e. chapter 5 electron packet

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The Electron Unit 5

Our Essential Question:

In our previous unit, The Atom, we looked for evidence to support what some have been saying for

over 2000 years: that the universe is composed of 82 naturally occurring types of atoms known as

elements. Evidence in support of this idea includes the experiments of Thomson and Rutherford, and

numerous images compiled since 1955.

But where are the electrons? Rutherford showed that they are in the vast space that exists outside

the nucleus of each atom. What purpose do the electrons serve? How are they organized? How do

they help explain the molecules, like water, that are all around us, and that we are made out of? To

find out we’ll complete our historical look at the atom with the discovery of Niels Bohr, aided by the

mathematical breakthroughs of Rydberg and Balmer.

Our plan for the week:

Day 1:

Lab: Flame Test Lab

Lesson: Light

Day 2:

Lab: Spectroscopy Lab

Lesson: Bohrs Epiphany

Day 3

Lesson: Electron Configuration

Lab: Problem Solving

Day 4

Review for Electron Test

Day 5:

Electron Test

Where are the electrons in an atom?

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1Unit 5 electrons Dr. B.’s ChemAdventure

Unit 5: Electrons

Where are the electrons in an atom?

X

And so First, we must understand…

The answer lies in this mystery:

?

Our Essential Question:


How does light travel?• What happens when we shine a

flashlight through a slit?

• What happens when we shine a flashlight through two slits?

?

3


S=wf

Young’s Double-Slit Experiment:

1. Light travels in waves Speed of light= 3 x 108 m/s

S-1Hz

2: really fast.

long wavelength

Low frequency

high frequencyshort wavelength

Speed of light (m/s) = wavelength (m) x frequency (1/s)

Instructions here; explanation here


• S=wf• 3 x 108 m/s = (w)(7.23 x 1014 s-1)

w = 3 x 108 m/s• 7.23 x 1014 s-1

• = 4.15 x 10-7 m =0.000000415 m• = 415 nm

What is the wavelength of violet light in nanometers;

f = 7.23 x 1014 s-1?L2 review scientific notation Here

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400-700 nm: Roy G. Biv

safe dangerous

often constant

The ElectromagneticSpectrum


Hydrogen

w(nm) =656, 486, 434, 410…What number is next??

Balmer

Rydberg

The story of Bohr’s Epiphany

−

=

22 n1

2101097.0

1wL1:Try it for n = 3

w = 656 nm

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X76543n?410435486656w

Bohr sees the connection

Between light and the electron

X397

−

=

22 n1

2101097.0

1w

Electron Emission Is Light


+P+P

22

33

44

656 nm

700 nm700 nm400 nm400 nm

--ee

Hydrogen Hydrogen AtomAtom

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--ee

--ee

55

--ee

486 nm

434 nm

--ee 410 nm

13:

absorption

3 2:Emission

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It all fits. Bohr • 1. Jimmy neutron: no

• Shells: yes• 2. evidence: emission

nucleus

28

1832


Where the electrons are exactly: Aufbau

order (to Argon)

(Soon)10And d3p663p3s22s,182p66and p2s22s,821s22s

3

1

# electrons

orbitals:# electrons

total

shelllevel

orbital # e’sConfig.

Video animation: 16:54- 17:20

Aufbau order to Ar:

Paired electrons

1s2 2s2 2p6 3s2 3p6

Nuc.1

23

2

again

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Electron Configuration of…1H: 1s1

2He: 1s2 3Li: 1s22s18O: 1s22s22p4

16S: 1s2 2s2 2p6 3s2 3p4

Orbitals and # electrons: s(2) p(6) d(10) f(14)

Orbital NotationA detailed way to show electron configuration

3Li: 1s2 2s1 Electron configuration

Orbital Notation

electrons pair up with opposite spinsFor each orbitalPauli Principle:

Hund’s Rule: electrons spread out within orbital groups


Orbital Notation of Carbon:

6C:

Hund’s Rule: Electrons spread out within orbitalsPlease give the electron configuration with orbital notation for Sulfur

1s2 2s2 2p2X

1s2 2s2 2p2No! Yes!

1s2 2s2 2p2 3s2 3p4

16S:

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Heisenberg’s Uncertainty Principle

• We can never know the exact position and velocity of an electron at the same time.

• Why?• (Even the smallest amount of energy

moves it unpredictably)


AufbauOrder

Note:4s< 3d

“Build up”

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Aufbau Order

7p66d105f147s26p65d104f146s25p64d105s24p63d104s23p63s22p62s21s2

12Mg: 1s22s22p63s2 9F: 1s22s22p5

21Sc:

41Nb:

1s22s22p63s23p64s23d1

1s22s22p63s23p64s23d104p65s24d3

[Ne]3s2 [He]2s22p5

[Ar]4s23d1 [Kr]5s24d3

[Ne][Ar][Kr][Xe][Rn]

“shorthandNotation”

[He]

Fill in blank table


Principles and rules of electron configuration

Pauli(opp. spins)

Hund’s Rule(spread out)

1s22s11s22p1Aufbau(build up)

Heisenberg(e-position uncertain)

GoodBadPrinciple or rule

1s2 2s2 2p2

1s2 1s2

1s2 2s2 2p2

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Valence Electrons• Outermost shell of electrons• Easily predictable from Periodic Table.1

2

2 3

2 22 2 2 2 2 2

4 5 6 78

2

All 2

(2)


Electron Dot Structures:• Quick look at valence electrons

Ne Li BeNo!: always spread

out all 8

Be

Try H,O,N,CCNH O

Valence electrons are the key to understanding:Chemical reactivity

Fortunately, it is nicely categorized in our next topic:The Periodic Table

X

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The Electron: Lab Experiment, Problem Sets, and Study Guide

This Unit includes two experiments, four problem sets, and a study guide. Bohr was fascinated by the brightly colored lights emitted by various chemicals when placed in a flame, and wondered if there was a chemical explanation. You wil create the same colors in the Flame Tests Lab, and use it to identify unknown samples. And between 1907-1913 Bohr and other scientists competed to explain the sharp lines of light that each element forms when heated or charged. Using spectroscopes, you will observe these lines as well in the Spectroscopy Lab, and will be asked to consider why they appear as they do for each element. Since there is a fundamental relationship between light and the electron, you will complete a problem set on Light, and this will be followed by several problems concerning how electrons are organized around the nucleus: Electron Configuration. This includes the idea of pairs, or orbitals organized around the nucleus, and shows us where they are most likely to be, including the direction of spin they are rotating in. We will see that it may be impossible to ever determine where they are and how fast they are moving at the same time.

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Names___________________________________________Period______________ lab5.1

Flame Tests Lab 20 points

Safety Notice: This lab is exciting, but please be cautious. Wear goggles. Assume all salts are toxic, as are all gases produced. Introduction: We have all seen the beautiful colors that can form when substances are placed in a flame. Pockets of gas in wood can form green and blue colors when they ignite. What is happening when this occurs? This answer was the key to unlocking the secrets of the electron, now known as quantum theory. In this experiment we will observe some of these colors, and will make some initial attempts to explain it. Finally, the color of the emitted light will be used to identify the unknown salts.

Materials: Bunsen burner; Paper Clips Beaker of water

Salts and unknowns:

1. __________ 2. __________ 3. __________ 4. __________

Unknown numbers:

Procedure 1. Goggles on please, and go to your stations. The rule all year will be that if the instructor is

wearing goggles, you are as well. 2. Listen to the Bunsen burner lesson 3. Each student should safely light the Bunsen burner and adjust the gas/air mixture. 4. Turn off Bunsen burner 5. Get your set of test tubes and unknowns 6. Dip the paper clip to see what color it turns on its own…if you see this color it may be due to the

metal in the paper clip. 6. Dip each of solutions using a paper clip and place in flame for less than 2 seconds each. Write

down the color of the flame, and estimate the wavelength in nanometers(a color chart will be available).

7. Fill in table 1 during testing. Identify the unknowns based on flame color. 8. Clean up: wet matches in trash. All stations will be inspected, including sinks. 9. Take your normal seats. 10. Answer questions at your desks. Turn in lab (one per group of two).

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Chart: Wavelength (in nanometers) of visible light

Table 1: Color and wavelength in nanometers of emission spectrum of salts and unknown.

Salt Flame color Estimated wavelength (nm)

1.

2.

3.

4.

Unknown #_____

Unknown #_____

Analysis 1. Each of the known compounds tested contains chlorine, yet each compound produced a flame of a

different color. What does this suggest? 2. We will learn this week that the movement of electrons in atoms produces the colors we observed.

What specifically may be going on with the electrons to produce color? (take your best guess)

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Team Names_________________ and ____________________ Period _____ Lab 5.2

Spectroscopy Lab

Introduction: In our previous lab we observed the vivid colors emitted by placing chloride salts in a flame. This was followed by a demonstration where we observed how a spectroscope (a prism, really) can divide up light into separate wavelengths. The purpose of this lab is to combine these two observations by repeating the flame test experiment, this time using a spectroscope. This experiment is similar to that performed by Niels Bohr and others, and begs the question: what does it all mean? How do the spectral lines relate to the structure of the atom? Safety: As before, this lab uses flames and toxic salts. Please wear goggles. Procedure:

1. Put on goggles. 2. Each group will perform a 5 minute experiment at each of 6 stations, and then proceed to the next. As precisely as you can, draw the component wavelengths observed at each station. Follow the instructions for each station, clean up, and be ready to move to the next station.

Station 1: Sunlight.

Each student should point through the spectroscope directly at the sun, and draw the component wavelengths observed. If weather permits, see if the colors are the same when you are not looking through a window.

Station 2: Artificial light

Each student should point through the spectroscope directly at the fluorescent lights, and draw the component wavelengths observed:

Station 3: Copper Chloride

Dip a paper clip into a copper chloride solution, and place it in the flame for less than two seconds while your partner observes the emission of light through the spectroscope. Repeat as necessary, but be cautious not to ignite the splint.

Station 4: Magnesium Combustion

Request a piece of magnesium metal from your instructor. Holding it in tongs, ignite the magnesium and observe the spectrum through the spectroscope. Warning: The light is extremely bright, and burns at 2000 degrees Celsius.

Station 5: Hydrogen gas

Turn on the hydrogen gas spectrum tube and observe the component wavelengths through the spectroscope. You should see individual spectral lines.

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Station 6: ______ Gas

Turn on the ________ gas spectrum tube and observe the component wavelengths through the spectroscope. You should see individual spectral lines. Data: Draw what you see through the spectroscope as accurately and precisely as you can. The marks are at 450, 550, and 650 nm.

Please answer the following questions at your normal seats: 1. Describe what you observed at each station:

1. 2. 3. 4. 5. 6.

2. Which light source provided the simplest spectrum? 3. Which light source provided the most complex or varied spectrum? 4. What were the wavelengths (in nanometers) of the individual lines from hydrogen in nanometers? 5. What were the colors of the individual lines from hydrogen? 6. Now that you have seen a variety of emission spectra, what do you believe causes the “lines”?

1. Sunlight 2. Fluorescent Light

700

3. Copper Chloride

4. Magnesium combustion 5. Hydrogen Gas

700

6._______

400 500 600 700 400 500 600 700 400 500 600 700

400 500 600 700 400 500 600 700 400 500 600 700

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Name_____________________________ Period_________ WS5.1

Wavelength worksheet

Please show your work, not just the answer . If you look down from Diamondhead in Hawaii, you will see waves rolling in at a steady rate. Some days they are nicely spread apart, meaning they have a long wavelength. Other days they come in more frequently; this is more dangerous for the surfers. The surfers prefer the long wavelength days. They know that as the wavelengths get shorter, their frequency gets higher, and there is more energy- more danger – to the high frequency waves. This is summarized in the diagram:

Light travels in the same way. It travels at a steady rate: about 300,000,000 meters per second, or 3 x 108 m/s. And as the wavelength decreases, the frequency must increase:

S = wf

S = speed of light = 3 x 108 m/s w = wavelength in meters (m) f = frequency in waves per second (Hz, or s-1)

In addition to a scientific calculator, you will need a wavelength chart to answer these questions. 2. An X-ray has a wavelength of 1.15 x 10-10 m. What is its frequency? 3. What is the speed and wavelength of an electromagnetic wave that has a frequency of 7.8 x 106 Hz?

Example. What is the frequency of green light, which has a wavelength of 4.90 x 10- 7 m?

Solution: 1-7-

8

s xxx m 10 x 4.90

m/s 10 x 3 ws f wf; s ====

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4. A popular radio station broadcasts with a frequency of 94.7 megahertz (MHz). What is the wavelength of the broadcast? (1 MHz = 1,000,000 Hz) 5. Cable television operates at a wavelength of about 1300 nanometers. Which wavelengths is this between in the Electromagnetic Spectrum? 6. Which is more dangerous, a radio wave or ultraviolet light? 7. The moon is 234,000 miles from earth. Light travels at 3 x 108 meters per second, and there are 1.62 kilometers in a mile.

When you shine a flashlight on the moon, how long does it take for the light to hit the moon? 8. The smallest particle of light is the photon. Max Planck discovered that the energy of light can be calculated, where it is simply equal to a constant number multiplied by the frequency of the light:

What is the energy of a photon of green light? (See question number 1)

9. What is the energy of a photon of light with a wavelength of 2 meters? 10. Since s = wf, and E = hf, can we calculate energy using wavelength, by combining the two formulas? Please show the combined formula. (Hint: note that f appears in both formulas).

E = hf Where E is the energy of the light in joules

h = Planck’s Constant = 6.626 x 10-34 joules .seconds f = the frequency of light in Hz (which is 1/seconds)

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Name____________________________ Period_______ WS5.2

The Bohr Model of the Atom Prior to the work of Niels Bohr, it was known that electrons existed outside of the nucleus, but beyond that very little was known. 1. What was the observation that Bohr based his research on? 2. The Balmer formula is : Solve this formula for n = 4. 3. The heart of Bohr’s discovery was that he was able to come up with real meaning to this formula. Draw a hydrogen atom with several energy levels (“shells”) around it and show electronic emission from the fourth shell to the second shell. 4. Draw diagrams indicating atomic emission and absorbance. 5. All of the visible atomic emissions for hydrogen enter the second energy level. What wavelength of light is emitted when an electron moves from the second energy level to the first energy level? What type of light is this?

−

=

22 n1

2101097.0

1w

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Name:_______________________________________ Period:______ WS5.3

Electron Configuration (L1 only) Directions: Draw the electron configurations with orbital notation for each of the following atoms. Example: Here is the electron configuration of Sulfur with orbital notation. 1. Scandium: 2. Gallium: 3. Silver: 4. Krypton: 5. Iron: 6. Bromine: 7. Californium 8. Write the electron configuration using shorthand notation of the following elements: a. sodium

1s2 2s2 2p2 3s2 3p4

16S:

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b. An oxygen anion, O- c. Radon 9. Two substances that have the same number of electrons are isoelectronic. For example, both the fluorine anion F- and neon have ten electrons, they are isoelectronic. a. The bromine anion is isoelectronic with what uncharged element? b. Argon is isoelectronic with which monocation?

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Name___________________________________ Period __________________ WS 5.4

Electron Configuration NOT! Worksheet (L1 only)

In this unit we have seen how the electrons are organized around the nucleus. It is a very detailed view of the electrons location, and various rules to help keep it all straight have been devised, and are shown below. In each problem below, the electron configuration is incorrect. Fix it, and explain what law or principle (not Principal!) was violated. EXAMPLE:

1. 1Hydrogen:

2. 17Chlorine

3. 39Yttrium (next page)

2s1

3s2 2p6 2s2 1s2

4d10 5s2 4p6 3d10 4s2 3p6 3s2 2p6 2s2 1s2

Unit 5 electrons Dr. B.’s ChemAdventure

Principles and rules of electron configuration

Pauli(opp. spins)

Hund’s Rule(spread out)

1s22s11s22p1Aufbau(build up)

Heisenberg(e-position uncertain)

GoodBadPrinciple or rule

1s22s22p2 1s22s22p2

1s2 1s2

Law Violated: Aufbau Principle Fixed:

Law Violated: __________ Fixed:

1s1

3p5

22

4. 8Oxygen 5. 106Seaborgium

2s2 1s2 2p4

6d4 5f14 7s2 6p6 5d10 4f14 6s2 5p6 4d10 5s2 4p6 3d10 4s2 3p6 3s2 2p6 2s2 1s2

Laws Violated: __________ Fixed:



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Name:_______________________________________ Period:______ WS 5.5

Electron configuration and orbital notation self test Chemical behavior is determined by electron position. It’s a simple statement, but it says a lot. Another way of saying it is “Chemistry is all about where the electrons are”. That’s why we’ve been spending the last week focusing on electrons. However, somehow it always seems to bog down in some weird world of 1s2 2s2 2p6, and the Pauli Principle, and we forget our goal: if we know where the electrons are we know how the substance will behave. Why Neon is stable, and sodium is very unstable, and in fact why all the elements and the substances they form behave the way they do. Let’s pick an element. We know that oxygen contains ___ protons. And since it is not charged, it contains _____ electrons. We know that ____ of the electrons occupy the first shell, and the other six are in the second shell. We know that the first shell consists of a _____ orbital that holds _____ electrons, and so we say that the electron configuration of that first shell is 1s2. For the second shell we have six electrons, and we have learned that the first two will occupy a ____ orbital, and the next four go into ____ orbitals. Thus the electron configuration of oxygen is____________________. We can go into more detail, and show the exact orbitals that the electrons are in, which even show the direction the electrons are spinning in. An atomic orbital is simply a ______ of electrons, and the Pauli Principle tells us that electrons prefer to pair up with _________ spins. The first shell of oxygen contains one orbital, which we draw with a box like this:_______, showing that the electrons are paired up with opposite spins. The second shell begins with one more orbital for the two electrons of the 2s subshell, for a total of four electrons so far. We have ______ more electrons in oxygen, and they will occupy the three p orbitals. We remember to apply _________’s rule and spread these electrons out as far as possible in those three boxes. Thus we can draw the electron configuration of oxygen with its orbital notation right above it: Note that this tells us that oxygen has four electrons in its outer (second) shell, and the two of them are unpaired….we also know from HONC that oxygen likes to form two bonds…a coincidence?? Let’s work out the electron configuration of nitrogen and see if we get three unpaired electrons: Nitrogen has _____ electrons, so the electron configuration with orbital notation is (be sure to spread out your p electrons): Does this orbital notation show 3 unpaired electrons?? If this makes sense, continue to the “how to ace it” guide.. If not, see me so we can do more examples.

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Howtoaceitunit5 How to ace the Electrons Exam

In this Unit our goal was to determine where the electrons are in atoms. To find out, we performed two experiments that revealed the sharp lines that excited pure elements produced. We then analyzed this data from a historical perspective, beginning with the work of Niels Bohr. For this we needed to review the properties of light, including frequency, wavelength, energy, and, common types. This involved the use of the speed of light equation (s = wf) and an understanding of the electromagnetic spectrum. We then showed how the key mathematical solutions of Balmer and Rydberg allowed Bohr to put it all together to postulate energy levels, where atomic emission explains light, and produces the spectral lines observed for all elements. This was followed by a detailed look at the electron around the nucleus. We found that not only do electrons reside in shells, there are also subshells or orbitals within each shell. We observed how they spread out within an orbital (Hund’s Rule), and even how they spin when near each other (the Pauli Principle). We learned the configurations of electrons for all elements following the Aufbau Order, and how to write it all down by electron position, configuration, or orbital notation. This can rapidly tell us how many electrons are in each shell and subshell, the spin of each electron, and the number of unpaired electrons. The limits of observation of the electron are a result of the Heisenberg Uncertainty Princliple, which states that it is impossible to measure the position and velocity of an electron simultaneously, due to the extreme sensitivity of the electron. Finally, we showed how valence is easy to determine using the periodic table, and that valence may be drawn using electron dot formulas, also known as Lewis Dot Formulas. During this study we found that the periodic table is well designed to show the number of valence electrons for any element. In our next unit we will apply this to our understanding of the periodic table. To dominate this test, review all of the material in his packet: The lessons, the labs, and the worksheets. Here is some of the key information you should know: To ace this exam you should know:

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1. Draw the symbols for Democritus, Aristotle, Paracelsus, Lavoisier, Dalton, Thomson, Rutherford, and

Bohr

2. What is the significance of each symbol? Try to assign one or two key words for each symbol.

3. What are the dangerous wavelengths of light?

4. How does light relate to electrons?

5. What is wavelength? Units?

6. What is frequency? Units?

7. Rearrange the speed of light equation to show what frequency is equal to.

8. The electromagnetic spectrum: what is it?

9. Frequency: how does it relate to energy and safety?

10. Wavelength- how does it relate to frequency?

11. Energy: which rays have the highest energy?

12. Safety: why are radio waves generally considered safe?

13. Types of radiation Really long waves include ___________ and _______________; really short waves include __________ and ____________. The ___________________ (long/short) waves are dangerous.

14. Convert 452 nanometers to meters (107 nm = 1m) 15. Use s = wf to find the frequency of 452 nm light. 16. (Level one only) The Balmer formula. Find it in your notes: 17. Significance

18. Solve for the n= 3 to n = 2 transition:

19. Atomic Emission Spectra: How did we observe it?

20. Emission vs. absorbance- what is the difference?

21. The Bohr model of the atom- draw a model

21.5 What is the difference between electron configuration, and orbital notation?

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22. Electron names to zirconium. For example, manganese has the symbol ____

23. L1 only: Electron configurations- all elements…do iodine using noble gas notation.

24. L1 only: Orbital notation: all elements. Do silicon. Include the number of valence electrons, and the

number of unpaired electrons.

25. The Heisenberg Uncertainty Principle. State what it is and why briefly.

26. L1 only: Orbitals: s, p, d, and f…how many electrons for each? How many orbitals for each?

27. L1 only: Aufbau principle. Give an example where it is broken, and fix it.

28. L1 only: Pauli exclusion principle. Break it and fix it.

29. L1 only: Hund’s Rule. Break it and fix it.

30. Lewis Dot Structures. Draw oxygen, for example

31. Valence Electrons. Do each column in the periodic table..

32. Why is it important to use scientific references, rather than websites, when writing a scientific

paper?

33. Where are the electrons in an atom?

e. chapter 5 electron packet

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