Wave Nature of Light

1. Draw a wave below and label the following parts: peak, trough, wavelength and amplitude

2. Draw two waves with different frequencies and circle the wave that has a higher frequency.

3. What is electromagnetic radiation? At what speed does electromagnetic radiation travel?

4. How are different types of electromagnetic radiation similar? How do they differ?

5. What does the wavelength of electromagnetic radiation represent? How is the wavelength of radiation related to the frequency of the radiation?

6. What do we mean by frequency of electromagnetic radiation? Is the frequency the same as the speed of electromagnetic radiation?

7. What are the basic SI units for (a) the wavelength of light, (b) the frequency of light, (c) the speed of light?

8. Arrange the following kinds of electromagnetic radiation in order of increasing wavelength: infrared, green light, red light, radiowaves, X-rays, ultraviolet light.

9. Why is the study of electromagnetic radiation important in Chemistry?

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© John Erickson, 2005 WS5-1LightWaves

The behavior of light indicates that it is comprised of waves. The distance between successive waves is called the wavelength (�) and the wavelength determines the type of light. The size of the waves determines the type of light. All of the various light waves move with the same speed, a value abbreviated (c) equal to 3.00 × 108 m/s. The frequency (�) that light waves pass a given point is measured in waves/second or simply ‘per second’ (1/s). The unit 1/s is also given the name hertz (Hz).

Answer the following questions about light waves. Show all work

1. What type of light has a wavelength of: a) 5.0 × 10-4 m? b) 2.4 × 10-8 m? c)12 mm?

2. An ultraviolet light wave is used to kill bacterial. It has a frequency of 1.2 × 1015 1/s. Find the wavelength.

3. An x-ray has a wavelength of 1.54 × 10-10 m. Find the frequency of this light.

4. A visible light wave has a frequency of 7.5 × 1014 1/s. Find the wavelength in nanometers (nm) and determine the color of the light.

5. One of the light waves produced when hydrogen is energized has a wavelength of 410.5 nm. What is the frequency of this light?

6. The frequency of light used to heat food in a microwave oven is 2.45 GHz (2.45 × 109 1/s). What is the wavelength of this light?

7. A radio wave broadcast on the AM dial has a wavelength of 280.4 m. Find the frequency of this radio wave in hertz. Convert the frequency to kilohertz.

8. What is the wavelength of a radio wave broadcast with a frequency of 99.5 MHz (FM 99.5)?

9. Pilots often use waves of about 2.340 m to communicate. What is the frequency of this wave?

10. The light used in night vision devices has a wavelength of about 25 micrometers (µm). What is the frequency of this light? In what part of the electromagnetic spectrum are these waves?

Name ________________ Light Waves Chem Worksheet 5-1

1 m = 1 × 103 mm

1 m = 1 × 106 µm

useful equations

c = � × � c = 3.00 × 108 m/s

1 m = 1 × 109 nm

1 m = 1 × 1010 Å c = 3.00 × 108 m/s

�

�

Violet Light Shorter wavelength (�) and higher frequency (�)

Red Light Longer wavelength (�) and lower frequency (�)

V I B G Y O R

400 500 600 700 750 nm

wavelength (m)

1020

10-9 10-7 10-5 10-3 10-1 101 103

1018 1016 1014 1012 1010 108 106 104

10-11

gamma rays x-rays

ultra- violet infrared microwaves radio/TV waves

visi

ble

frequency (1/s)

1 MHz = 1 × 106 Hz 1 GHz = 1 × 109

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Particle Nature of Light

1. What did Max Planck propose about the nature of energy?

2. If human height were quantized in one-foot increments what would happen to the height of a child as she grows up?

3. How is the frequency of electromagnetic radiation related to its energy according to Planck’s equation?

4. What is a photon?

5. How is the energy carried per photon of light related to the wavelength of light according to Planck’s equation? Does short-wavelength light carry more energy or less energy than long-wavelength light?

6. Explain the photoelectric effect. How does it demonstrate that light has particulate properties?

7. Does energy have mass? Explain your answer.

8. Does matter exhibit wave like properties? Explain your answer.

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© John Erickson, 2005 WS5-2PlancksEq

E = h × � h = 6.63 × 10-34 J·s

useful equations

c = � × � c = 3.00 × 108 m/s

1 m = 1 × 109 nm 1 kJ = 1000 J

example

Light with a wavelength of 525 nm is green. Calculate the energy in joules for a green light photon.

- find the frequency: υλ ×=c λc

v =

nmm

nm

smv

9

8

1011

525

/1000.3

××

×= sv /11071.5 14×=

- find the energy: υ×= hE )/11071.5)(10626.6( 1434 ssJE ×⋅×= − photonJE /1078.3 19−×=

Problem-Solving Strategy Known Unknown Frequency (�) Energy (E) Wavelength (�) Frequency (�) Energy (E) Energy (E) Frequency (�) Wavelength (�)

Max Planck theorized that energy was transferred in chunks known as quanta, equal to h�. The variable h is a constant equal to 6.63 × 10-34 J·s and the variable � represents the frequency in 1/s. This equation allows us to calculate the energy of photons, given their frequency. If the wavelength is given, the energy can be determined by first using the wave equation (c = � × �) to find the frequency, then using Planck’s equation to calculate energy.

Use the equations above to answer the following questions.

1. Ultraviolet radiation has a frequency of 6.8 × 1015 1/s. Calculate the energy, in joules, of the photon.

2. Find the energy, in joules per photon, of microwave radiation with a frequency of 7.91 × 1010 1/s.

3. A sodium vapor lamp emits light photons with a wavelength of 5.89 × 10-7 m. What is the energy of these photons?

4. One of the electron transitions in a hydrogen atom produces infrared light with a wavelength of 7.464 × 10-6 m. What amount of energy causes this transition?

5. Find the energy in kJ for an x-ray photon with a frequency of 2.4 × 1018 1/s.

6. A ruby laser produces red light that has a wavelength of 500 nm. Calculate its energy in joules.

7. What is the frequency of UV light that has an energy of 2.39 × 10-18 J?

8. What is the wavelength and frequency of photons with an energy of 1.4 × 10-21 J?

Name ________________ Planck’s Equation Chem Worksheet 5-2

λν c= νhE =

hE=ν

νhE =

λν c=

4

Line Spectrum of Hydrogen and the Bohr Model

1. What does the ground state of an atom represent?

2. What happens to the electron in a hydrogen atom when it absorbs energy and enters an excited state?

3. When an atom in an excited state returns to its ground state, what happens to the excess energy of the atom?

4. What is the continuous spectrum? How is it produced?

5. How is the emission spectrum of hydrogen different than the continuous spectrum of white light? Why is this significant?

6. How does the Bohr model of a hydrogen atom fit with the idea that energy is quantized?

7. Why is the Bohr model of the atom considered to be fundamentally incorrect?

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Quantum Mechanical Model and Atomic Orbitals

1. What information does a wave function give us about an electron outside the nucleus of an atom? Can we know how an electron moves given a wave function?

2. What does Heisenberg’s Uncertainty Principle state?

3. Describe the difference between an orbit (as described by the Bohr Model) and an orbital (as described by the wave mechanical model).

4. Describe the significance of quantum numbers.

5. What does the principal quantum number (n) tell us about an orbital?

6. What are sublevels?

7. Draw a picture of an s-orbital and a p-orbital.

8. Which orbital is the first to be filled in any atom? Why?

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Chemistry WS5-5OrbitalDiagrams

1s 2s 2p Se:

3s 3p 4s 3d Each orbital is half-filled before being completely filled.

Two electrons in the same orbital must have opposite spin.

Electrons fill the lowest available energy levels first.

This violates Hund’s rule. One electron should be distributed to each of the 3p orbitals before doubly filling any.

1s 2s 2p 3s 3p

Boxes drawn for various sublevels

s sublevel: 1 orbital

p sublevel: 3 orbitals

d sublevel: 5 orbitals

f sublevel: 7 orbitals

1s 2s 2p 3s 3p 4s

1s 2s 2p 3s 3p

1s 2s 2p

3d

5s 5p 4d

1s 2s 2p 3s 3p

4s 1s 2s 2p 3s 3p

4s 4p 3d

An orbital diagram uses boxes with arrows to represent the electrons in an atom. Each box in an orbital diagram represents an orbital. Orbitals have a capacity of two electrons. Arrows are drawn inside the boxes to represent electrons. Two electrons in the same orbital must have opposite spin so the arrows are drawn pointing in opposite directions. The following is an orbital diagram for selenium. In writing an orbital diagram the first step is to determine the number of electrons. Normally this is the same as the number of protons, which is known as the atomic number. Next the boxes are drawn for the orbitals. Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau rule. The Pauli exclusion principle requires that electrons in the same orbital have opposite spin. Hund’s rule states that orbitals in a given sublevel are half-filled before they are completely filled.

Write the name and symbol for the elements with the following orbital diagrams.

1. 4.

2. 5.

3. 6.

There is an error with each of the following orbital diagrams. Explain the error.

7. 8.

Write orbital diagrams for the following. You may abbreviate using a noble gas.

9. hydrogen

10. boron

11. sodium

12. krypton

13. chromium

14. phosphorus

15. carbon

16. cobalt

17. platinum

18. plutonium

19. oxygen

20. potassium

Name ________________ Orbital Diagrams Chem Worksheet 5-5

[Kr]

7s [Rn]

5f

[Ar]

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© John Erickson, 2005 WS5-6ElectronConfig

An electron configuration is simply a list of the orbitals that contain electrons for a given element.

The orbital designation is followed by a superscript number that tells how many electrons are found in that orbital. The following designation represents an atom with electrons found in the 1s, the 2s, the 2p, and the 3s orbitals. There are a total of 11 electrons in the atom. This represents the element sodium.

The orbitals of an atom fill in a specific sequence. The pattern fits very nicely with various regions of the periodic table. The table is been sectioned into blocks which are labeled: s block, p block, d block, and f block. The rows of each block are labeled as well. Using this shortcut, electron configurations can be determined easily. The element manganese is the fifth element in the 3d row. The orbitals before the 3d orbital are all filled so it has full 1s, 2s, 2p, 3s, 3p, and 4s orbitals. Since manganese is the fifth element in the 3d row we designate the 3d orbital with 5 electrons.

Electron configurations can be abbreviated by writing the element symbol for the previous noble gas in brackets, followed by the remaining electrons. For example, rather than writing all of the electrons in antimony (element 51), the first 36 electrons are represented by [Kr]. The remaining electrons are notated using orbital names and superscript numbers.

Write the name and symbol for the atoms with the following electron configurations.

1. 1s22s22p63s23p64s23d104p4 4. 1s22s22p63s23p1

2. 1s22s22p63s23p64s23d104p65s24d105p66s1 5. [Rn]7s25f9

3. 1s22s22p63s23p64s23d104p65s24d7 6. [Xe] 6s24f145d106p2

Write complete electron configurations for the following substances.

7. nitrogen 10. nickel

8. magnesium 11. tin

9. niobium 12. chlorine Write abbreviated electron configurations for the following elements.

13. arsenic 19. sulfur

14. thulium 20. zirconium

15. rubidium 21. argon

16. einsteinium 22. iron

17. platinum 23. polonium

18. molybdenum 24. bohrium

Name ________________ Electron Configuration Chem Worksheet 5-6

Orbital name Number of electrons

Ex. 1s22s22p63s1

Complete configuration Sb 1s22s22p63s23p64s23d104p65s24d105p3

Kr 1s22s22p63s23p64s23d104p6

Sb [Kr] 5s24d105p3 Abbreviated configuration

s block p block

d block

f block

1s 2s

3s

5s

6s

7s

2p

3p

4p 5p

6p

3d 4d

5d

6d

4f 5f

4s

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Electron ConfigurationsMatching:

a) orbital b) principal quantum number c) valence shell

d) kernel e) orbital diagram

1. One of these is assigned to each energy level in the Bohr model of the atom.

2. Used to describe the placement of electrons in energy levels and sublevels.

3. Part of the atom exclusive of the outer shell of electrons

4. Electrons with the same energy and opposite spin

5. The outermost principal energy level

a) d b) p c) f d) n e) s

6. Symbol for the principal quantum number of an energy level

7. Symbol for the fourth energy sublevel

8. Symbol for the third energy sublevel

9. Symbol for the first energy sublevel

10. Symbol for the second energy sublevel

a) electron configuration b) Pauli exclusion principle c) orbital

d) uncertainty principal e) quantum numbers

11. Four of these are used to describe the location of an electron

12. Specifies that when two electrons occupy the same orbital, they must have opposite spins

13. Arrangement of units of negative charge among the various orbitals of the atom

14. Specifies that it is impossible to know location and velocity of a subatomic particle at the same time

15. Region in space where an electron with specified energy may be found.

Multiple Choice:

16. The characteristic bright-line spectrum of an element is produced when electrons

a) fall back to lower energy levelsb) are gained by a neutral atomc) are emitted by the nucleus as beta particlesd) move to higher energy levels

CHEMISTRY

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17. When an atom goes from an excited state to the ground state, the total energy of the atom

a) increases b) decreases c) remains the same

18. The lowest sublevel in each principal energy level is represented by the symbol

a) f b) p c) s d) d e) n

19. What is the maximum number of electrons in the nth principal energy level

a) n b) 2n c) n2 d) 2n2

20. What is the number of kinds of sublevels in the energy level that has the principal quantum number of 2?

a) 2 b) 3 c) 4 d) 8

21. What is the number of electrons permitted in the d sublevel of the third energy level?

a) 3 b) 9 c) 10 d) 14 e) 18

22. What is the number of orbitals in the 4f sublevel?

a) 1 b) 4 c) 7 e) 16 e) 14

23. The arrangement that represents the lowest energy of electrons in an atom is called the

a) ground state b) excited state c) valence electrons d) kernel electrons

24. When an orbital is occupied by two electrons, the electrons must have

a) the same charge and the same spin b) opposite charge and the same spinc) the same charge and opposite spin d) opposite charge and opposite spin

25. The part of the atom that contains electrons that are not the valence electrons

a) nucleus b) standing wave c) orbital d) kernel

26. In the charge-cloud model, energy sublevels are divided into

a) energy levels b) orbits c) orbitals d) configurations

27. An orbital may contain more than one

a) electron b) proton c) energy level d) orbit

28. According to the Heisenberg uncertainty principle, which two characteristics of a small particle cannot beknown precisely at the same time?

a) mass and velocity b) location and motionc) wavelength and diameter d) radius and distance from its nearest neighbor

29. Which sublevel contains electrons with the highest energy?

a) 3p b) 2p c) 3s d) 4s

CHEMISTRY

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30. Which electron transition is accompanied by the emission of energy?

a) 1s to 2s b) 2s to 2p c) 3p to 3s d) 3p to 4p

31. What is the number of completely filled orbitals in an atom of fluorine in the ground state?

a) 4 b) 5 c) 6 d) 9

32. What is the total number of occupied principal energy levels in an atom of Aluminum in the ground state?

a) 2 b) 3 c) 4 d) 5

33. Which electron configuration represents a neutral atom of carbon in an excited state?

a) 1s2 2s2 2p2 b) 1s2 2s2 2p3 c) 1s2 2s1 2p3 d) 1s2 2s2 2p6

34. Compared to an atom of chlorine in an excited state, a chloride ion has

a) one more electron b) one fewer electron c) the same number of electrons

35. Which particle contains a total of 2 protons, 2 neutrons, and 2 electrons

a) a cation b) an anion c) the kernel of an atom d) a neutral atom

36. The electron configuration 1s2 2s2 2p5 could represent the

a) ground state of a fluorine atom b) ground state of a magnesium ionc) excited state of a magnesium atom d) excited state of a fluoride ion

37. What is the number of valence electrons in an atom that has the electron configuration1s2 2s2 2p6 3s2 3p2

a) 2 b) 3 c) 4 d) 6

38. Which is the electron configuration of a neutral atom in the ground state with a total of six valenceelectrons?

a) 1s2 2s2 2p2 b) 1s2 2s2 2p4 c) 1s2 2s2 2p6 d) 1s2 2s2 2p6 3s2 3p6

39. Which applies to an atom that has taken on one additional electron?

a) It has entered the excited state b) It has become an ionc) It has absorbed energy d) It has acquired an additional valence shell

40. What additional information must be known in order to determine whether or not an atom with 8 protonsand 8 electrons is in the ground state?

a) the number of neutrons b) the electron configurationc) the atomic mass d) the mass number

CHEMISTRY

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12

The Structure of the Periodic Table

Directions:1. Label the columns with the both the 1 – 18 and IA – VIIIA designations.2. Write the period numbers in the left margin.3. Draw a heavy line to separate the metals from the nonmetals.4. Color the groups listed in the key with different colored pencils, indicating colors in the key.5. Write the atomic numbers in the upper left hand corner of elements 57, 58, 71, 72, 89, 90, 103, and 104. Write the

symbols of lanthanum and actinium in the correct boxes. Draw an arrow to indicate where the lanthanide and actinideseries fit into the main periodic table.

6. Write the symbols of the elements {Ag, Al, C, Ca, Cl, Cu, F, Fe, H, He, Hg, K, Mg, N, Na, O, P, Pb, S, Zn} in theproper locations.

K E Y

alkali metals

alkaline earth metals

halogen family

noble gases

transition metals

lanthanide series

actinide series

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Periodic Trends

1. Discuss the importance of Mendeleev’s periodic law.

2. Identify each element as a metal, metalloid, or nonmetal.a) fluorineb) germaniumc) zincd) phosphorouse) lithium

3. Give two examples of elements for each category.a) noble gasesb) halogensc) alkali metalsd) alkaline earth metals

4. What trend in atomic radius do you see as you go down a group/family on the periodic table?What causes this trend?

5. What trend in atomic radius do you see as you go across a period/row on the periodic table?What causes this trend?

6. Circle the atom in each pair that has the largest atomic radius.

a) Al B b) S O c) Br Cl d) Na Al e) O F f) Mg Ca

7. Define ionization energy.

8. Is it easier to form a positive ion with an element that has a high ionization energy or anelement that has a low ionization energy? Explain.

9. Use the concept of ionization energy to explain why sodium form a 1+ ion (Na+) butmagnesium forms a 2+ ion (Mg2+).

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11. What trend in ionization energy do you see as you go across a period/row on the periodictable? What causes this trend?

12. Circle the atom in each pair that has the greater ionization energy.

a) Li Be b) Na K c) Cl Si

d) Ca Ba e) P Ar f) Li K

13. Define electronegativity

14. What trend in electronegativity do you see as you go down a group/family on the periodictable? What causes this trend?

15. What trend in electronegativity do you see as you go across a period/row on the periodictable? What causes this trend?

16. Circle the atom in each pair that has the greater electronegativity.

a) Ca Ga b) Li O c) Cl S

d) Br As e) Ba Sr f) O S

10. What trend in ionization energy do you see as you go down a group/family on the periodictable? What causes this trend?

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