S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 5 | Page 1

NAME__________________________________________ PERIOD____________ UNIT 5 NOTES: MODELS OF ATOMIC THEORY

STUDENT OBJECTIVES: Your fascinating teachers would like you amazing learners to be able to… 1. Perform calculations involving the speed of light (c), the wavelength (λ), and the frequency (ν) of

electromagnetic radiation (EMR). 2. Perform calculations involving the energy, the wavelength (λ), and the frequency (ν) 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 EMR. 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 and atom or an ion. 12. Be able to identify the highest-energy electrons, the valence electrons, and the number of valence electrons. 13. Using the periodic table, describe the periodic trends in atomic radii, ionic radii, ionization energy, reactivity,

and electronegativity.

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LEGGETT Atom EMR 5-1 2013 (6:56) http://youtu.be/jJdSr25-Q_U http://vimeo.com/73061084

I. LIGHT AND ELECTRON CONFIGURATION:

Much of what we know about the atom has been learned through experiments with light; thus, we need to know some fundamental concepts of light in order to understand the structure of the atom, especially the placement of the electrons.

1. Characteristics of Light: Has "Dual" nature (or split personality) - there are times in which it behaves like a continuous WAVE and other times when it behaves like a PACKET of light (a blip of light!!!)

a. Light as a Wave i. 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)

ii. Wavelength is represented by the Greek letter, ______________ ( ): 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)

iii. Frequency is represented by the Greek letter, ______________ ( ): It is the number of wave crests which pass a given point in 1 second. Its units are, most commonly, cycles / second , simply sec-1. The unit, Hertz (Hz) is also used as a unit label.

b. Electromagnetic Spectrum i. Any visible light will have a wavelength of between ____________ nm and ____________ 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 and - the velocity is always the same (the speed of light!). The colors of visible light in order from longest wavelength to shortest spell out ROY G BIV.

ii. NOT ALL LIGHT IS VISIBLE! The entire range of light energy is considered the Electromagnetic Spectrum.

Example 5-1. Label the following electromagnetic spectrum from lowest energy on the left hand side to highest energy on the right hand side.

700 nm R O Y G B I V 400 nm

RED MARTIANS INVADE ROYGBIV USING X-RAY GIZMOS

http://images.wikia.com/looneytunes/images/b/b7/Daffy-duck-looney-tunes-bugs-

bunny-duck-dodgers1.jpg

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Leggett PAP Chem Atomic Models 5-2 (14) http://youtu.be/ldjnjFyxXms http://vimeo.com/30786351

Example 5-2. Do you think a longer or shorter wavelength would have the HIGHER frequency?

Example 5-3. What mathematical relationship can you draw about wavelength and frequency?

II. CALCULATIONS WITH LIGHT

Note: BE CAREFUL OF CHOOSING SPEED OF LIGHT FOR YOUR FORMULA!!! You must make the units of

wavelength match up with your units for speed of light. Remember that “c” can also be 3.0 x 1017 nm/s.

Units of frequency can be reported as 𝟏 𝒔⁄ (per second), second –1, or Hertz. They are interchangeable with

each other!

Example 5-4. If the wavelength is known to be 550 nm, what is its frequency?

Example 5-5. If the frequency of light is known to be 9.45 x 1014 s-1, is the light visible? If not, is it UV or IR? HINT: Calculate the wavelength in nm.

Example 5-6. If the wavelength of light is 7.23 x 10-5 nm, what is the frequency of light?

Example 5-7. What is the energy of a photon of light whose frequency is 7.85 x 1015 Hz? Is the light visible? How do you know?

Example 5-8. If a light has a wavelength of 550 nm, what is the energy of one photon of this light?

Example 5-9. If one photon of light is known to have energy of 3.33 x 10–19 J, is it visible?

FORMULAS ON THE PERIODIC TABLE:

c = Ephoton = h λ

hcEphoton

CONSTANTS ON THE PERIODIC TABLE:

c = 3.0 x 108 m/s

(speed of light in meters)

h= 6.63 x 10-34 J•sec

(Planck’s constant)

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LEGGETT Atom Bohr & AES 5 3 2013 (13:49) http://youtu.be/kao4hoyCU_c http://vimeo.com/73061083

III. BOHR ATOM: Don’t Forget! Bohr said that electrons are in energy levels – which is true! And his

research with light & electron energy levels is very important and lead to a lot of what

we know about the atom! But he also proposed the planetary model – which is not

true.

a. Electron cloud made up of _______________________________________ which are like the ______________ of a ladder and are BUT they

are not ____________________________________.

b. Amount of energy possessed by an electron determines its ___________________________________.

c. Levels farther from nucleus represent _________________________ energy.

d. Electrons like to hang out where they are most __________________, which is their _______________ __________________. This is the

exact amount of energy they like to possess – it’s the electron’s home. All electrons have their own ground state – on Level

1, Level 2… or Level 8!

e. When electrons gain a “quanta” (amount) of energy (through sources such as extreme heat or electricity), the electron will

____________________ up energy levels in a process called _________________________. This excited state is very

_____________________________!!!

f. Electrons can’t stay excited at the higher energy level forever! When they fall back to their ground state, they will lose the

extra energy in the form of a packet of light emitted – in a process called ________________________________. The larger the fall,

the larger energy of light released!

n= ___

n= ___

n= ___

n= ___

n= ___

e- ____ energy

Process___________

e- ____ EMR (energy)

Process___________

Ground State:__________________________

n= ___

n= ___

n= ___

n= ___

n= ___

e- ____ energy

Process___________

e- ____ EMR (energy)

Process___________

Ground State:__________________________

When the electrons lose the energy

they gained, they release LIGHT

(EMR)…

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g. Each possible energy fall corresponds to a __________________________________________________ amount…

which means that it corresponds to a different _____________________________________ of __________________.

Some of the falls can be detected by the human eye (as they fall in the visible range), but other falls

that are not visible must be detected by instruments (as they are infrared or ultraviolet).

h. Different elements have different energies associated with their energy levels. This means that

similar falls in __________________________________________ will produce __________________________

______________________.

i. When we observe a color being produced from an element undergoing relaxation, we are observing

all of the VISIBLE falls simultaneously occurring. We can use things like a prism or a spectroscope

to break apart that light into its individual colors – resulting in a _______________________________________!

j. We can use the resulting bright-line spectra to identify an element!

Example 5-10. Based on the Bright-Line Spectra shown above, identify the elements found in the mixture.

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Leggett PAP Chem Atomic Models 5-4 (12) http://youtu.be/_nkXrbKnxuU http://vimeo.com/30789661

IV. SCHRODINGER MODEL: Quantum Mechanical Model (1926)

a. Schrodinger ___________________________ with Bohr about quantized energy levels.

b. However, Schrodinger ____________________ in the fact that he does not define an _______________ path called an orbit.

Schrodinger’s model describes a region in space with a high ___________________ of finding an electron.

Unfortunately, we still call this an ______________________.

ANALOGY: It is probable that you will be in the region of space defined by my classroom during this specified

period BUT we don’t know where you are in the room and how you are moving around the room. We’re just 90%

sure we will find you (maybe you are absent, in the office, etc.). My classroom is the orbital!

c. ELECTRON LOCATION: Electrons are distributed around the nucleus in specific ways and each electron has

four coordinates to describe its most probable location or distribution: 4 quantum numbers, n, l, m, and s.

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 = __________________________________________ (Principle Quantum Number)

Indicates how far from the nucleus the electron is located and its energy. It is designated by the Principle Quantum Number, n.

The higher the value of n, the ____________________ away from the nucleus and the _________________ the energy level

The number of electrons in any energy level: ______________________________ (n=energy level)

Example 5-11. Fill in the chart with how many electrons each energy level can hold.

First Energy Level Second Energy Level Third Energy Level Fourth Energy Level

Leggett PAP Chem Atomic Models 5-5 (8.5) http://youtu.be/0vjpiQ3mIoI http://vimeo.com/30789797

l = _______________________ (Second Quantum Number)

** Each energy level is divided into ___________________________ which

are designated by the Second Quantum Number, l. Actually not

a number at all, but rather a letter!

** Indicates the _______________________ of the region.

1. Sublevels are labeled by the letters s, p, d, and f.

2. The first energy level will have 1 sublevel and its label is "s".

3. The 2nd energy level will have 2 sublevels; "s" and "p".

4. The 3rd energy level will have 3 sublevels; "s" , "p", and "d"

5. The 4th energy level will have 4 sublevels; "s", "p", "d", "f"

Here is a trick for remembering the

order:

Stupid People Die First….

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“s” sublevels have a _____________________ shape. “p” sublevels have a _____________________ shape. “d” and “f”

sublevels have really weird shapes!!!

http://solarwiki.ucdavis.edu/@api/deki/files/85/=Single_electron_orbitals.jpg

Leggett PAP Chem Atomic Models 5-6 (10) http://youtu.be/VEUW5al2jFw http://vimeo.com/30953891

mℓ = ________________________ (Third/Magnetic Quantum Number)

Each sublevel contains __________________ which are regions of high ____________________________ for finding

electrons… the mℓ indicates the _____________________________________ of those orbitals!

The number of orbitals is related to the sublevels (ℓ)

Each orbital can hold only ______________ electrons

The surface of the orbital is drawn to represent area

where any particular electron can be found

______________ of the time.

Each sublevel has a set number of orbitals:

s sublevel - ________ orbital

p sublevel - ________ orbitals

d sublevel - ________ orbitals

f sublevel - ________ orbitals

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Leggett PAP Chem Atomic Models 5-7 (7) http://youtu.be/nlLtXoKpd2A http://vimeo.com/30954046

Example 5-12. Complete the table:

Energy Level Sublevel(s) Orbitals (use lines to depict) # e– per sublevel # e– per energy

level

1

2

3

4

ms = _____________________ (Fourth Quantum Number)

** The electrons within an orbital must have _____________________ spins to overcome their repulsion, which we represent with opposite arrows.

** Spin is either + ½ (______) or - ½ (_____)

We are going to use these FOUR Quantum Numbers to describe the location of electrons, like your ticket describing your seat for a sporting event! Each get more and more specific. Think of it like this:

n = stadium level (lower, platinum, nosebleed) = energy

level

ℓ = section number = sublevel

mℓ = row = orbital

s = seat number = spin

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LEGGETT Atom Models 5 8 2013 (12:46) http://youtu.be/ngE5R2Ke2KI http://vimeo.com/73061085

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)!!!

V. ELECTRON CONFIGURATIONS AND ORBITAL DIAGRAMS Knowing the rules for electron arrangement allows us to write what is called the 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. ____________________________________ provide a lot of information (since they show ALL FOUR quantum numbers), but they become annoying to draw.

______________________________________ are a shorter way of showing arrangements of electrons around the nucleus. They don’t provide all of the information that an orbital diagram would, but they do show us quite a bit! Each element has its own electron configuration. Electron configuration gives specifically the energy level, sublevel, and number of electrons in that sublevel.

WE HAVE SOME BACKGROUND INFO WE NEED TO GO OVER BEFORE WE CAN WRITE THESE ELECTRON CONFIGURATIONS & ORBITAL DIAGRAMS…

Remember that “n” provides information on both energy and distance. The relationship between “n” and DISTANCE stays true! Electrons in higher “n” energy levels are more

likely to be found further from the nucleus. 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 – meaning it doesn’t fill up in the order

we quite expect. We will learn the order for filling sublevels in just a minute. FOR NOW: trust us! We are going to ultimately show you that the periodic table provides us with the

perfect arrangement to break the code of filling order!

The _________________ electrons are the electrons that are in the highest energy level, “n”.

The ________________________________ electrons are in the last sublevel that was filling.

ALSO: There are some rules we must follow for arranging electrons…

Aufbau Principle: Electrons are placed in orbitals of ______________ energy first.

Pauli Exclusion Principle: An orbital may hold only ______________ 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 – meaning 4 different quantum numbers. REMEMBER: We designate those spins with arrows! ↑ for +1/2 ↓ for ─1/2

Hund’s Rule: When filling a sublevel (which contains orbitals of equal energy), one electron enters each orbital

until each _____________________, then they _______________________!

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ABOUT IONS:

Most metals will ______________ electrons to become ____________________________ with the nearest noble gas. When they lose electrons, the ion will be smaller than the original atom.

Non-metals will ______________ electrons when they are with metals to become _____________________ with the nearest noble gas. When they gain electrons, the ion will be larger than the original atom.

In Noble gases, electrons have maximized attraction to the nucleus and minimized repulsion with each other as much as possible, due to their electron configuration. They are STABLE and do not react – aka INERT. When substances become ions, they want to copy-cat the ideal electron configurations of the noble gases!!!

Example 5-13. For each of the following, predict the charge of the most likely ion and indicate which sublevel gains or loses electrons.

Group

Electron Configuration of

Valence Electrons?

Number of Valence

Electrons?

Lose or Gain Electrons to

look like Noble Gas?

Common Charge of Ions?

Would resulting ion be larger or smaller than the original atom?

1

2

13

15 non-metals

16 non-metals

17 non-metals

http://mr-moss.wikispaces.com/file/view/noble_gases.jpg/134023945/394x275/noble_gases.jpg

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PERIODIC TABLE and ELECTRON CONFIGURATION/ORBITAL DIAGRAMS

Example 5-14. Fill in the missing outer configurations in the periodic table above.

Example 5-15. Answer the following question about the periodic table above.

a) How many electrons does an s- sublevel hold? ______

b) How many elements are in a row in the “s” block? ______ Hmmm…

c) How many elements are in a row in the “p” block? ______

d) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “p” block? ______ Hmmm…

e) How many elements are in a row in the “d” block? ______

f) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “d” block? ______ Hmmm… are you noticing a recognizable pattern?

g) How many elements are in a row in the “f” block? ______

h) Remember that each orbital can hold TWO electrons…How many orbitals would be needed for the “f” block? ______ Hmmm…

i) What is the correlation between the period number and the energy level, “n”, for the “s” and “p” blocks?

j) What is the correlation between the period number and the energy level, “n”, for the “d” block?

k) What is the correlation between the period number and the energy level, “n”, for the “f” block?

Group 1 (IA) Group 18 (VIIIA)

1 2

1s1

1s2

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

2s1

2p1

2p3

2p4

2p6

11 12 13 14 15 16 17 18

3s2

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

3d2

4p2

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54

5s1

4d3

4d5

4d8

5p5

55 56 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86

6s2

5d10

6p3

6p6

87 88 103 104 105 106 107 108 109 110 111 112 113 114 115 116

6d4

7p1

7p4

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

Periodic Table of the ElementsP

eri

od

1

2

3

4

5

6

7

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Example 5-16. (a) Write the electron configuration for Cobalt. Then, fill out the orbital diagram above for Cobalt.

(b) What are the highest-energy electrons in Cobalt?

(c) What are the outer-most (valence) electrons in Cobalt?

(d) How many valence electrons does Cobalt have?

1s

2s

3s

4s

5s

2p

3p

4p

5p

3d

4d

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Example 5-17. Work the following regarding Sulfur.

(a) Write the electron configuration.

(b) Write the orbital diagram across from left to right (rather than upwards).

(c) What are the highest-energy electrons?

(d) What are the outer-most (valence) electrons?

(e) How many valence electrons are there?

(f) What is the most likely ion for Sulfur? Did we lose or gain electrons to form the ion?

(g) What would be the new electron configuration for the Sulfur ion?

(h) Would the resulting ion be larger or smaller than the original atom?

(i) With which noble gas is the Sulfur ion isoelectronic?

Example 5-18. Work the following regarding Strontium.

(a) Write the electron configuration.

(b) Write the orbital diagram across from left to right (rather than upwards).

(c) What are the highest-energy electrons?

(d) What are the outer-most (valence) electrons?

(e) How many valence electrons are there?

(f) What is the most likely ion for Strontium? Did we lose or gain electrons to form the ion?

(g) What would be the new electron configuration for the Strontium ion?

(h) Would the resulting ion be larger or smaller than the original atom?

(i) With which noble gas is the Strontium ion isoelectronic?

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Example 5-19. Answer the following regarding the Aluminum isotope above.

(a) What is the most likely ion for Aluminum? Did we lose or gain electrons to form the ion?

(b) What would be the new electron configuration for the Aluminum ion?

(c) Would the resulting ion be larger or smaller than the original atom?

(d) With which noble gas is the Aluminum ion isoelectronic?

NUCLEUS

Isotope: ALUMINUM – 27 (NO CHARGE) Nuclear Symbol:

Atom or Ion? _____________________________________ ______________________

# of Protons: _____________ # of Neutrons: _____________ # of Electrons: _____________

Electron Configuration: _____________________________________________________________________

Orbital Diagram:

Circle the Highest-Energy Electrons.

Box the Valence Electrons.

Number of Valence Electrons:

_____

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Example 5-20. Answer the following regarding the Bromine isotope above.

(a) What is the most likely ion for Bromine? Did we lose or gain electrons to form the ion?

(b) What would be the new electron configuration for the Bromine ion?

(c) Would the resulting ion be larger or smaller than the original atom?

(d) With which noble gas is the Bromine ion isoelectronic?

NUCLEUS

Isotope: BROMINE - 71 (NO CHARGE) Nuclear Symbol:

Atom or Ion? _____________________________________ ______________________

# of Protons: _____________ # of Neutrons: _____________ # of Electrons: _____________

Electron Configuration: _____________________________________________________________________

Orbital Diagram:

Circle the Highest-Energy Electrons.

Box the Valence Electrons.

Number of Valence Electrons:

_____

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Example 5-21. Answer the following regarding the Zinc isotope above.

(a) What is the most likely ion for Zinc? Did we lose or gain electrons to form the ion?

(b) What would be the new electron configuration for the Zinc ion?

(c) Would the resulting ion be larger or smaller than the original atom?

(d) Is there any noble gas that the Zinc ion is isoelectronic with?

NUCLEUS

Isotope: ZINC - 64 (NO CHARGE) Nuclear Symbol:

Atom or Ion? _____________________________________ ______________________

# of Protons: _____________ # of Neutrons: _____________ # of Electrons: _____________

Electron Configuration: _____________________________________________________________________

Orbital Diagram:

Circle the Highest-Energy Electrons.

Box the Valence Electrons.

Number of Valence Electrons:

_____

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Example 5-22. Let’s see how to include the “F” Block!!! Write the electron configuration for the following.

a) Tungsten ( #74)

b) Gold (#79)

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. 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. Note: You cannot use these if a question says to write the “complete” electron configuration!

For iron (Fe): 1s22s22p63s23p64s23d6 BECOMES: [Ar] 4s23d6

Example 5-23. Write the abbreviated electron configurations for the following elements.

a) Antimony (#51) b) Palladium (#46)

c) Mercury (#80)

IDENTIFYING ELEMENTS USING ELECTRON CONFIGURATIONS

Using either the total count of electrons (by totaling up the exponents on a complete electron configuration), or by looking at the last component of the electron configuration, we can determine what element we have by using the electron configuration.

Example 5-24. Identify the neutral elements that have the following electron configurations. 1s22s22p63s23p3

1s22s22p63s23p64s23d104p65s24d7

[Xe] 6s24f145d106p3

[Rn]7s25f146d5

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VI. PERIODIC TRENDS – Use these notes to help with your Inquiry!

GROUPS (aka Families) run VERTICALLY on the periodic table. Trends (such as reactivity and some physical properties) tend to be consistent within a group.

PERIODS run HORIZONTALLY on the periodic table. As you work through the periods on the periodic table, trends tend to repeat themselves from period to period (such as number of valence electrons).

However, there are some specific trends that can be summarized by looking at the periodic table as a whole. The following contains a description of those trends. Realize with all of these trends, we are giving general descriptions of how these trends work… as with anything in chemistry; there are exceptions to the rules!

ATOMIC SIZE (RADIUS)

The radius is ½ of the distance measured from the center of one atom to the center of an adjacent identical atom. The bulk of this distance is the “electron playground”!

Atomic size decreases across a period, due to the increased number of protons having a greater pull on the valence electrons more as you move across.

Atomic size increases down a group, as you are adding more energy levels as you go down.

This would mean that Fr (Francium) would theoretically have the largest atomic size. He (Helium) would have the smallest atomic size.

EX: Which would have a larger atomic size – Ca, Br, or Ba?

IONIZATION ENERGY

This refers to the amount of energy required to remove one outer-most (or VALENCE) electron. Since 8 valence electrons indicate a “full” outer shell (s & p sublevels), the closer the number of valence

electrons is to 8, the more difficult it will be to pull off an electron - meaning it will take more ionization energy.

Ionization energy increases across a period, because the atom is more tightly held together as the valence sublevels become full… meaning it will take more energy to pull an electron off.

Ionization energy decreases down a group, because if you have more energy levels, the outer-most electrons are further away, meaning they aren’t held on as tight… meaning it takes less energy to pull them off.

This would mean that He (Helium) would theoretically have the highest ionization energy. (Yes, he only has two valence electrons, but that’s all he can handle!) Fr (Francium) would have the lowest ionization energy.

EX: Which would have a higher ionization energy – Ca, Br, or Ba?

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ELECTRONEGATIVITY

This refers to the attraction an atom has to gain a nearby electron. The closer an atom is to 8 valence electrons (in the s & p sublevels), the higher the electronegatvitiy (because the nucleus has a stronger “+” charge!) However, would an atom that already has 8 valence electrons want to gain another?

Electronegativity increases across a period, because as you move across, the atom has a greater pull to gain another electron (to achieve the 8 valence electrons).

Electronegativity decreases down a group, because as the valence electrons get further from the nucleus, it has less of a pull to gain more electrons.

Noble (Inert) gases have no electronegativity, because they already have 8 valence electrons (filled s & p sublevels), so they have no pull to gain any more! This is also part of the reason they are considered “INERT” – meaning unreactive.

This would mean that F (Fluorine) would theoretically have the highest electronegativity. Fr (Francium) would have the lowest electronegativity.

EX: Which would have a higher electronegativity – Ca, Br, or Ba?

USE THE PERIODIC TABLE BELOW TO MAKE SUMMARIES ABOUT PERIODIC TRENDS!!!