Atomic Structure

1

UNIT 1: ATOMIC STRUCTURE OUTCOMES

All important vocabulary is in Italics and bold.

Explain the difference between isotope and radioisotope.

Determine average atomic mass using isotopes and their relative abundance.

Describe the electromagnetic spectrum in terms of frequency, wavelength, and energy.

Use Plank’s equation and the speed of light to solve equations for electromagnetic waves.

Explain the basics of wave-particle duality.

Define the relationship between quantum, photon and electron.

Describe how a produced line spectra relates to the Bohr diagram for a specific element.

Explain the problems with the Bohr model.

Explain the quantum arrangement of electrons in an atom.

Include: energy level, sublevel, and orbital.

Write the electron configuration or orbital box diagrams for a variety of atoms and ions

Relate the electron configuration of an element to its valence electron(s) and its position on the periodic

table.

Draw basic Lewis dot structures of atoms and compounds.

Using VSEPR, predict bond shape from electron arrangement in Lewis dot diagram.

Describe factors that affect electron position around a nucleus.

Include: nuclear charge, distance, shielding.

Explain periodic trends using above factors.Include: atomic / ionic radii, ionization energy, and electronegativity.

Additional KEY Terms

Half-life Radioactive decay

Spectroscopy Emission Quantum

Line spectrum Absorption Spectra

Threshold energy Quantum number (n)

Shorthand notation

2

COMMON EQUIPMENT USED IN A CHEMISTRY LABORATORY

Beaker Erlenmeyer Flask Round-bottom (Florence) flask

Test tube Graduated cylinder Eye-dropper

Ring support Tongs Funnel

Double Burette clamp Bunsen burner scoopula

3

Mr. Storie 40S Chemistry Atomic Structure - Worksheets

Ring stand Burette Volumetric flask

Test tube rack Buchner funnel Mortar and pestle

Wire gauze Filter flask Watch glass

Utility clamp Test tube holder Hot plate

4

WORKPLACE HAZARDOUS MATERIALS INFORMATION SYSTEM

(WHMIS)

The WHMIS symbols were developed so all dangerous materials used in all

workplaces and schools, would be labeled in the same way. Pay careful attention to

any warning symbols on the products or materials that you handle in a laboratory.

5

METRIC CONVERSION CHART:

Prefix Symbol Factor of base unit Examples

giga G 1 000 000 000 = 109

1 billion

mega M 1 000 000 = 106

1 million

kilo k 1 000 = 103

km, kL

1 thousand

hecto h 100 = 102

1 hundred

deca da 10 = 101

Ten

1 = 100

deci d 0.1 = 10-1

1/10th

centi c 0.01 = 10-2

cm

1/100th

milli m 0.001 = 10-3

mg

1/1000th

, mL

micro µ 0.000 001 = 10-6

µg, µL

1/100 000th

nano n 0.000 000 001 = 10-9

1/1 000 000th

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Answer the following significant figure questions:

1. How many significant figures in each of the following?

a) 657 f) 4.26 k) 28.07

b) 509 g) 0.115 l) 2600

c) 0.0082 h) 0.0001 m) 0.00800

d) 706 000 i) 0.05009 n) 14.50

e) 10 000 j) 5.000 o) 540 000

2. Round the following to 3 significant figures.

a) 32460 f) 10250

b) 0.0052430 g) 0.032457

c) 0.0007045 h) 3.097×103

d) 6.045 i) 8.09200×10−4

e) 90310 j) 609 500

3. Perform the following calculations, rounding the answer to the correct number of significant figures.

a) (3.10)(48.9)(1.2)(901.5)

b) 979.6 ÷ 7000

c) 1.8 × 2.97 × 162.3

d) (45.5 – 31.06)(6.2) =

e) (1.3637 × 106)(5.61 × 10

−2)

f) 3.2 x 10 2 + 6.32

g) 0.21 + 4.33 + 0.008

h) 1489.75 + 6.3 - 158

i) (9.0 x 10 -5

)(0.3)

j) (0.0056)(2.6 x 10 3)

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4. Perform the following operations and give the answer to the correct number of significant digits.

a) 15.1 + 75.32

b) 178.90456 – 125.8055

c) 4.55 × 10-5

– 3.1 x 10-5

d) 0.000159 + 4.0074

e) 1.805 × 104 + 5.89 × 10

4

f) 0.000817 - 0.0000481

g) 8.166 × 105 – 7.819 × 10

5

h) 45.128 + 8.50187 – 42.18

i) 5.677 × 10-6

+ 7.785 × 10-6

j) 8.75 × 10-9

+ 6.1157 × 10-9

k) 1.99 3.1 l) 1200.0 3.0

m) 5.32 x 10-4

4.218 × 10-8

n) 45.32 x 2.3 o) 0.02400 6.000 p) 12.4 x 0.30

q) (5.50 x 108) × (4 x 10

5)

r) 7.4 3 s) 4.75 × 5

t) 2.5 × 6.700 0.891

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Write the proper chemical names OR write the proper chemical formula for the following compounds:

1. Name each of the following.

a. NaC2H3O2

b. Ni(NO3)2

c. HgF2

d. Sn3(PO4)2

e. Ca(OH)2

f. O2

g. NO

h. CO2

i. NO2

j. Ca(ClO)2

k. ZnS

l. FeCO3

m. Li2SO4

n. Ba(NO2)2

o. CCl4

p. NH3

q. P3N5

r. N2O

2. Write the chemical formula of the following compounds.

a. calcium sulfite

b. ammonium dichromate

c. potassium permangante

d. lead (II) iodide

e. magnesium acetate

f. carbon tetrachloride

g. dichlorine monoxide

h. sulfur trioxide

i. ammonia

j. iron (II) hydroxide

k. potassium thiocyanate

l. manganese (VII) oxide

m. rubidium perchlorate

n. cadmium nitrate

o. chlorine gas

p. nitrogen dioxide

q. diphosphorous decaoxide

r. bromine pentafluoride

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3. Write balanced equations from the word equations below (you may need to figure the products):

a. Zinc and lead (II) nitrate react.

b. Aluminum bromide and chlorine gas react to form aluminum chloride and bromine gas.

c. Sodium phosphate and calcium chloride react.

d. Potassium metal and chlorine gas combine.

e. Copper and sulfuric acid react to form copper (II) sulfate, water and sulfur dioxide.

4. Answer the following stoichiometry questions:

a. How many litres are in 285 grams of tricarbon octahydride gas at STP?

b. Find the mass of 9.36 x 1023

formula units of calcium hydroxide.

c. Find the number of atoms of oxygen in 1.5 x 10-5

moles of Mg3(PO4)2.

5. Explain the difference between ionic, non-polar covalent and polar covalent bonding.

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ANSWER THE FOLLOWING ATOMIC STRUCTURE QUESTIONS:

1. Complete the following table:

Symbol Z A p+

e-

no

Atomic charge

F 0

Hg 0

3 0

55 0

83.8 0

47.8 0

31 0

25 0

Br -

Pb +4

Na +1

57 139 55

51 23 22

P -3

18 +2

2. Identify the numbers of protons, neutrons and electrons in a neutral atom of each of the following:

a. U

b. U

c. S

3. Complete the following table:

Isotope Symbol Z p+

no

A

Hydrogen-2 H

Zn

92 142

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4. Calculate the average atomic mass of each element.

Element Symbol Mass Mass (µ) Relative Number (exactly) Abundance (%)

Carbon C-12 12 12 98.98

C-13 13 13.003 1.11

Silicon Si-28 28 27.977 92.21

Si-29 29 28.976 4.70

Si-30 30 29.974 3.09

5. Using the graph below, calculate the average atomic mass of copper.

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ANSWER THE FOLLOWING ATOMIC SPECTROSCOPY:

1.

2.

3.

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ANSWER THE FOLLOWING EMR and WAVE QUESTIONS:

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.

3. Why is light called electromagnetic radiation?

4. Compare and contrast line and continuous spectra.

5. A certain substance strongly absorbs infrared light having a wavelength of 6.50 x 10-6

m (µm

– micrometres). What is the frequency of the light?

6. Ozone protects the Earth’s inhabitants from the harmful effects of ultraviolet light arriving from the

sun. This shielding is maximum for light of frequency 1.02 x 1015

Hz. Calculate the wavelength.

7. A typical radar signal has a wavelength of 3.19 cm. What is the frequency in Hertz.

8. EM waves used in FM broadcasting have a frequency of about 100 megahertz (MHz - 106).

AM broadcasting waves are 1 megahertz. Calculate the wavelength of these waves.

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9. Hydrogen atoms are ionized by radiation of 9.12 x 10-8

m. One photon can remove one electron.

Calculate the ionizing energy in joules per electron.

10. A violet spectral line consists of light of wavelength 4.1 x 10-7

m. Calculate the frequency of the light

emitted. Calculate the energy per mole of photons.

11. To remove an electron from an atom requires an input of 1.04 x 10-17

J.

a. What is the frequency of the required radiation? b. What is the wavelength of this ionizing radiation?

12. Sound is a wave phenomenon. If the velocity of sound in air is 330 m/s, calculate the wavelength of

sound corresponding to the middle “C” of a piano which has a frequency of 262 Hz.

13. Visible light includes all radiation of wavelengths of about 4000 Å (violet) to 7000Å (red). Calculate

the range of visible frequencies. (1 Å = 1 x 10 -10

m)

14. What is the energy in joules of one photon of yellow light having a wavelength of 5.8 x 10-5

cm?

15. An atomic emission is equal to 1260 kJ/mol. What is the frequency of the emitted radiation?

Answer Key

5. 4.62 x 1013

Hz, 6. 2.94 x 10- 7

m, 7. 9.40 x 109 Hz, 8. 3.00 m, 300 m, 9. 2.18 x 10

- 18 J/e

-, 10. 7.3 x 10

14 Hz,

2.9 x 105 J/mol, 11. 1.58 x 10

16 Hz, 1.90x 10

-8 m, 12. 1.26 m, 13. 7.5 x 10

14 Hz – 4.3 x 10

14 Hz, 14.

3.4 x 10-19

J/photon, 15. 3.20 x 1015

Hz

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LINE SPECTRA INVESTIGATION

PURPOSE:

To observe the line spectra of some elements and to gain an understanding of continuous versus line spectra.

INTRODUCTION:

In this investigation you will observe both continuous and line spectra. In order to do this you will need to

observe light from different sources through a diffraction grating. A diffraction grating (spectroscope) is very

similar to a prism in that it separates light into its component colours to produce a spectrum.

APPARATUS:

Gas discharge tube Spectroscope

Incandescent light source Power source

PROCEDURE:

1. Observe an incandescent light through you spectroscope. Draw and label the spectrum you observe.

2. Using your spectroscope, observe the spectra for several gas discharge tubes. Sketch each spectrum, and be

sure to indicate the colour of each observed line.

ANALYSIS:

1. Classify each spectrum you drew as a continuous spectrum or a line spectrum.

2. What is the difference observed between a continuous and a line spectrum?

CONCLUDING QUESTIONS:

1. Why did each gas show a colour only when electricity was passed through the discharge tube?

2. What do the different colours in a line spectrum represent?

3. Suggest an explanation for the differences in the brightness of the lines in the line spectra.

4. Why do different substances show different spectra?

5. How does the Bohr model of the hydrogen atom explain the spectrum of hydrogen?

6. What is one practical application of passing an electrical charge through a gas?

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ANSWER THE FOLLOWING QUESTIONS ON ENERGY LEVELS AND EMR:

1. Arrange the following types of radiation from lowest to highest energy: x-rays, infrared light, blue light,

red light, radio waves, gamma rays.

2. What trends exist moving from radio waves to gamma rays on the EMR spectrum?

3. Consider red light having a wavelength of 6.6 x 10-5

cm.

a. What is the frequency of this light? b. What is the energy of one photon of this light?

4. How does the spectrum of a hydrogen discharge tube differ from that of an incandescent light source?

5. Some children are playing on the “staircase” shown below, jumping from one stair to another. How many

different possible jumps (up-only) can the children make?

6. To remove an electron from lithium requires energy – called “ionization energy.” The first ionization

energy of lithium is 518 kJ/mol. a. Determine the frequency of one photon of this energy.

b. Determine the wavelength of this ionizing energy.

7. The total energy required to boost electrons in a mole of H atoms from level n=2 to n=4 is +246 kJ/mol. a. Calculate the frequency and wavelength of this radiation per photon.

b. Is this radiation visible?

8. A hydrogen atom emits radiation of frequency 2.92 x 1015

Hz.

a. What is the wavelength of this radiation? b. What is the energy of this radiation in J/pho and in kJ/mol?

ANSWERS

3. 4.5 x 1014

Hz, 3.0 x 10-19

J/pho, 5. 15, 6. 1.30 x 1015

Hz, 2.31 x 10-7

m, 7. 6.18 x 1014

Hz, 4.85 x10-7

m

8. 1.02 x 10-7

m, 1.93 x 10-18

J/pho, 1160 kJ/mol

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ANSWER THE FOLLOWING QUESTIONS ON ELECTRON CONFIGURATIONS:

1. State the names of the three main rules/principles use to construct an energy-level diagram. Describe each

in your own words. 2. How can the periodic table be used to help complete energy-level diagrams?

3. 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 4. 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+

5. Complete orbital box diagrams for a potassium ion and a chloride ion. Which noble gas atom has the same

diagram (*ISOELECTRIC) as these ions?

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6. It takes 7.21 x 10-19

J of energy to remove an electron from an atom. Can light with a wavelength of 300 nm

do this?

7. If it takes 492 kJ to remove 1 mole of electrons from gold atoms, how much energy is needed to remove a

single electron from gold? What is the minimum wavelength of light capable of doing this?

8. 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

i. Te 9. 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+

10. Write shorthand electron configurations for the following:

a. Sr b. Sn

c. Fe2+

d. Cl-

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ANSWER THE FOLLOWING QUESTIONS ON ELECTRON CONFIGURATIONS:

1. Briefly explain why barium has 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 1s22s

22p

63s

23p

64s

2

element B 1s22s

22p

63s

23p

5

element C 1s22s

22p

63s

23p

64s

1

element D 1s22s

22p

63s

23p

6

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?

5. Which of these elements would have the highest value for the second ionization energy? a. K b. Si c. Ar d. Br

6. Which of the following has the largest atomic radius and which has the smallest? a. nitrogen b. antimony c. arsenic

7. For each of the following properties indicate which has the larger value fluorine or bromine:

i) atomic radius

ii) first ionization energy

iii) ionic radius

8. Arrange the following from largest to smallest. Explain the order. Ne, Mg2+

, F–, Na

+, O

2–

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9. Rank each of the following sets of elements in order of decreasing first ionization energy.

a. Cl, Br, I

b. Na, Li, Cs

c. Ga, Ge, Se

d. S, Cl, Br

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

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

a) non-polar covalent bond b) polar covalent bond c) ionic bond.

12. 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

13. 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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ANSWER THE FOLLOWING QUESTIONS ON ELECTRON CONFIGURATIONS: 1. Draw the Lewis structure for each of the following.

a. NH3

m. H2O

b. SO3

n. Cl2O

c. BrF

o. ClO3–

d. O2

p. CO32–

e. PO43–

q. SiF4

f. MgF2

r. BH3

g. NF3

s. HOBr

h. CH4

t. AsH3

i. BeF2

u. K2O

j. NO3-

v. HCHO

k. HCN

w. N2H2

l. N2

x. SeCl6

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2. What is meant by single, double and triple bond? 3. For diatomic O, N, and F molecules:

a. Draw the electron dot structure for the molecule.

b. draw orbital box (energy-level) diagrams for each atom.

c. Circle the electrons that are involved in bonding.

d. Describe the bond as single, double or triple.

4. For the following molecules: water, ammonia, methane, carbon dioxide, complete the following: a. Draw the electron dot structure for the molecule.

b. Draw orbital box (energy-level) diagrams for each atom.

c. Circle the electrons that are involved in bonding.

5. Draw the electron dot structure you would predict for the molecule carbon monoxide.

a. Does each atom have a stable octet? *Carbon monoxide has a coordinate covalent bond

(When necessary, a pair of electrons can be donated from one atom to make a needed covalent bond).

b. Which atom contributes both electrons to form the third bond in the necessary triple bond?

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Lewis Structures, VSEPR, Polarity, IM Forces

For each molecule:

a. Provide a Lewis Structure.

b. Determine the bonding and lone pairs and molecular shape.

c. Determine whether the molecule is polar or non-polar. (Does it have polar bonds? Is it symmetrical or

non-symmetrical?)

1) carbon tetrafluoride

2) CH2F2

3) O2

4) H2O

5) C2 H2

6) H2CS

7) NF3

8) BF3 9) carbonate ion 10) nitrate ion

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ANSWER THE FOLLOWING QUESTIONS ON VSEPR CONFIGURATIONS:

For each molecule:

a. Provide a Lewis Structure.

b. Determine the bonding and lone pair numbers and the molecular shape.

c. Determine whether the molecule is polar or non-polar (Does it have polar bonds? Is it symmetrical or

non-symmetrical?)

1. PF5

2. SF4

3. NO2-

4. XeF4

5. IO4-

6. NH4+

7. OCl2

8. CH2O

9. SeH2

10. CH3-

11. CH3+

25

ATOMIC UNIT REVIEW

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.

3. Compare and contrast line and continuous spectra.

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

5. Write complete electronic configurations for the following atoms and ions a. P

b. Ca

c. Cu d. Rh e. Sb

3+

f. Ni2+

6. How many unpaired electrons are there in each of the following: a. Mn

b. As

c. Sr

d. Tl+

7. For each of the following properties indicate which has the larger value fluorine or bromine.

i) atomic radius

ii) first ionization energy

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

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

a) non-polar covalent bond b) polar covalent bond

c) ionic bond.

10. 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

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For each molecule:

a. Provide a Lewis Structure.

b. Determine the bonding and lone pair numbers and the molecular shape.

c. Determine whether the molecule is polar or non-polar (Does it have polar bonds? Is it symmetrical or

non-symmetrical?)

1) OH-1

2) SF2

3) IBr3

4) HOF

5) CH2F2

6) H3P

7) AsF6-1

8) HCO3-1

27

28

29

30

Energy Sublevels Total # e- capacity

Level

1 s 1 2

2 s,p 1+3 = 4 8

3 s,p,d 1+3+5 = 9 18

4 s,p,d,f 1+3+5+7 = 16 32

n n types n2 2n2

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Group Electronegativity Table 1

1 2 H He

2.2 2

13 14 15 16 17

3 4 5 6 7 8 9 10 Li Be B C N O F Ne

1.0 1.6 2.0 2.5 3.0 3.4 4.0

11 12 13 14 15 16 17 18 Na Mg Al Si P S Cl Ar

0.9 1.3

3 4 5 6 7 8 9 10 11 12

1.6 1.9 2.2 2.6 3.2

19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

0.8 1.0 1.4 1.5 1.6 1.7 1.6 1.8 1.9 1.9 1.9 1.6 1.8 2.0 2.2 2.6 3.0

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

0.8 1.0 1.2 1.3 1.6 2.2 2.1 2.2 2.3 2.2 1.9 1.7 1.8 2.0 2.0 2.1 2.7

55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La* Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

0.8 0.9 1.1 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.8 1.9 2.0 2.2

87 88 89 104 105 106 107 108 109 Fr Ra Ac** Rf Db Sg Bh Hs Mt

0.7 0.9 1.1

32

33