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Practice: Lewis Dot Structure

1. Draw the Lewis structure for each of the following.

a) NH3 b) SO3

c) BrF d) O2 e) PO4

3– f) CH4 g) HCN h) N2 i) Cl2O j) ClO3

- k) CO3

2- l) SiF4

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Lesson 6: Ionization Energy and Periodic Trends Goals:

Describe the factors that affect the force on an electron.

Explain the trends in atomic radius, ionic radius and ionization energy.

A Force on an Electron

Atomic Radii

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Atomic Radii (pm) of Group 1 Atom

Ionic Radii

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Periodic Trends and Ionization Energy

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Practice: Ionization Energy and Periodic Trends

1. Briefly explain why barium has a lower first ionization energy than calcium.

2. Given the following elements and their electron configuration. Which element will have the lowest first ionization energy? Why?

Element A 1s22s22p63s23p64s2

Element B 1s22s22p63s23p5

Element C 1s22s22p63s23p64s1

Element D 1s22s22p63s23p6

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3. The first four ionization energies (IE) for the element aluminum are as follows:

IE1 = 577 kJ/mol

IE2 = 1 817 kJ/mol

IE3 = 2 745 kJ/mol

IE4 = 11 580 kJ/mol How many valence electrons does aluminum have?

4. An atom has the following successive ionization energy levels.

IE1 = 737 kJ/mol

IE2 = 1 450 kJ/mol

IE3 = 7 731 kJ/mol

IE4 = 10 545 kJ/mol

IE5 = 13 627 kJ/mol How many valence electrons does this element have? Explain.

5. Which of these elements would have the highest value for the second ionization energy? Why?

K

Si

Ar

Br

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6. Which of the following has the largest atomic radius and which has the smallest? Explain.

Nitrogen

Antimony

Arsenic

7. For each of the following properties:

atomic radius

first ionization energy

ionic radius Indicate which has the larger value fluorine or bromine.

8. Arrange the following from largest to smallest. Explain the order. Ne, Mg2+, F-, Na+, O2-

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