Pay attention to EVERYTHING;

Write down only things that are UNDERLINED;

These notes will be available on line the day before the test as a review

• Why the electron?• The Study of Electrons• Planck’s Constant and Einstein’s Equation• Heisenberg’s Uncertainty Principle• Quantum Numbers• Aufbau Diagram• Pauli’s Exclusion Principle• Orbital Filling• Lewis e- Dot Notation

• Why all this fuss over the almost mass-less electron?• Because it is the electron (e-) that actually takes part in

chemical reactions.• It is the first thing to come in contact with other electrons

from other atoms• It is the reason that atoms and molecules react, and their

behavior dictates how quickly that happens if at all, how much energy is released or absorbed in the process, and what substances may be formed as a result

• Therefore, the importance of the electron cannot be understated!

Once scientists realized that it was indeed the e- that played such an important role, they realized that it was important to somehow categorize these electrons.

In the early 1900’s, Max Planck, a German Physicist, found that matter could absorb or emit energy in discrete amounts called “quanta”. He developed the relationship…

• Where E is the change in energy • n is an integer (1,2,3) • h is Planck’s constant , which is (6.626 x

10-34 J.s)• and = the frequency of the

electromagnetic radiation absorbed or emitted by the atom

E = n h

• Because it gives us quantitative information about the color of light emitted by a particular atom when energy is absorbed, then released, by that atom

• It helps us understand the concept of “atomic fingerprint” (that each atom has its own color pattern when photons of energy are analyzed)

• Atomic fingerprints help us identify unknown substances and help us better understand our universe (we know what chemistry exists in a star, for example, by analyzing light coming from it)

• From Zumdahl, page 283:

• The blue color in fireworks is often achieved by heating copper (I) chloride (CuCl) to about 1200oC. Then the compound emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum) that is emitted at 4.50 x 102 nm by CuCl?

• Using E = h (c = speed of light)• v = c / or 2.9979 E 8 m/s = 6.66 E 14 s-1

• 4.50 E -7 m• So E = h = (6.626 E -34 J.s) (6.66 E 14 s-1) = • 4.41 E -19 J

• That is the size of a quantum of energy for CuCl

• Einstein, studying Planck’s works, suggested that electromagnetic radiation can be viewed as a stream of “particles” called photons (“bundles of light energy”)

• So… Ephoton = h = hc / • By substitution, Einstein • developed the famous equation • in 1905, linking mass to energy

• E = mcE = mc22

So we were saying that …

it is important to know as much as possible about the electron…

• Werner Heisenberg, one of the primary developers of quantum mechanics, stated in 1927 that…

• "it is impossible to accurately determine both the momentum () and position of an e- or any other very small particle“

• = m v (m = mass; v = velocity)

• So, we now know that the more we know about an object’s position, the less we know about it’s momentum,

• And the more we know about an object’s momentum, the less we know about it’s position!

• Scientists developed a set of Quantum Numbers that increase our understanding about electrons.

• Scientists, needing to know as much about the e- as possible, developed a statistical approach to determine the probability of describing e-’s at any given moment.

• This desire to describe the electron caused a new field of study, known as “Quantum Mechanics”

• a ‘quantum’ = “very small”

• a "quantum" is a quantity or amount, but in chemistry will relate to the energy of one photon

• A photon is a "bundle of light energy"

• Quantum numbers help identify a SPECIFIC e-

• According to the "Pauli Exclusion Principle", no 2 e-'s can have the same set of four quantum numbers (like “addresses”)

• Analogy: no 2 students have the same address

• 1= state• 2= city• 3= street name• 4= house #

• describes the main energy level an e- occupies• "n" may be any positive integer... 1,2,3,4,...7+… • …or may be a letter... K,L,M,N,...Q+• an energy level may also be called a "shell"• “n” indicates whether the e- is near or far from

the nucleus• n = 1 2 3 4 5 6 7

CLOSEST TO NUCLEUS FARTHEST FROM NUCLEUS

• ...the average energy of the level• ...the size of the orbital (path)• ...the average distance from the nucleus• ...the amount of e-'s that can fit in the level

• “2n2" tells the capacity of e-'s in each level• ...so when n = 1, the capacity is_____• ... when n = 2, the capacity is_____• ... when n = 3, the capacity is_____• ... when n = 4, the capacity is_____

2

8

18

32

• In general, only e-'s from the outermost occupied energy level participate in chemical reactions.

• The outer energy level is called the "valence" level, and

• e-'s in the valence level are called "valence e-'s"

• describes "sublevels" or "subshells"• designates the shape of the region in space• there are 4 basic shapes:

sublevel s p d f (g,h,i, etc)

name sharp principal diffuse funda-mental

e- capacity 2 6 10 14 more

shape: sphere dumbbell lobes

• designates a particular orbital within a given sublevel (s,p,d,f) in which an e- resides

• an orbital has a capacity of 2 e-'s. • If an orbital is full, e- spins are

opposite• each subshell is composed of one

or more orbitals

s sublevel Has a capacity of 2 e-’s

And therefore has ___ orbital

p sublevel Has a capacity of 6 e’s

And therefore has ___ orbitals

d sublevel Has a capacity of 10 e’s

And therefore has ___ orbitals

f sublevel Has a capacity of 14 e’s

And therefore has ___ orbitals

1

3

5

7

• capacity of s sublevel = ___ electrons; • capacity of p sublevel = ___ electrons; • capacity of d sublevel = ___ electrons;• capacity of f sublevel = ___ electrons.• • How to remember? EASY! • JUST LOOK AT THE PERIODIC TABLE!

(See next slide!)

26

1014

s pdf

• e-'s behave as though each was spinning about its own axis

• "s" or specifies the direction of the spin about its own axis

• use "ms" = + 1/2 or "ms" = -1/2

• There are 3 important ways chemists identify & label electrons.

• 1. Using "electron configuration notation"

• 2. Using “orbital notation”

• 3. Using “electron dot notation”

• This is the most useful notation for electrons.

• It helps us identify the element and position it on the Periodic Table

• It helps us see how many valence e- there are

• It uses the “Aufbau Diagram” (aka: the “Diagonal Rule”

•• 1s• 2s 2p• 3s 3p 3d• 4s 4p 4d 4f• 5s 5p 5d 5f• 6s 6p 6d• 7s 7p

Start here for each new element!Each of these

represents a

“sublevel”.

How many electrons fit in each of these sublevels?

2

2

2

2

2

2

2

6

6

6

6

6

6

All s sublevels can hold 2 e-’s

All p sublevels can hold 6 e-’s

All d sublevels can hold 10 e’s

10

10

10

10 All f sublevels can hold 14 e’s

14

14

• Let’s write the e- config. nota. for Neon.

• We need to know the atomic number (10)

• We must follow the Aufbau diagram until we have accounted for all 10 electrons.

• 1s2 2s2 2p6

• Let’s write the e- config. nota. for Sc.

• We need to know the atomic number (21)

• We must follow the Aufbau diagram until we have accounted for all 21 electrons.

• 1s2 2s2 2p6 3s2 3p64s2 3d1

• Uses the diagonal rule (Aufbau Diagram) coupled with boxes (or circles) & arrows (notice that 4s gets filled before 3d = diagonal rule)

• Based on "Hund's Rule" = e-'s do not pair up until there is already

Each orbital is a place where up to 2 electrons exist.

Electrons can pair up if they have opposite spins.

The fact that their spins are opposite means we use up ( ) and down ( ) arrows to show each separate electron.

These are “orbital boxes”; their capacity is 2 e-This is a 1s sublevel and it has one orbital, so can fit 2 e-’s

This is a 3p sublevel containing 3p orbitals; capacity of 6 e’s

Step 1 – Name the element and the number of electrons it has

Step 2 – Remembering Hund’s Rule (orbitals within a sublevel don’t fill up until each orbital of the sublevel has one e- in it)

Step 3 – Use the Aufbau Diagram (ie: 4s gets filled before 3d)) Let’s Do An Orbital Filling Notation for…Neon (Ne) Z = 10

20Ca

5B

18Ar

30Zn

• shows the s & p electrons from the valence (outermost) energy level, since these are the ones most likely to enter into chemical reactions

• since s has a capacity of 2 e-'s and p has a capacity of 6 e-'s, …

X 1 The first e- goes here

2 The second e- goes here

3 The third e- goes hereThe sixth e- goes here 6

The fourth e- goes here 4

The fifth e- goes here 5

The seventh e- goes here 7

8 The eighth e- goes here

• it is important to note the number "8" as being the number of e-'s that most elements want to have in their valence shell

• we call "8" an "octet" of e-'s; atoms are stable when they have an octet of e-'s

• the dots must be drawn in a specific order, which will help you draw molecular structures later on in Chemistry

Draw the e- dot notation for 19K

1 – Write the e- config. nota. for K.

1s2 2s2 2p6 3s2 3p6 4s1

2 - What is the OUTERMOST VALENCE LEVEL? Ans.: level 4 is the highest level entered

3 – Using the order on the previous page, place a dot around the symbol for K representing each outer energy level s and p electron (which in this case is just 1)

Draw the e- dot notation for 34Se

1 – Write the e- config. nota. for Se.

1s2 2s2 2p6 3s2 3p6 4s2 3d104p4

2 - What is the OUTERMOST VALENCE LEVEL? Ans.: level 4 is the highest level entered

3 – Using the order on the previous page, place a dot around the symbol for Si representing each outer energy level s and p electron (which in this case is 6)

• the order: X

• Li N Ne Na C

• Be O F B H

• find these elements on the periodic table. Above the column on the periodic table in which you find these elements, write the number of electron dots each has.

• What relationship do you find?

Quantum numbers for the first Four Levels of Orbitals

n l Orbital Designation

mlNumber of Orbitals

1 0 1s 0 1

2 0 2s 0 1

1 2p -1,0,+1 3

3 0 3s 0 1

1 3p -1,0,+1 3

2 3d -2,-1,0,+1,+2 5

4 0 4s 0 1

1 4p -1,0,+1 3

2 4d -2,-1,0,+1,+2 5

3 4f -3,-2,-1,0,+1,+2,+3 7

• "Paramagnetic" - has an odd # of e-'s, at least one unpaired e-

• weakly attracted to magnetic field• "Ferromagnetic" - from Fe = Ferrum = Iron• STRONGLY magnetic behavior due to

interactions b/w paramagnetic atoms in the solid state

• Ferro- is about 1 million times stronger than para- (for simplicity, 1s, 2s, and 2p sublevels are not shown)

• the more unpaired e-'s in a subshell you have, the more magnetic

• Emission Spectra - produced when e-'s leave "ground state" (low)

• a) as an atom absorbs E, its e-'s move to higher E levels b) as e-'s drop back to lower E levels, they give back this E

• c) the l of the emitted radiation reflects the size of the e-'s jump

• (E is increased as l decreased)• Electrons Elevated By Absorbing E•

• de Broglie: "light can exhibit both wave & particle properties"

• l = h/mv if E = mc2 and E = hv, hv = mc2

• if v = c/l, then hc/l = mc2, and solving for l,

• we get l = h/mc

• "Bright Line Spectra" Lab• "emission spectra"• when an electric current is passed thru a gas

in a vacuum tube• at very low P's, the light that the gas emits is

dispersed by a• prism into distinct lines. This shows an

atomic "fingerprint"• heavy elements produce complex spectra;

light elements• produce the fewest lines

• - i.e.: a green line of l 4.86 E -7 m is observed in the emission spectrum of H. Calculate the E of 1 photon of green light

• We know l of light; calculate the v.• lv = c• v = c/l; • 3.00 E 8 m/sec = 6.17 E 14 s-1

• 4.86 E -7 m• E = hv; (6.63 E -34 J s)

(6.17 E 14 s-1) = 4.09 E 19 J

• - e-'s can have Q. N.'s 1,2, and 3 (n, l, ml) equal, but ms distinguishes one e- from another in the same orbital

•

•

• How e-'s and magnetism are related• Spin Q.N. (ms) can be +1/2 or -1/2• Spinning produces a tiny electromagnet• "Diamagnetic" - e-'s are paired• substances experience NO attraction for a

magnetic field - they have the same #'s of e- of each spin; spins cancel out

• 1s 2s 2px 2py 2pz 3s 3px 3py 3pz 4s 3d• 18Ar