3-atomic structure overview characteristics of atoms interaction b/tw matter and light...

3-Atomic Structure

Overview

• Characteristics of Atoms

• Interaction b/tw matter and light– Photoelectric Effect

• Absorption and Emission Spectra

• Electron behavior

• Quantum numbers

Atomic Structure

• Atomic orbitals– Orbital energies– Electron configuration and the periodic

table

• Periodic table– Periodic properties– Energy

Characteristics of Atoms

• Atoms possess mass• Atoms contain positive nuclei• Atoms contain electrons• Atoms occupy volume• Atoms have various properties• Atoms attract one another• Atoms can combine with one another to form

molecules

Atomic Structure

• Atomic structure studied through atomic interaction with light

• Light: electromagnetic radiation– carries energy through space– moves at 3.00 x 108 m/s in vacuum– wavelike characteristics

Electromagnetic Spectrum

Visible Spectrum

Wavelength () & Frequency ()

= number of complete cycles to pass given point in 1 second

amplitude

Energy

c = x = 3.00 x 108 m/s

long wavelength low frequency

short wavelength high frequency

Low Energy

High Energy

Energy

Mathematical relationship:

E = hE = energy

h = Planck’s constant: 6.63 x 10–34 J s

= frequency in s–1

Energy

Mathematical relationship:E = hc = x

hc

E =

Energy: directly proportional to frequencyinversely proportional to wavelength

Problems 3-1, 2, & 3

1. a) Calculate the wavelength of light with a frequency = 5.77 x 1014 s–1 b) What is the energy of this light?

2. Which is higher in energy, light of wave-length of 250 nm or light of 5.4 x 10–7 m?

3. a) What is the frequency of light with an energy of 3.4 x 10–19 J?b) What is the wavelength of light with an energy of 1.4 x 10–20 J?

Photoelectric Effect

• Light on metal surface

• Electrons emitted

• Threshold frequency, o

If < o, no photoelectric effect

If > o, photoelectric effect

As , kinetic energy of electrons

Photoelectric Effect

Einstein: energy frequency

If < o electron doesn’t have enough energy to leave the atom

If > o electron does have enough energy to leave the atom

Energy is transferred from light to electron, extra is kinetic energy of electron

Ephoton = hphoton = ho + KEelectron

KEelectron = hphoton – ho

Animation

Problem 3-4

A given metal has a photoelectric threshold frequency of o = 1.3 x 1014

s1. If light of = 455 nm is used to produce the photoelectric effect, determine the kinetic energy of the electrons that are produced.

Bohr Model

Line spectra

Light through a prism continuous spectrum:

Ordinary white light

Bohr Model

Line spectra

Light from gas-discharge tube

through a prism line spectrum:

H2 discharge

tube

Line Spectra (emission)

White light

H

He

Ne

Line Spectra (absorption)

Light source

Gas-filled tube

Bohr Model

For hydrogen:

22 n

1

2

1 C C = 3.29 x 1015 s–1

Niels Bohr: Electron energy in the atom is quantized.

2n n

1RE H n = 1, 2, 3,….

RH = 2.18 x 10–18 J

Bohr ModelEatom = Eelectron = h

E = Ef – Ei = h

2

18

2Hn n

10 x 18.2

n

1RE

22

fi n

1

n

1

h

R

h

E H

Line spectrum

Photoelectric effect: n

Minus sign: free electron has zero energy

Bohr Energy Levels

Electrons

• All electrons have same charge and mass

• Electrons have properties of waves and particles (De Broglie)

mu

hparticle

Heisenberg Uncertainty Principle

Cannot simultaneously know the position and momentum of electron

x = h

Recognition that classical mechanics don’t work at atomic level.

Schrödinger Equation

Erwin Schrödinger 1926

Wave functions with discrete energies

Less empirical, more theoretical

n En

n wave functions or orbitals

n2

probability density functions

Quantum Numbers

Each orbital defined by 3 quantum numbers

Quantum number: number that labels state of electron and specifies the value of a property

Quantum Numbers

Principal quantum number, n (shell)

Specifies energy of electron (analogous to Bohr’s n)

Average distance from nucleus

n = 1, 2, 3, 4…..

Quantum Numbers

Azimuthal quantum number, (subshell) = 0, 1, 2… n–1

n = 1, = 0

n = 2, = 0 or 1

n = 3, = 0, 1, or 2

Etc. 0 1 2 3 4

s p d f g

Quantum Numbers

Magnetic quantum number, m

Describes the orientation of orbital in space

m = –….+

If = 2, m = –2, –1, 0, +1, +2

Problem 3-5

Fill in the quantum numbers in the table below.

n m

3 0 0 3s

2 –2, –1, 0, 1, 2

0

2p

Schrödinger Equation

Wave equations: Each electron has & E associated w/ it

Probability Density Functions: 2

-graphical depiction of high probability of finding electron

Probability Density Functions

energy

2 probability density function

s, p, d, f, g

1s 2s

3s

Node: area of 0 electron density

Link to Ron Rinehart’s page

Probability Density Functions

2p

Node: area of 0 electron density

3p

nodes

Link to Ron Rinehart’s page

Electrons and Orbitals

Pauli Exclusion Principle: no two electrons in the same atom may have the same quantum numbers

Electron spin quantum number ms = ½

Electrons are spin paired within a given orbital

Electrons and Orbitals

n = 1

= 0, m = 0, ms = ½

2 electrons possible:

1,0,0,+½ and 1,0,0,–½

2 electrons per orbital

1s1 H

1s2 He

Electrons and Orbitals

n = 2

= 0, m = 0, ms = ½2,0,0, ½2 electrons possible

n = 2

= 1, m = –1,0,+1, ms = ½2,1,–1, ½ 2,1,0, ½ 2,1,+1, ½6 electrons possible

Electron Configurations

n = 1

1s 2 electrons possible

H 1e– 1s1

He 2e– 1s2

Electron Configurations

n = 2

2s 2 electrons possible

Li 3e– 1s2 2s1

1s

2s

Be 4e– 1s2 2s2

1s

2s

Electron Configurations

n = 2

2p = 1, m = –1, 0, +1

3 x 2p orbitals (px, py, pz): 6 electrons possible

B 5e– 1s2 2s2 2p1

1s

2s

2p

Electron Configurations

n = 2

2p = 1, m = –1, 0, +1

3 x 2p orbitals (px, py, pz): 6 electrons possible

B 5e– 1s2 2s2 2p1

1s

2s

2p

Electron Configurations

n = 2

2p = 1, m = –1, 0, +1

C 6e– 1s2 2s2 2p2

1s

2s

2p

Hund’s Rule: for degenerate orbitals, the lowest energy is attained when electrons w/ same spin is maximized

Problem 3-6

Write electron configurations and depict the electrons for N, O, F, and Ne.

Electron Configurations

n = 3

3s, 3p, 3d

Na 11e– 1s2 2s2 2p63s1

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

Mg 12e– 1s2 2s2 2p63s2

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

Al 13e– 1s2 2s2 2p63s23p1

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

Si 14e– 1s2 2s2 2p63s23p2

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

P 15e– 1s2 2s2 2p63s23p3

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

S 16e– 1s2 2s2 2p63s23p4

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

Cl 17e– 1s2 2s2 2p63s23p5

1s

2s

2p

3s

3p

Electron Configurations

n = 3

3s, 3p, 3d

Ar 18e– 1s2 2s2 2p63s23p6

1s

2s

2p

3s

3p

Electron Configurations

3d vs. 4s

Filling order 1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f 5g

6s 6p 6d

7s 7p

Electron Configurations

1s

2s2p

3s

3p

4s

4p

3dK

Electron Configurations

1s

2s2p

3s

3p

4s

4p

3dCa

Electron Configurations

1s

2s2p

3s

3p

4s

4p

3dSc

Electron Configurations

1s

2s2p

3s

3p

4s

4p

3dTi

Link to OSU site

Problem 3-7

Write the electron configurations for the transition metals V – Zn. Fill in the corresponding boxes to denote the electronic spin.