bohr model of the atom bohr’s atomic model of hydrogen bohr - electrons exist in energy levels...

Bohr’s Atomic Model of Hydrogen

Bohr - electrons exist in energy levels AND defined orbits around the nucleus.

Each orbit corresponds to a different energy level.

The further out the orbit, the higher the energy level

De Broglie

Heisenburg

Modeled electrons as waves

Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron

Electrons exist in orbital’s of probability

Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

Schrödinger Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom

quantum mechanical model of the atom – current model of the atom treating electrons as waves.

Quantum Numbers

Wave Equation generates 4 variable solutions n - size l – shape: azimuthal

quantum

m – orientation s – spin

Address of an electron

n – Primary Quantum Number

Describes the size and energy of the orbital

n is any positive #

n = 1,2,3,4,….

Found on the periodic table

Like the “state” you live in

l – Orbital Quantum Number

Sub-level of energy

Describes the shape of the orbital

l = 0,1,2,3,4,….(n-1)

“City” you live in

n = 3

l = 0,1,2n = 2

l = 0,1n =

1

l = 0

l – Orbital Quantum Number

# level = # sublevels 1st level – 1 sublevel 2nd level – 2 sublevels 4th level = 4 sublevels

Energy Sublevels

• Labeled s, p, d, or f– Based on shape of the atom’s orbitals

– Each sublevel can only contain at most 2 e-

m – Magnetic Quantum Number

Describes the orientation of the orbital in space

Also denotes how many orbital's are in each sublevel

For each sublevel there are 2l +1 orbital's

m = 0, ±1, ±2, ±3, ±l

“Street” you live on

Look at Orbital's as Quantum Numbers

l = 0 m = 0

Can only be one s orbital

l = 1 m = -1, 0, +1

For each p sublevel there are 3 possible orientations, so

three 3 orbital's

Assigning the NumbersAssigning the Numbers The three quantum numbers (n, l, and m) are integers. The principal quantum number (n) cannot be zero. n must be 1, 2, 3, etc. The angular quantum number (l ) can be any integer between 0 and n - 1. For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1, +2.

Energy Energy LevelLevel

Possible Possible sub-sub-

levelslevels

Number of Number of Sub-levelsSub-levels

nn

No. of No. of OrbitalsOrbitals

nn22

No. of No. of ElectronElectron

ss

2n2n22

44 s, p, d, fs, p, d, f 44 1616 3232

33 s, p, ds, p, d 33 99 1818

22 s, ps, p 22 44 88

11 ss 11 11 22

Hog Hilton Time

Read the scenarioComplete the questionsCompleted packet due tomorrowHW: Finish Packet

Energy Level Diagrams

Aufbau Principle

• Electrons occupy the lowest energy level orbital available.

Aufbau Principal Lowest energy orbital

available fills first

“Lazy Tenant Rule”

Pauli Exclusion PrinciplePauli Exclusion Principle

No two electrons in an atom can have the same four quantum numbers.

Wolfgang Pauli

Every house has a different address

Pauli Exclusion Principle

No two electrons have the same quantum #’s

Maximum electrons in any orbital is two

()

Hund’s Rule

When filling degenerate orbital's, electrons will fill an empty orbital before pairing up with another electron. Empty room rule

RIGHT WRONG

Outermost sub-shell being filled with electrons

The order of sublevel filling is arranged according to increasing energy level. Electrons first fill the 1s sublevel followed by the 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p and 6s

Increasing Energy

1s

2p

6s

4s

5s

3s

2s

4d

3d

5p

4p

3p

Think about piggy palace….

p ______ ______ ______

3 s ______

p ______ ______ ______

2 s ______

1 s ______

An energy diagram for Neon

Incre

asin

g E

nerg

y

Electron Spin

1s2

2s2 2px22py

22pz2

2p61s2

2s2

Electron Configuration Notation

Orbital Notation shows each orbital O (atomic number 8)

____ ____ ____ ____ ____ ____

1s 2s 2px 2py

2pz 3s

1s22s22p4 electron configuration!

Write the orbital notation for S S (atomic number 16)

____ ____ ____ ____ ____ ____ ____ ____ ____

1s 2s 2p 3s 3p

1s22s22p63s23p4

How many unpaired electrons does sulfur have? 2 unpaired electrons!2 unpaired electrons!

Electron Configuration

Shorthand way of writing electron configuration of atoms

10Ne: 1s2 2s2 2p6

Elemental Symbol and atomic number

Principal energy level

Energy sublevel

Number of electrons

Shorthand Configuration

S 16e-

Valence Electrons

Core Electrons

S 16e- [Ne] 3s2 3p4

1s22s22p63s23p4

Longhand Configuration

[Ar]4s23d104p2

Example - Germanium

X X X X X X X X X X X X X

Let’s Practice P (atomic number 15)

1s22s22p63s23p3

Ca (atomic number 20) 1s22s22p63s23p64s2

As (atomic number 33) 1s22s22p63s23p64s23d104p3

W (atomic number 74) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4

Noble Gas Configuration

[Ne] 3s23p3

[Ar] 4s2

[Ar] 4s23d104p3

[Xe] 6s24f145d4

Your Turn N (atomic number 7)

1s22s22p3

Na (atomic number 11)

1s22s22p63s1

Sb (atomic number 51)

1s22s22p63s23p64s23d104p65s24d105p3

Cr (atomic number 24)

1s22s22p63s23p64s23d4

Noble GasConfiguration

[He] 2s22p3

[Ne] 3s1

[Kr]5s24d105p3

[Ar] 4s23d4

End of information for the test on Thursday 1/14

Valence Electrons

As (atomic number 33) 1s22s22p63s23p64s23d104p3

The electrons in the outermost energy level.

s and p electrons in last shell

5 valence electrons

Full energy levelFull sublevelHalf full sublevel

1

2

3

4 5

6

7

Copper Expect: [Ar] 4s2 3d9

Actual: [Ar] 4s1 3d10

Silver Expect: [Kr] 5s2 4d9

Actual: [Kr] 5s1 4d10

Chromium Expect: [Ar] 4s2 3d4

Actual: [Ar] 4s1 3d5

Molybdenum Expect: [Kr] 5s2 4d4

Actual: [Kr] 5s1 4d5

Exceptions are explained, but not

predicted!

Atoms are more stable with half full

sublevel

Atoms create stability by losing, gaining or sharing electrons to obtain a full octet

Isoelectronic with noble gases

1

2

3

4 5

6

7

+1 +

2-3 -2 -1

0

+3

+4

Atoms take electron configuration of the closest noble gas

Na (atomic number 11) 1s22s22p63s1

1s22s22p6 = [Ne]1

2

3

4 5

6

7

Na

1 Valence electronMetal = Loses

Ne

P-3 (atomic number 15)

1s22s22p63s23p6

Ca+2 (atomic number 20)

1s22s22p63s23p6

Zn+2 (atomic number 30)

1s22s22p63s23p63d10

Last valence electrons (s and p)

Full Octet

Element Configuration notation

Orbital notation Noble gas notation

Lithium 1s22s1 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s1

Beryllium 1s22s2 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2

Boron 1s22s2p1 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p1

Carbon 1s22s2p2 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p2

Nitrogen 1s22s2p3 ____ ____ ____ ____ ____

1s 2s 2p

[He]2s2p3

Oxygen 1s22s2p4 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p4

Fluorine 1s22s2p5 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p5

Neon 1s22s2p6 ____ ____ ____ ____ ____ 1s 2s 2p

[He]2s2p6

Half of the distance between nucli in covalently bonded diatomic molecule

"covalent atomic radii"

Periodic Trends in Atomic Radius

Radius decreases across a period Increased effective nuclear charge dueto decreased shielding

Radius increases down a group Addition of principal quantum levels

Determination of Atomic Radius

Table of Atomic Radii

Increases for successive electrons taken from the same atom

Tends to increase across a period

Electrons in the same quantum level do not shield as effectively as electrons in inner levels

    Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group

Outer electrons are farther from thenucleus

Ionization Energy: the energy required to remove an electron from an atom

Affinity tends to increase across a period

Affinity tends to decrease as you go down in a period

Electrons farther from the nucleusexperience less nuclear attraction

Some irregularities due to repulsive forces in the relatively small p orbitals

Electron Affinity - the energy change associated with the addition of an electron

A measure of the ability of an atom in a chemicalcompound to attract electrons

Electronegativities tend to increase across a period

Electronegativities tend to decrease down a group or remain the same

Electronegativity

Cations Positively charged ions

Smaller than the corresponding atomAnions

Negatively charged ions Larger than the corresponding atom

Ionic Radii

Table of Ion Sizes

Summary of Periodic Trends