TEACHER NOTES 2011 UNIT 5 | 1 UNIT 5 – Models of Atomic Theory
Student Objectives
You have a team of enthusiastic teachers longing to help you be able to…
1. Perform calculations involving the speed of light (c), the wavelength (λ), and the frequency (f) of electromagnetic radiation (emr) 2. Perform calculations involving the energy, the wavelength (λ), and the frequency (f) of electromagnetic radiation. 3. Place forms of electromagnetic radiation in correct energy order on an electromagnetic spectrum. 4. Compare the energy, frequency, and wavelength of electromagnetic radiation. 5. Label an energy level diagram, indicating both excitation and relaxation 6. Describe the four quantum numbers used to define the region of space that has a 90% probability of finding an electron with a given energy. 7. Explain the information provided by each of the quantum numbers. 8. Distinguish between an energy level, sublevel, and an orbital. 9. List the allowable quantum numbers for each energy level. 10. Draw an orbital diagram using the Aufbau principle, Hund’s rule, and the Pauli Exclusion Principle. 11. Write the electron configuration for an atom.
LIGHT AND ELECTRON CONFIGURATION
Much of what we know about the atom has been learned through experiments with light; thus, you need to know some fundamental concepts of light in order to understand the structure of the atom, especially the placement of the electrons.
I. Light A. Characteristics of Light 1. Has "Dual" nature (or split personality) 2. There are times in which it behaves like a WAVE and other times when it behaves like a PARTICLE (called a photon)
B. Light as a WAVE: You must be able to describe its wavelength, frequency, and velocity (speed) 1. Speed of light in a vacuum is the most accurately known constant, c, in the universe.
c = 3.0 x 108 meters/sec OR 3.0 x 1017 nm/sec (assuming a vacuum)
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 2 2. Wavelength is represented by the Greek letter, lambda ( ). It is the length between corresponding parts of adjacent crests and can be expressed in ANY units of length you choose (feet, inches, meters, kilometers, etc)
3. Frequency is represented by the Greek letter, nu () or the letter f. It is the number of wave crests which pass a given point in 1 second. Its units are:cycles / second, sec-1,or Hertz (Hz).
II. Electromagnetic Spectrum:
A. Visible Light: considered to have a wavelength of between 700 nm and 400 nm. The longest wavelength in visible is red (about 700 nm) and the shortest visible wavelength is violet (about 400 nm). The only thing that makes one color of light different from another is its wavelength and frequency; the velocity is always the same (the speed of light!).
B. Continuous Spectrum: Has all the energies of visible light. The colors of visible light in order from longest wavelength to shortest spell out: ROY G BIV.
C. Entire Range of light energy: Includes ALL energies of light. Become very familiar with the electromagnetic spectrum!!!
Example 5-1: Label the following electromagnetic spectrum from low energy on the left hand side to high energy on the right hand side.
AM/FM& TV MICRO- INFRA- VISIBLE ULTRA- X-RAYS GAMMA (RADIO) WAVES RED (IR) VIOLET RADIATION (UV) Long Short
700 nm R O Y G B I V 400 nm Low f High f
Low E I better memorize High E this by the next quiz… argh!
RED MARTIANS INVADE ROYGBIV USING X-RAY GIZMOS
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 3 III. Light as a WAVE:
c = f where c = speed of light (3.0 x 108 m/sec OR 3.0 x 1017 nm/sec)
Note: BE CAREFUL OF THE UNITS YOU USE!!! If speed of light is in meters, then wavelength must be in meters.
Example 5-2: Does a longer or a shorter wavelength have the HIGHER frequency? Shorter wavelength
Example 5-3: Are wavelength and frequency directly or inversely related? Inverse
Example 5-4: If the wavelength is known to be 550 nm, what is its frequency?
3.0x1017 f 5.45x1014s 1 550
Example 5-5: If the frequency of light is known to be 9.45 x 1014 s-1, is the light visible? HINT: Calculate the wavelength in nm.
3.0x1017 14 317nm 9.45x10 NO, it is not visible
Example 5-6: If the wavelength of light is 7.23 x 10-5 cm, is it visible?
NO, it is not visible
1x10 2 m 1nm 7.23x10 5 cm 7.23x10 2 nm 9 1cm 1x10 m
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 4 IV. Light as Particles: When light behaves as a particle, it is called a PHOTON. We are interested in how much energy they have. The unit of energy we will use is the Joule, J, and the equation which allows us to calculate the energy of one photon is:
hc E hf OR E where h= Planck's Constant = 6.63 x 10─34J•sec photon photon
Example 5-7: What is the energy of a photon of light whose frequency is 7.85 x 1015 Hz? Is this visible? E 6.63x10347.83x1015 5.19x1019J/photon
3.0x1017 38.2 nm NOT Visible 7.85x1015
Example 5-8: If a light has a wavelength of 550 nm, what is the energy of one photon of this light?
34 17 6.63x10 3.0x10 19 Ephoton 3.62x10 J/photon 550
Example 5-9: If one photon of light is known to have energy of 3.33 x 10─19 J, is it visible?
hc 6.63x10343.0x1017 -19 597nmYES E photon 3.33x10
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 5 V. Bohr Atom
Don’t Forget!! Bohr said that electrons are in energy levels!!! But he also proposed the planetary model. Silly ole Bohr.
A. electrons possess certain fixed amounts of energy B. electron cloud made up of energy levels which are like the rungs of a ladder BUT they are not EQUALLY SPACED
C. amount of energy possessed by an electron determines its energy level D. levels farther from nucleus represent higher energy E. in order to move from one energy level to another, electrons must gain or lose an exact amount of energy known as a Quantum.
**The ground state is where the electron normally hangs out. The excited state is where the electron has more energy than it normally has (meaning it has gained some quanta of energy!)
n= ___5 n= ___4
n= ___3
n= ___2
n= ___1
e- ____gains energy e- ____loses EMR (energy) Process______Excitation, light Process______Relaxation, light
absorption emission
GROUND STATE: GroundAll electrons State:______in lowest energy available EXCITED STATE: When an electron temporarily moves to a higher energy level
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 6 VI. Emission Spectra v. Absorption Spectra
Niels Bohr studied the emission spectrum of hydrogen to determine that electrons are on different energy levels. Every element has a unique emission/absorption spectrum. This is one way to identify an element.
Emission Spectra: the light emitted as electrons fall to lower energy levels Electrons emit photons of a specific energy. If the wavelength is in the visible range, we see bands of light against a black background.
Absorption Spectra: the light absorbed by electrons as they move to higher energy levels Electrons absorb photons of a specific energy. If the wavelength is in the visible range, we see black bands in the rainbow.
VII. Schrodinger Model: Quantum Mechanical Model (1926) A. AGREES with Bohr about quantized energy levels
B. DIFFERS: does not define an EXACT path called an orbit. The math describes a region in space with a high PROBABILITY of finding an electron. Unfortunately, we still call this an ORBITAL.
ELECTRON LOCATION: The location of the electron is described by 4 quantum numbers: n, l, ml, and ms. Think of these as being the “address” for an electron’s probable location! (State, city, street, house number… each gets more and more specific.) n= ENERGYLEVEL (Principle Quantum Number)
The higher the value of n… the further away from the nucleus the electrons are, and the higher the energy level The maximum number of electrons in any energy level: #e- = 2n2 Nucle us (where n=energy level) n = 1
n = 2 n = 3 Example 5-10: Fill in the chart below. n = 4
Energy n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7 Level (n) Maximum 2 8 18 32 50 72 98 number of electrons
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 7 l = SUBLEVELS (Second Quantum Number)
Each energy level is divided into SUBLEVELS Sublevels are regions in space with a specific SHAPE. Sublevels are labeled by the letters s,p,d,f,g,h… (The only ones we are going to work with and that you need to know are s, p, d, and f.) Electrons in the sublevels have a specific amount of energy. Energy of electrons in the sublevels: s < p < d < f. (s-electrons have the least energy; f- electrons have the most energy.) Each energy level (n) has one more sublevel than the previous energy level.
Here is a trick for remembering the order of the sublevels:
Silly People Do Flips!
Example 5-11: Fill in the chart below.
Energy Level (n) Sublevels Present 1 s 2 s, p 3 s, p, d 4 s, p, d, f
mℓ = ORBITALS (Magnetic Quantum Number)
Each sublevel contains ORBITALS which are regions of high PROBABILITY for finding electrons… the mℓ indicates the ORIENTATION of those orbitals! Each orbital can hold only TWO electrons The surface of the orbital is drawn to represent area where any particular electron can be found 90% of the time.
Example 5-12: Fill in the chart below. Sublevel # of orbitals Maximum # of electrons on Shape sublevel (= # orbitals x 2) s 1 2 Spherical p 3 6 Dumb-bell d 5 10 Weird! f 7 14 Even weirder!
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 8
s sublevel = 1 orbital = 2 electrons
p sublevel = 3 orbitals = 6 electrons
d sublevel = 5 orbitals = 10 electrons
f sublevel = 7 orbitals
= 14 electrons
Example 5-13: Complete the table:
Energy Sublevel(s) Orbitals (use lines to # e─ per sublevel Total # e─ per Level depict) energy level 1 s ____ 2 2
2 s ____ 2 8 p ______6
3 s ____ 2
p ______6 18
d ______10
4 s ____ 2
p ______6 32 d ______10
f ______14
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 9 ms= SPIN Quantum Number
The electrons within an orbital must have OPPOSITE spins to overcome their repulsion, which we represent with opposite arrows. Spin is either + ½ (______) or - ½ (______)
We are going to use these Quantum Numbers to describe the location of electrons, like your ticket describing your seat for a sporting event! Each gets more and more specific. Think of it like this:
n = stadium level (platinum, nosebleed) = energy level
l= section number = sublevel
ml= row = orbital
ms = seat number = spin
We will look very soon at how to use all of these numbers to describe an electron’s probable location, just like level, section, row, and seat number would help us find you at the game (assuming of course you are not out getting nachos!) !!!
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 10 VIII. Electron Configurations and Orbital Diagrams: Knowing the rules for electron arrangement allows us to write what is called electron configuration which gives a description of where in that atom the electrons are actually located. We can also draw an orbital diagram which gives the same information with more detail. Cool SAT word!
Orbital Diagrams provide a lot of information, but they become cumbersome to draw. Chemists use electron configurations to help model that structure of an atom or an ion. This in turn will help us with bonding in the next unit. Electron configurations show arrangements of electrons around the nucleus. Each element has its own electron configuration. Electron configuration gives specifically the energy level, sublevel, and number of electrons in that sublevel.
The outermost or valence electrons are the electrons that are in the highest energy level, “n”.
The highest energy electrons are in the last sublevel that was filling.
So, the valence electrons do not necessarily have the highest energy, even though they are
farthest away from the nucleus!
Rules for Arranging Electrons
1. Aufbau Principle Electrons are placed in orbitals of LOWEST energy first.
In a perfect world we would have the following ENERGY order: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f etc. UNFORTUNATELY, things get a little mixed up in terms of energy. The diagram to the right shows the filling order for sublevels.
2. Pauli Exclusion Principle:
An orbital may hold only TWO electrons. Think of it like an “exclusive club” with only 2 members! Why can’t we have an orbital with more electrons? Well, NO TWO ELECTRONS CAN HAVE THE SAME 4 QUANTUM NUMBERS. The combination of 4 quantum numbers is for that electron only! If they have the same energy (n), are in the same sublevel (ℓ), and are in the same orbital (mℓ), then they must have opposite spins. REMEMBER: We designate those spins with arrows! ↑ for +1/2 ↓ for -1/2
3. Hund’s Rule: When filling a sublevel (which contains orbitals of equal energy), one electron enters each orbital SEPARATELY (and with the SAME SPIN) until all the orbitals have one electron, then they PAIR UP!
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 11 Orbital Diagrams Notice that we have the same spin!
5p 4d 5s 4p 3d 4s 3p 3s
2p 2s
1s
Example 5-14: The orbital diagram for Zirconium (Zr) is drawn above. Answer the following questions about zirconium. a) How many electrons does zirconium have? 40 b) Why can’t the 2 electrons in the 4d sublevel be placed in the same orbital? Hund’s Rule c) In which sublevel are the electrons with the highest energy? 4d d) In which sublevel are the valence electrons? 5s e) How many valence electrons does rubidium have? 1 f) Electrons are always removed from or added to the sublevel that is farthest from the nucleus (the valence electrons!) to form an ion. Circle the electron(s) in the orbital diagram that is/are removed when zirconium becomes the zirconium ion.
Example 5-15: Draw the orbital diagram for sulfur.
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 12 PERIODIC TABLE and ELECTRON CONFIGURATION/ORBITAL DIAGRAMS
Periodic Table of the Elements Group 1 (IA) Group 18 (VIIIA) 1 2 1 2 1 1s 1s
Group 2 (IIA) Group 13 (IIIA) Group 14 (IVA) Group 15 (VA) Group 16 (VIA) Group 17 (VIIA) 3 4 5 ATOMIC NUMBER 5 6 7 8 9 10 1 1 3 4 6 2 2s 2p 2p 2p 2p
11 12 13 14 15 16 17 18 2 3 3s
Group 3 (IIIB) Group 4 (IVB) Group 5 (VB) Group 6 (VIB) Group 7 (VIIB) Group 8 (VIIIB) Group 9 (VIIIB) Group 10 (VIIIB) Group 11 (IB) Group 12 (IIB) 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 2 2
4 3d 4p Period
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 1 3 5 8 5 5 5s 4d 4d 4d 5p
55 56 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 2 10 3 6 6 6s 5d 6p 6p
87 88 103 104 105 106 107 108 109 110 111 112 113 114 115 116 4 1 4 7 6d 7p 7p
57 58 59 60 61 62 63 64 65 66 67 68 69 70 4f1 4f14
89 90 91 92 93 94 95 96 97 98 99 100 101 102 5f13
Example 5-16: Fill in the missing outer configurations in the periodic table above.
Example 5-17: Answer the following question about the periodic table above.
a) How many electrons does an s- sublevel hold? 2
b) How many elements are in a row in the “s” block? 2 Hmmm…
c) How many elements are in a row in the “p” block? 6
d) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “p” block? 3 Hmmm…
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 13 e) How many elements are in a row in the “d” block? 10
f) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “d” block? 5 Hmmm…are you noticing a recognizable pattern?
g) How many elements are in a row in the “f” block? 14
h) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “f” block? 7 Hmmm…
i) What is the correlation between the period number and the energy level, “n”, for the “s” and “p” blocks? n = period number j) What is the correlation between the period number and the energy level, “n”, for the “d” block? n = period number ─ 1 k) What is the correlation between the period number and the energy level, “n”, for the “f” block? n = period number ─ 2
Example 5-18: Work the following regarding Zn.
a) Write the electron configuration for the Zn atom. 1s22s22p63s23p64s23d10
b) Draw an orbital diagram for Zn. Draw energy increasing from left to right rather than down to up.
_____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____
1s 2s 2p 3s 3p 4s 3d
c) In which sublevel are the valence electrons? 4s
d) How many valence electrons does zinc have? 2
e) The electrons with the highest energy are in which sublevel? 3d
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 14 Example 5-19: Work the following regarding bromine.
a) Write the electron configuration for the Br atom. 1s22s22p63s23p64s23d104p5
b) Draw the orbital diagram for bromine.
_____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____ _____
1s 2s 2p 3s 3p 4s 3d
_____ _____ _____
4p
c) Which electron(s) is/are the highest-energy electrons? 4p5
d) Which electron(s) is/are physically farthest from the nucleus? 4p5
Example 5-20: Write the electron configuration for the following elements:
a) Carbon 1s22s22p3
b) Sodium 1s22s22p63s1
c) Sulfur 1s22s22p63s23p4
d) Calcium 1s22s22p63s23p64s2
Example 5-21: Identify the element that has the given electron configuration.
a) 1s22s22p63s23p3 P
b) 1s22s22p63s23p64s23d104p65s24d7 Ru
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 15 ABBREVIATED ELECTRON CONFIGURATIONS
Writing these long, repetitive configurations can be a pain in the neck and takes up a lot of space. Space is at a premium on a periodic table, so scientists developed an abbreviated form.
Example 5-22: Compare the electron configurations below. What do the configurations of these elements have in common?
antimony (#51) 1s22s22p63s23p64s23d104p65s24d105p3
palladium (#46) 1s22s22p63s23p64s23d104p65s24d8
They are identical up to the point of 5s2
To write abbreviated electron configurations, look to the noble gas in the prior period. Brackets around the symbol indicate that the element has the same electron configuration of that noble gas PLUS whatever follows.
MODEL: Fe: 1s22s22p63s23p64s23d6 BECOMES: [Ar] 4s23d6
Example 5-23: Write the abbreviated electron configurations for the following elements.
a) antimony (#51) [Kr] 5s24d105p3
b) palladium (#46) [Kr] 5s24d8
Note: These are great, but you can only write them when explicitly instructed!
CHEMISTRY TEACHER NOTES 2011 UNIT 5 | 16 IX. LEWIS DOT STRUCTURES OF ELEMENTS
A. Valence Electrons: Outer s and p electrons – the electrons most likely to be involved in bonding.
Valence electrons are most helpful for predicting the oxidation # (often referred to as charge) for elements in the s and p blocks. The To be studied in oxidation number is a method used for tracking electrons, and is equal to the ionic charge for cations (+ ions) and anions (— ions). our next unit!
B. LDS of ATOMS: To draw the LDS (Lewis Dot Structure) for atoms, simply imagine a square around the element symbol and then place the valence electrons around the symbol. Remember Hund’sRule: put one in each spot, then double up!
Example 5-24: Draw the Lewis Dot Structures for the Main Group elements (Groups 1, 2, 13 – 18) in the periodic chart below.
I only have 2 electrons total, so I can’t have 8 valence electrons like everybody else!
CHEMISTRY