Atomic Theory Unit II

Lesson Date Topic WS #

1. Early Atomic Theory 1

2. Bohr Evidence 1

3. Bohr Diagrams 1

4. Quantum Theory 2

5. Mass Spectrometer/ Elegant Universe-1 3

6. Elegant Universe-2/Periodic Chem 4

7. Periodic Chemistry 5

8. Classifying Matter Lab/ Elegant Universe-3

9. Classifying and Naming Formulas 1 6

10. Classifying and Naming Formulas 2 7

11. Electron Dot Diagram Structural Formula 1 8

12. Electron Dot Diagram 2 9

13. Practice Test 1 10

14. Practice Test 2 11

Worksheet # 1 Early Atomic Theory

Briefly describe each atomic theory listed below. Include a diagram.

1. The Four-Element Theory

2. Dalton’s Atomic Theory

3. The Thompson Atom

4. The Rutherford Atom

5. The Bohr Atom

For each of the above, give some evidence that led to the atomic model. Briefly explain how each evidence is accounted for in the atomic theory.

6. The Four Element Theory

(a) evidence (b) explanation within theory

7. Dalton’s Atomic Theory

(a) evidence (b) explanation within theory

8. The Thompson Atom

(a) evidence (b) explanation within theory

9. The Rutherford Atom

(a) evidence (b) explanation within theory

10. The Bohr Atom

(a) evidence (b) explanation within theory

Draw Bohr atomic diagrams for the following atoms. Be sure to include protons, neutrons and electrons.

11. Oxygen 12. Calcium

13. Silver 14. Barium

15. Cs 16. I

17. Na 18. V

19. Cl^- 20. Al³⁺

21. Se^2- 22. Ca²⁺

Worksheet #2 Quantum Mechanics

1. What is the main difference between the Bohr theory of the atom and the Quantum Mechanical Theory?

2. How many electrons will fill the smallest orbital in quantum mechanical theory?

3. How is a 3s orbital different than a 2s orbital in terms of shape and distance from the nucleus?

4. Explain what happens to the energy when an electron falls from a 3s orbital to a 2s orbital.

Use your Quantum Periodic Table to write quantum electron configurations for each element below.

5. F

6. K

7. C

8. Kr

9. S

10. Rb

11. Co

12. P

13. Ca

14. Al

15. Ag

16. 1s²2s²2p⁶3s¹

17. 1s²2s²2p⁶3s²3p⁵

18. 1s²2s²2p⁶3s²3p⁶3d⁹4s²

19. 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰5s²5p⁵

20. 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶4d¹⁰5s²5p⁶6s²

21. Give the formula of four chemical species that are isoelectronic (same electron configuration) as Ar.

Worksheet #3 Mass Spectrometry

Calculate the average atomic mass for each element. Round off to the correct number of sig figs. Write down the atomic mass from the periodic table rounded off to the same number of sig figs. Show all your work.

Isotope Mass % Abundance Average Mass Atomic Mass(table)

1.

¹⁴N 14.0030744 99.6340

¹⁵N 15.000108 0.366001

2.

²⁰Ne 19.992404 90.92

²¹Ne 20.993849 0.257

²²Ne 21.991385 8.82

3.

⁴⁶Ti 45.952633 7.93

⁴⁷Ti 46.95176 7.28

⁴⁸Ti 47.947948 73.94

⁴⁹Ti 48.947867 5.51

⁵⁰Ti 49.944789 5.34

4.

⁵⁴Fe 53.93962 5.8202

⁵⁶Fe 55.93493 91.660

⁵⁷Fe 56.93539 2.1901

⁵⁸Fe 57.93327 0.33001

5. Silver has two common isotopes. One is 106.90508 amu and 51.35 % and the other is 48.65 %. If the average atomic mass is 107.9730 amu, what is the atomic mass of the other isotope.

6. Copper has two common isotopes. One is 62.92959 amu and 69.09 % and the other is 30.91 %. If the average atomic mass is 63.5472 amu, what is the atomic mass of the other isotope.

7. Complete the chart below.

Protons Electrons Neutrons

²⁸Si

²⁹Si

³⁰Si

8. Write a quantum electron configuration for each of the following.

a) Ne

b) Mg

c) Ti

d) Cr

e) Sr

f) Ag

g) Br

9. What was the first atomic theory to account for the Law of Conservation of Mass?

Explain how the theory accomplished this.

10. What was the first atomic theory to account for electromagnetic radiation (light)?

Explain how the theory accomplished this.

11. What was the first atomic theory to account for the small, dense nucleus?

Explain how the theory accomplished this.

12. What was the first atomic theory to have a wave theory for the electron?

Explain how the theory accomplished this.

13. What was the first atomic theory to account for positive and negative charges in matter? Explain how the theory accomplished this.

Worksheet #4 Periodic Chemistry

1. Define the following:

a) Oxidation

b) Reduction

c) Anion

d) Cation

e) Atom

f) Chemical family

g) Period

2. Why are noble gases stable?

3. Why are non-noble gases un-stable or reactive?

4. Draw Bohr diagrams for the following chemical species.

a) He b) K

c) K⁺ d) S^2-

e) P^3- f) Li⁺

5. Fill in the chart below.

Symbol atom, Protons Neutrons Electrons Valence stable or

cation or Electrons reactive?

anion

Mg²⁺ Cation 12 12 10 8 Stable

Mg

F

F^-

Ne

C

Be

Be²⁺

N^3-

Worksheet #5 Ionic Chemistry

1. Symbol Atom, Protons Neutrons Electrons Valence stable or

Cation or Electrons reactive?

Anion

Ga atom 31 39 31 3 reactive

Ga³⁺

Br

Br^-

Kr

Ca

Ca²⁺

P

P^3-

2. What happens to protons, electrons and neutrons as you move form left to right within a row on the periodic table?

3. Write half-reactions to show how each atom forms an ion. Label each as oxidation or reduction. The first two are done for you.

a) K → K⁺ + 1e^- oxidation

b) N₂ + 6e^- → 2N^3- reduction

c) P

d) O₂

e) Ca

f) Br₂

g) I₂

h) Al

i) Ba

j) Cs

k) Mg

l) Zn

m) Ga

n) Cl₂

o) F₂

4. Describe five properties of:

a) Metals

b) Non-metals

5. Draw Bohr diagrams for each of the following.

a) Na b) Na⁺

c) O d) O^2-

e) Ca f) Ca²⁺

Worksheet # 6 Classifying and Naming Formulas 1

1. Complete the table.

Salt Base Acid Covalent

Nonacid

Litmus

Conductivity

2. Put each formula into the table below.

Ca(OH)₂ NH₄OH CH₃OH C₁₂H₂₂O₁₁

HCl PI₃ K₂SO₄ RbOH

H₃PO₄ NaOH CaCl₂ Li₂SO₄

H₂SO₃ BaF₂ BCl₅ CH₃COOH

H₂CO₃ CsOH S₂Cl₂ Fr₂S

Salt Base Acid Covalent Nonacid

3. Draw Bohr diagrams for each of the following.

a) K⁺ b) P^3-

4. Write half-reactions to show how each atom forms an ion. Label each as oxidation or reduction. The first two are done for you.

a) Ca → Ca²⁺ + 2e^- oxidation

b) O₂ + 4e^- → 2O^2- reduction

c) I₂

d) N₂

e) Cs

f) Ba

g) Al

h) F₂

i) H₂

j) Na⁺ + 1e^- ® Na_(s)

k) N^3- ® N₂ + 6e^-

l) Ca²⁺

m) Ba²⁺
Worksheet # 7 Classifying and Naming Formulas 2

1. Complete the following table by classifying and naming each compound

Formula Classification Name

CuS_(s)

H₃PO_4(s)

P₂O_5(s)

NH₄OH_(s)

Al₂O_3(s)

MgSO_4(s)

HCl_(g)

HCl_(aq)

H₂SO_4(l)

H₂SO_4(aq)

NI_3(s)

N₃Cl_3(s)

CO_(g)

K₂CrO_4(s)

H₂Cr₂O_7(aq)

H₂O_(l)

CrCO_3(s)

HBr_(g)

P₃O_5(s)

.

Complete the following table by classifying and naming each compound.

2. Formula Classification Name

HI _(aq)

(NH₄)₃PO_4(s)

NCl_3(l)

Ba(OH)_2(s)

Rb₂SO_4(s)

CuCl_2(s)

Al₂O_3(aq)

N₃Cl_3(aq)

CO_(g)

H₂SO_3(aq)

CuSO₄ ^. 6H₂O_(aq)

H₃PO_3(s)

Mg₃(PO₄)_2(aq)

HCH₃COO_(aq)

HF_(aq)

N₂O_5(aq)

Na₃PO₄ ^. 5H₂O_(aq)

Ni(NO₃)_3(aq)

SO_(g)

Use your Quantum Periodic Table to write quantum electron configurations for each element below.

3. Sr

4. V

5. 1s²2s²2p⁶3s²

6. 1s²2s²2p⁶3s²3p³

7. 1s²2s²2p⁶3s²3p⁶4s²3d⁴

Pick the best answers. Answers can be used more than once.

Answers: Four Element Theory

Dalton’s Atomic Theory

Thomson’s Atomic Theory

Rutherford’s Atomic Theory

Bohr’s Atomic Theory

Quantum Mechanical Theory

8. The 1^st model of the atom to explain the gold foil experiment

9. The 1^st model to explain light

10. The 1^st model to account for the Law of Constant Composition

11. The 1^st model to have a small, dense nucleus

12. The 1^st model to have an electron as a wave

13. Non-scientific Theory

14. The 1^st model to have electrons

15. The 1^st model to account for the Law of Conservation of Mass

16. Modern theory of the atom

17. The 1^st model to claim the atom is mainly “empty space”

Worksheet #8 Electron Dot Diagrams

Draw structural and electron-dot diagrams for each.

Structural Dot-Diagram

CH₄

CI₄

S₂

P₂

C₂Cl₆

C₂F₄

NF₃

CS₂

N₂Cl₂

HCN

CH₄N₂O

(Symmetrical)

C₆H₆

(Cyclic)

CF₄

N₂Cl₄

NBr₃

Name each compound

1. CH₃COOH_(aq)

2. HBr_(aq)

3. HF_(g)

4. HNO_3(aq)

5. HClO_4(aq)

Write the quantum electron configurations for the following.

6. Cl^-

7. Sr²⁺

8. I

Write a dissociation equation for each to show how each ionizes in water.

9. CH₃COOH_(l)

10. HNO_3(l)

11. Al₂(SO4)_3(s)

12. Co₃(PO₄)_2(s)

Name each compound above.

13.

14.

15.

16.

17. Classify and name the following compounds.

NaOH

BaF₂

BCl₅

CH₃COOH

H₂CO₃

CsOH

S₂Cl₂

BaCl₂

Worksheet # 9 Electron Dot Diagrams

Draw structural and electron-dot diagrams for each.

ClO₃^- PO₄^3-

IO₃^- BrO₃^-

CN^- NO₃^-

SO₄^2- CaCO₃

Li₂SO₄ CCl₄

NI₃ NSCl

NH₄⁺ H₃O⁺

NaCl ClO₃^-

Draw structural and electron-dot diagrams for each.

BrO₄^- PO₃^3-

IO₄^- NO₃^-

HCN SO₃^2-

CO₃^2- CaS

Na₂SO₄ NCl₃

N₂ O₂

Cl₂ C₂H₆

C₂H₄ C₂H₂

Draw electron dot diagrams for each ionic compound

LiCl

Na₂O

K₂S

BaO

GaH₃

Worksheet # 10 Practice Test 1

1. Classify as stable or reactive.

Na N⁺ Ne Cl^- S^2- S^3-

P P^3- Ca Ca⁺² NaCl N^3-

2. Describe a metal and a nonmetal in terms of gaining or losing electrons.

3. Why are noble gases always stable?

4. Determine the number of valence electrons for:

Ca Ca²⁺ Cl Cl^-

O O^2- Al

5. Draw a Bohr diagram for

Ca Ca²⁺ N N^3-

6. Determine the number of protons, neutrons, and electrons in each.

Protons Neutrons Electrons

S

S^2-

Al

Al³⁺

⁴²Ca

7. Classify as ionic or covalent compounds.

HCl CH₃OH H₂O NH₄OH

NaCl MgSO₄ CoCl₂ H₃PO₄

NH₃ P₂O₅ Ba(OH)₂

8. Classify the above compounds into

Acids Covalent Non-acids Salts Bases.

9. Calculate the average atomic mass for magnesium using the following percentage abundance data.

²⁴Mg 78.70% (24.00 amu)

²⁵Mg 10.13% (25.00 amu)

²⁶Mg 11.17% (26.00 amu)

10. Write the formulas for each ionic compound.

Magnesium chloride

Silver phosphate

Cobalt III oxide

Zinc phosphate

Calcium nitride

Copper I nitrate.

11. Name each ionic compound.

Fe₂(SO₄)₃ CoCl₃

Na₂O AgCl

Na₃PO₄ CaF₂

NH₄OH Ca(NO₃)₂

K₂Cr₂O₇ MgCrO₄

12. Name each covalent compound.

P₂O₅ N₂O₃

CO CO₂

SO₂ P₃O₅

C₆H₆ H₂SO_4(l)

HCl_(l) HNO_3(l)

H₂CrO_4(aq) HF_(aq)

H₃PO_4(l) H₂CO_3(aq)

13. Complete the chart below.

Protons Neutrons Electrons Reactive # of valence

or stable? Electrons

Li⁺

S^2-

²²Mg²⁺

14. Write dissociation equations showing how each salt or acid dissolves in water and forms ions. The first one is done for you.

MgCl₂ → Mg²⁺ + 2Cl^-

FeCl₃

Co₂(SO₄)₃

HClO₃

H₂Cr₂O₇

16. Name each acid:

HCl (aq) HClO₃ (aq)

HNO₃ (aq) HBr(aq)

17. Classify each as formula units or molecules.

NaF H₂O

CH₄ NH₄Cl

CaSO₄ CH₃OH

H₂SO₄

18. Indicate the solutions that conduct electricity.

NaCl(s) NaCl(aq)

Ca(OH)₂(aq) HCl(aq)

NH₄OH(aq) C₆H₁₂O₆(aq)

19. Consider the following electron configuration 1s²2s²2p⁶. Determine the element and some ions that have the above electron configuration.

Element

Cation

Anion

20. Describe why NaCl_(s) doesn’t conduct electricity. Describe what happens to NaCl_(s) when it is dissolved in water. Why does it conduct electricity?

21. Write the quantum electron configuration for each (1s²2s²p.....)

He Ar

Na Na⁺

Cl Cl^-

K K⁺

Br Br^-

22. What two particles make up most of the mass within an atom?

23. I am an atom with 35p 40n. Who am I?

24. I am a cation with 56p 81n & 54e. Who am I?

25. Define isotopes.

26. In Rutherford’s Gold foil experiment some particles were completely undeflected and others were radically deflected. Describe the significance of each in terms of the structure of the atom.

27. Define ionic and covalent bonding

28. How many valence electrons are in the calcium ion?

29. How many valence electrons are in the fluoride ion?

30. What is the name of the family that has and electron configuration of:

a) s²p⁵

b) s¹

c) s²p²

31. Consider the following electron configuration 1s²2s²2s²2p⁶3s²3p⁶ . Determine the element and some ions that have the above electron configuration. These are called isoelectronic.

Element:

Cations:

Anions:

Draw electron dot diagrams for the following using brackets for ions. Write a dissociation equation first.

32. NaCl

33. Li₂O

34. CaF₂

Name and classify each compound as an acid, molecular, salt, or base.

35. CuSO_4(aq)

36. P₂O₄(s)

37. H₂SO_4(aq)

38. H₂Cr₂O_7(aq)

39. H₂Cr₂O_7(l)

40. Ca(OH)₂ ^. 5H₂O

41. HBr_(aq)

42. Calculate the average atomic mass for neon if there are three naturally occurring isotopes and they are:

²⁰Ne mass = 19.9924404 amu abundance = 90.92 %

²¹Ne mass = 20.993849 amu abundance = 0.2570 %

²²Ne mass = 21.991385 amu abundance = 8.820 %.

Show some work if you want some marks. Round to an appropriate number of sig figs.

43. Name each compound below. Assume all compounds are aqueous.

Ca(OH)₂ NH₄OH

CH₃OH C₁₂H₂₂O₁₁

HCl PI₃

K₂SO₄ RbOH

H₃PO₄ NaOH

CaCl₂ Li₂SO₄

SiO₂ BaF₂

BCl₅ CH₃COOH

H₂CO₃ CsOH

S₂Cl₂ Fr₂S

Fe₂(SO4)₃ ZnCl₂

Co₃(PO₄)₂ Ag₂Cr₂O₇

Worksheet # 11 Practice Test 2

Balance each equation.

1. ___C₁₆H₃₄ + ___O₂ → ___CO₂ + ___H₂O

2. ___Ga + ____H₂SO₄ → ____H₂ + ____Ga₂(SO₄)₃

Write a balanced equation including phase symbols.

3. Solid carbon reacts with chlorine gas to produce liquid tetracarbon decachloride.

Write chemical formulas for each ionic or molecular compound.

4. Strontium sulphide

5. triphosphorous hexoxide

6. Osmium IV sulphide

Name each chemical formula

7. Sn(CO₃)₂^.5H₂O

8. Si₃F₈

9. NaHCO₃

Classify the following as acids, bases, salts, and molecular non-acids. Name each.

10. Sn(SO₄)₂

11. Ca(OH)₂

12. CH₃COOH

13. S₂O₅

Round off each measured number to three significant figures.

14. 0.0056349

15. 539663

Add or subtract the measured quantities.

16. 153.267 + 0.53493

17 4.5631 x 10²⁴ + 2.36 x 10²³

Simplify the following rounding to the correct number of significant figures.

18. (5.6 x 10 ^-24) (5.37894 x 10^-25)(6.532 x 10¹⁵)

(2.059378 x 10²⁴)(5.23 x 10²²)(9.37894 x 10^-13)

Use unit analysis and the conversion factors to perform the following conversions:

2.210 lb = 1.000 kg 14 lb = 1 stone (defined)

2000 lb = 1 ton (defined) 1.61 km = 1.00 mile

4.54 L = 1.00 gallon 16 oz = 1 lb (defined)

19. 236 oz to stone

20. 8.53 stone to oz

21. 25.6 mi/h to km/s

State the model of the atom is best described by each statement below.

22. First model to account for The Law of Conservation of Mass.

23. The first theory to explain the emission of photons.

24. First model to account for positive and negative charges.

25. First model to account for the wave properties of electrons.

26. First model to include a small dense nucleus.

27. Describes the atom as a small dense nucleus surrounded with electrons, which are not in orbitals.

28. Describes the atom as a small dense nucleus surrounded with electrons, which are in spherical orbitals.

29. Describes the atom as a spherical atom that is indestructible and combines in simple whole number ratios to form compounds.

30. Describes electrons as being contained in a complex 3D orbitals as negative clouds of vibrational energy

31. Non-scientific theory that delayed modern theories of the atom for 1800 years and was shown to be incorrect.

What did the evidence tell us about the nature of the atom?

32. 99 % of alphas in the gold foil experiment were completely un-deflected.

33. 1 % of alphas in the gold foil experiment were radically deflected.

34. Flame spectroscopy of an element produces an emission spectrum consisting of 4 photons.

35. A beam of negative particles is produced in a Crooke’s tube.

36. There are five naturally occurring isotopes of Germanium. Complete the chart to show the number of protons neutrons and electrons.

protons neutrons electrons At. Mass Abundance

⁷⁰Ge 69.92428 20.52%

⁷²Ge 71.92174 27.43%

⁷³Ge 72.9234 7.760%

⁷⁴Ge 73.92115 36.54%

⁷⁶Ge 75.9214 7.760%

37. Calculate the average atomic mass of Germanium. Show some work if you want some marks. Round to an appropriate number of significant figs.

38. Write the quantum electron configurations for the following atoms or ions.

39. F

40. Ga

41. Br^-

42. Rb⁺

Determine the element that has the following electron configuration.

43. 1s²2s²2p⁶3s²

Determine a cation that has the following electron configuration.

44. 1s²2s²2p⁶3s²3p⁶

Determine an anion that has the following electron configuration.

45. 1s²2s²2p⁶3s²3p⁶

Complete the following chart.

Symbol p e n valance el. stable/unstable atom/cation/anion

46. Na

47. P^-3

48. Xe

49. Sr⁺²

Name and classify each compound as an acid, molecular, salt, or base.

50. CuSO_4(aq)

51. P₂O_4(s)

52. H₂SO_4(aq)

53. H₂CO_3(aq)

54. H₂CO_3(l)

55. Ba(OH)₂^. 2H₂O

56. HF_(aq)

Write structural diagrams and electron-dot diagrams for each.

57. CCl₄

58. S₂

59. NH₃

60. CO₂

Write electron-dot diagrams for each.

61. NaCl

62. SO₄^2-

63. Na₃PO₄

64. NO₃^-

65. NH₄⁺

66. IO₄^-