Electrons in Atoms & Periodic Table Chapter 13 & 14 Assignment & Problem Set

Chapter 13 & 14 Assignment & Problem Set

Warm-Ups

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Answer: Answer:

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Chapter 13 & 14 Assignment & Problem Set

Study Guide: Things You Must Know

Vocabulary (know the definition and what it means):

 electronic structure

 Rutherford gold foil experiment

 Bohr planetary model

 quantized energy level

 bright line or emission spectrum

 ground state

 excited state

 photon

 Quantum theory or Wave theory

 atomic orbital

 electron cloud

 electron energy level or shell

 electron configuration

 valence electrons

 metals, nonmetals, metalloids

 representative elements

 transition elements

 inner transition elements

 alkali metals

 alkaline earth metals

 halogens

 noble or inert gases

 periodic trend

 atomic radius

 ionization energy

 ionic radius

 electronegativity

 metallic character of an element

 polar and nonpolar bond

Learning Objectives:

 the historical development of atomic theory, especially the contributions of Rutherford, Bohr and Schrodinger

 Bohr’s planetary model of the atom and how this explains the bright line (emission) spectrum of hydrogen

 the electron energy levels in an atom are quantized

 electrons absorb energy moving to higher energy levels and release energy moving to lower energy levels

 when electrons move to a lower energy level, the energy is released in the form of light (photons). The energy of the light is related to the difference in the starting and ending energy levels of the electron

 in the quantum or wave model of the atom, an atomic orbital is a region in space with high probability of finding an electron, and each atomic orbital can hold two electrons

 how to interpret electron configurations as given on the Regents chart

 how to write the electron configuration of an element in an excited state

 that atoms become ions by losing or gaining electrons to attain a Noble gas electron configuration

 how to explain neon lights, bright line spectra, and flame tests in terms of electron transition between energy levels

 similarities and differences between Mendeleev’s Periodic Table the modern Periodic Table

 the information given for each element on the Regents Periodic Table

 special names for elements in Groups 1, 2, 17, 18

 transition metal salts dissolve in water to give colored solutions

 what is meant by atomic radius, and how to determine atomic radii of elements using Table S

 how to explain the trend in atomic radius across a Period in terms of nuclear charge

 how to explain the trend in atomic radius down a Group in terms of electron energy levels

 the definition of ionization energy, and how to determine ionization energy of elements using Table S

 how to explain trends in ionization energy across a Period and down a Group in terms of the distance of valence electron from the nucleus

 cations are smaller than their atom (less electrons) and anions are larger than their atom (more electrons)

 the definition of electronegativity, and how to determine electronegativity of elements using Table S

 how to explain trends in electronegativity across a Period and down a Group in terms of the distance the valence electron are from the nucleus

 the definition of metallic character

 how to explain trends in metallic character across a Period and down a Group in terms of ionization energy Key Reference Tables

 Periodic Table

 Table S: Properties of Selected Elements

Chapter 13 & 14 Assignment & Problem Set

•Read Chapter 13, except skip “Light and Atomic Spectra” p372-375, “The Quantum Concept and the Photoelectric Effect” pp376-378, and “Quantum Mechanics” pp381-382. Be sure not to skip over “An Explanation of Atomic Spectra” pp379-380.

•Read Chapter 14

•Lab 12: Flame Test

•Lab 13: A Tour of the Periodic Table

•Warm-ups and problems will be collected before you take the test. Answer all problems in the space provided.

For problems involving an equation, carry out the following steps: 1. Write the equation.

2. Substitute numbers and units. 3. Show the final answer with units. There is no credit without showing work.

Atomic Models

1. List a major contribution of each of these scientists to the understanding of the atom:

a. Dalton

b. Thomson

c. Rutherford

d. Bohr

e. Schrodinger

2. Which subatomic particles did Thomson include in the plum-pudding model of the atom?

3. How did Bohr explain why an electron does not fall into the nucleus?

4. What was Bohr’s explanation for the emission spectrum of hydrogen, and why are only a few wavelengths of light observed in this process?

Quantum Mechanical Model of the Atom 5. What is an atomic orbital?

6. The energies of electrons are said to be quantized. Explain what this means.

Chapter 13 & 14 Assignment & Problem Set

7. An atom of an element has two electrons in the first energy level and five electrons in the second energy level.

Write the electron configuration for this atom and name the element.

8. Give the symbol and names of the elements that correspond to these configurations.

a. 2-8-1 b. 2-8-5 c. 2-8-18-9-2 d. 2-8-18-32-12-2

9. Write the electron configuration for each atom.

a. lithium

b. fluorine

c. potassium

10. Write the electron configuration for the elements that are identified only by these atomic numbers.

a. 15 b. 9 c. 18

11. What is the maximum number of electrons that can go into each of the following energy levels?

a. 1^st energy level b. 2^nd energy level c. 3^rd energy level d. 4^th energy level

12. How many electrons are in the second energy level of an atom of each of the following elements?

a. chlorine b. phosphorus c. potassium

13. How many valence electrons do these atoms have?

a. barium b. sodium c. aluminum d. nickel

Atomic Emission Spectra

14. Compare the ground state and the excited state of an electron.

15. Explain the origin of the atomic emission spectrum (bright line spectrum) of an element.

Chapter 13 & 14 Assignment & Problem Set

16. List the colors of the visible spectrum (colors of the rainbow) in order. State which color is the highest wavelength and state which color is the highest energy.

The Periodic Table

17. What is the difference between the way Mendeleev organized the Periodic Table and the way it is done today?

18. Explain how an element's outer electron configuration is related to its position in the periodic table.

19. Why do the elements potassium and sodium have similar chemical and physical properties?

20. Which groups of elements are the representative elements?

21. Which of the following are representative elements? Na, Mg, Fe, Ni, Cl?

22. Categorize each element as a representative element, a transition metal, or a noble gas.

a. 2-8-18-1

b. 2-8-18-6

c. 2-8-18-18-8

23. Which of the following are transition metals: Cu, Sr, Fe, Au, Al, Ge, Co?

24. Write the Regents electron configurations of these atoms:

a. fluorine b. zinc c. aluminum d. barium

25. Write the electron configuration of these elements:

a. the inert gas in period 3 b. the element in Group 14, period 4

Trends in Atomic and Ionic Size

26. What is the trend in atomic radius going left to right across any period? Explain this trend in terms of nuclear charge.

Chapter 13 & 14 Assignment & Problem Set

27. What is the trend in atomic radius going down a group? Explain this trend in terms of electron energy levels.

28. Indicate which element in each pair has the greater atomic radius.

a. sodium, lithium

b. strontium, magnesium

c. carbon, germanium

d. selenium, oxygen

29. Arrange these elements in order of increasing atomic size: sulfur, chlorine, aluminum, and sodium.

30. How does the ionic radius of a typical metallic atom compare with its atomic radius? Explain.

31. How does the ionic radius of a typical anion compare with the radius for the corresponding neutral atom?

Explain.

32. Circle which particle has the larger radius in each in each atom/ion pair?

a. Na, Na⁺ b. S, S^-2 c. I, I^- d. Al, Al⁺³

33. The Mg⁺² and Na⁺ ions each have ten electrons surrounding the nucleus. Which ion would you expect to have the smaller radius? Why?

Trends in Ionization Energy

35. What is the trend in ionization energy going left to right across any period? Explain in terms of atomic size.

36. What is the trend in ionization energy going down a group? Explain in terms of atomic size.

Chapter 13 & 14 Assignment & Problem Set

37. Circle the element in each pair that has the greater first ionization energy?

a. sodium, potassium

b. magnesium, phosphorus

c. lithium, boron

d. magnesium, strontium

e. cesium, aluminum

38. Would you expect metals or nonmetals to have higher ionization energies? Why?

Trends in Electronegativity

40. Which element has the greatest electronegativity? What is its value?

41. What is the trend in electronegativity going down any group? Explain.

42. What is the trend in electronegativity going left to right across any period? Explain.

43. Is a C-O bond polar? Explain in terms of electronegativity.

44. Which bond is more polar, C-O or C-H? Explain in terms of electronegativity difference.

Metallic Character

45. What is meant by metallic character?

46. Circle the more metallic element in each pair.

a. Na or Cs b. Na or Al c. Si or S d. Sr or F

Chapter 13 & 14 Assignment & Problem Set

Review

47. Give the number of protons and electrons in each of the following:

a. Cs b. Ag⁺ c. Se^-2

48. Name or give the formula for the following, and state whether each is molecular (M) or ionic (I).

a. N2O4

b. Cs2O

c. tin(II) sulfate

d. ammonium carbonate

e. Pb3(PO4)2

49. Convert the following quantities.

a. 3.7 mg to kg

b. 132 km to mm

c. 62 m/s to km/hr

50. Chlorine has two isotopes: Cl-35 which is 75.8 % abundant and Cl-37 which is 24.2% abundant. Calculate the atomic mass of chlorine?