Electrons in Atoms

Quantum Mechanics

Periodic Trends

Chemical Bonding

So why does potassium

explode in water?

Dalton’s Thompson’s Rutherford’s

Bohr’s carbon Quantum Model

of helium shows

2 electrons in the

1s orbital or “1s2”

Bohr proposed electrons orbit in paths of fixed energy called energy levels.

12.1 Development of Atomic Models

Size of Atom and Subatomic particles

Quantum Theory Symphony

Quantum Mechanical Model

The quantum model of the atom

is based on the solution to the

Schrödenger equation.

One way to visualize the model’s

electron levels is to imagine a

ladder where the higher rungs or

levels are closer together.

7 rungs = 7 energy levels

Principle quantum numbers (n)

correspond to the energy levels

Atomic Sublevels or S and P orbitals Probability cloud models (left)

show where it is most likely to

find electrons. For each

principal quantum number (n)

there are the same number of

sub levels. (n=2, 2 sublevels)

Helium n=1, had 1 suborbital

the “s” orbital. Level 2, n=2,

has 2 sub levels, “s” and “p.”

(The px,py and pz orbitals are

found on energy levels 2-7)

Each energy level, like level 2

shown here in green. has a

spherical “s” orbital.

1. How many sublevels are there

for each principle quantum

number?

2. Which energy levels contain px,

py and pz orbitals?

3. How are “s” orbitals different

than “p” orbitals?

Ask Your Neighbor:

Electron Configuration of Hydrogen

1s

2s 2px 2py 2pz

3p

3s

4s 3 d

H 1s1

How are all group IA

elements similar in

electron configuration?

Light will bend and reflect at the

interfaces between different

materials.

Prism lenses bend light. White light is a

blend of all wavelengths of visible light.

Each color has a different wavelength

(λ) lambda = wavelength

.4µm

.5µm

.6µm

.7µm

1000 µm = 1 mm

The shortest wavelength also has the

highest energy, hence UV light can harm

us if the wavelength is too short!

http://science.hq.nasa.gov/kids/imagers/ems/visible.html

What are emission spectra?

How is each element’s emission spectra

unique?

Evidence of Energy Levels and Suborbitals

http://phet.colorado.edu/new/simulations/sims.php?si

m=Neon_Lights_and_Other_Discharge_Lamps

Discrete lines = quanta of energy Why does each element have its own

“signature” emission spectrum? Tell neighbor.

A. Each element has a different number of protons and electrons.

B. Each element has unique nuclear attraction for electrons in shells.

C. Each atom’s first energy level is a unique distance from nucleus.

D. Distance between outer energy levels in atom is unique to each element.

E. Electrons emit photons whose frequency is proportional to energy lost.

Predictions Based on

Models of the Atom

• Click here

http://baestudent.s-cool.co.uk/animations_interactions.asp

http://intro.chem.okstate.edu/WorkshopFolder/Electronconfnew.html

Electron Configuration Rules

• Aufbau principle- Electrons enter orbitals of

lowest energy first.

• For order of orbitals from lowest to highest

learn the ZIG ZAG RULE.

Zig-Zag Rule

7s 7p

6s 6p 6d 6f

5s 5p 5d 5f

4s 4p 4d 4f

3s 3p 3d

2s 2p

1s

#e- Atomic #

8 92 118 at 7p filled

18 60 zig 6

32 40 zig 5

32 22 zig 4

18 12 zig 3

8 4 zig 2

2 2 zig 1

Electron Configuration of Helium

1ss

2s 2px 2py 2pz

3p

3s

4s 3 d

He 1s2 (filled)

Hund’s Rule

In a single energy level, electrons must occupy only one orbital until each orbital has an electron.

Since electrons have the same charge, they have strong repulsion forces, pushing them into different suborbitals of the same energy.

Pauli exclusion principle

Orbits fill with electrons of opposite spin. (½+

and ½-)

Electrons (e-) have strong, negative repulsion forces, but they are less repulsive if they have opposite spin.

1ss

2s 2px 2py 2pz

3p

3s

4s 3 d

N 1s2 2s2 2p3

Electron Configuration of Nitrogen

Predict the electron

configuration of Ne.

1ss

2s 2px 2py 2pz

3p

3s

4s 3 d

Ne 1s2 2s2 2p6

Electron Configuration of Neon

Filled 2nd energy level

(8 electrons = octet)

Predict the electron

configuration of Ar.

1ss

2s 2px 2py 2pz

3p

3s

4s 3 d

Ar 1s2 2s2 2p6 3s2 3p6

Electron Configuration of Argon

3rd energy level

(8 electrons = octet)

Filled 2nd energy level

(8 electrons = octet)

http://intro.chem.okstate.edu/WorkshopFolder/Electronconf

new.html

Observe this Shockwave Electron Configuration which introduces us to the mystery of some periodic trends.

1. What happens to the size of the atom as

the energy levels are filled?

2. When do the energy levels change the

most?

Photoelectric Effect

• Click here for photoelectric effect

simulation

High frequency light frees electrons

from reactive metals

• Patterns in the physical and

chemical properties of the

elements are called trends.

•Periodic means cycle or

repeating pattern.

Periodic Trends

Periodic Trends in Atomic radius

Group trends- Radius increases as electrons (must fill new energy levels) and are added to atom.

Atomic mass, and number increase in this direction also.

r

What happened to energy levels as p+

and e- increased across a row?

How does this affect atomic radius

across a row?

Period trends- Radius decreases as electrons fill

across same energy level. Filled inner levels

shield outermost electrons from the nucleus (So in any period, between the nucleus and outer

electrons, there is the same number of electrons.)

This trend is opposite for atomic mass &

number.

Periodic Trends in Atomic radius

Where are the metals vs. the nonmetals?

Metals have low electronegativity

Nonmetals have high

electronegativity.

Electronegativity Metal vs. Nonmetal

Where are the smallest atoms in a period?

Big atoms have lower electronegativity

Electronegativity vs. atomic size

Electronegativity explained

Valence electrons of small atoms that are closer to the nucleus than larger atoms, tend to be held to the nucleus with stronger forces of attraction. Usually the farther they are away, the weaker the forces of attraction.

e-

e-

High or strong attraction to valence electrons in a

bond = High Electronegativity.

• This is the amount of energy it takes to

remove the first or outer most electron.

• Look on your periodic table at first

ionization potential in V, or on page 362-3

in textbook.

• How easy is it to remove electron from the

Group I & II metals? From the halogens?

First Ionization Energy

First Ionization Energy

How might the first ionization energy

compare to the electronegativity across

the first period?

Th

ey a

re very

simila

r!

Trends Important to Bonding Ionization energy was used to help determine

electronegativity.

Electronegativity is a scale in the units of

Paulings, developed or calculated to show the

degree one element tends to have the bonding

electron(s) in a pair of oppositely charged ions.

http://jcrystal.com/steffenweber/JAVA/jpt/jpt.html

NaCl

The electronegativity of nonmetal Cl- is 3.12

And the electronegativity of metal K+ is .82

The difference between the two is 2.30 Pauling units

We determine the percent ionic character of the bond to

be 74%. By definition this is considered ionic bonding

since it is more than a 2.0 difference.

For example

0.82 3.12

To determine bond type: Calculate

the difference in Electronegativity

between these Element Pairs

Na and F

K and F

Li and F C and H

N and H

S and H

3.0

3.05

3.16

0.35

0.94

0.48

C and F 1.43

Difference < 2.0 = covalent Difference > 2.0 = ionic

(Formula Units of

bonded ions) (Molecules of bonded atoms)

1. Ca +2 4. Br-1

2. Al +3 5. S-2

3. K 6. N

1. Ar 1s2 2s2 2p6 3s2 3p6

2. Ne 1s2 2s2 2p6

3. 1s2 2s2 2p6 3s2 3p64s1

4. Kr [Ar] 4s2 3d10 4p6

5. Ar 1s2 2s2 2p6 3s2 3p6

6. 1s2 2s2 2p3

Write electron configurations of the following:

Use electron dot diagrams to determine

chemical formulas of the ionic compounds

formed when the following elements combine.

Example:

• K and I

• Ca and S

-2 +2 +1 -1