Atomic Structure Part II

Electrons in Atoms

Radiant energy travels in the form of waves that have both electrical and magnetic properties.

These electromagnetic waves can travel through empty space, as you know from the fact that radiant energy from the sun travels to Earth every day.

Electrons in Atoms

To understand the relationship between the nature of the atomic structure (arrangement of electrons in the atom) and how elements emit light when heated, it is necessary to first understand the nature of light.

Electrons in Atoms I. Light and Quantized Energy -

Properties of Waves 1. Definition of Radiation:

Energy that exhibits wave-like (or oscillating) behavior as it travels through space

Electrons in Atoms I. Properties of Waves 1. Definition of Radiation:

Electrons in Atoms 2. Wavelength (λ) distance from peak

to peak, length of one complete wave Abbreviation Lamda (λ)

Electrons in Atoms 2. Wavelength (λ) distance from peak to

peak, length of one complete wave 3. Frequency (ν)

a. number of peaks that pass at a given point each sec b. can be called cycles per second (peak/sec) c. cps now called 1 Hertz (Hz)

Figure 06.31-09UNEOC

Electrons in Atoms - Cont. 4. Velocity (C = speed of light) a. distance a given peak moves in

a unit of time b. velocity (m/s) = frequency x

wavelength

c = ν x λ

II. Behavior of Light A.  Newton (1600) thought light consisted of particles

(beam of light is a stream of particles)

B. Maxwell (1864) thought light was a wave

phenomenon. • Calculated the velocity of the propagation of an electromagnetic wave and found it was the same for light

II. Behavior of Light 1. some say light is like waves, some say

its like particles 2. modern theory says that it behaves as

both "wave/particle duality"

II. Behavior of Light 3. Max Planck (early 1900's) said: a. light is made up of bundles of energy called

photons (or quanta) b. the energy of each photon is proportional to

the frequency of the light (Quantum Theory) A quantum is… the minimum amount of energy that can be gained or lost by an atom (electron in an atom.)

•  example: CONTINUOUS SPECTRUM

*** when white light is passed through a prism, it is separated into a band of colors from red è violet. It's called a continuous spectrum

c. the work of Planck & Einstein led to

E=energy,

ν = frequency,

h=planks constants (6.6262x10-34J·sec) ���J is the symbol for joule the SI unit for energy

E = h x ν

Energy of a quantum is related to… the frequency of the emitted radiation by this equation

c. the work of Planck & Einstein led to

According to Planck’s theory, for a given frequency….. matter can emit or absorb energy only in whole-number multiples of hv, that is 1hv, 2hv, 3hv, and so on.

E = h x ν

c. The photoelectric effect

In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface. (example: solar calculator)

c. The photoelectric effect

Einstein said light can both wavelike and particle like natures. That is, while a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles or energy, called photons.A Photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

III. Bright line spectrum •  A. a spectrum that shows separate bright

lines, each with a specific wavelength •  B. bright-line spectra occur when an

element is heated and the colored light given off is viewed through a spectroscope. Each element has a unique set of lines, characteristic of that element (like a fingerprint)

Line-Emission Spectrum

ground state

excited state

ENERGY IN PHOTON OUT

Fireworks? Hmmm…

IV. Electromagnetic Spectrum •  A. visible light (like the continuous spectrum)

is only one type of radiation. All other types are not visible to the human eye.

HIGH ENERGY

LOW ENERGY

Long (λ) ���Low Freq. (υ)

Short (λ) ���High Freq. (υ)

Electromagnetic Spectrum

LOW ENERGY

HIGH ENERGY

R O Y G. B I V

red orange yellow green blue indigo violet

Electromagnetic Spectrum

B. all forms of electromagnetic radiation travels at the speed of light.

1. speed of light = 3.00 x 108 meters/sec 2. use formula:

c = ν x λ 3. each line spectrum has a particular

frequency (ν ). If know wavelength (λ), we can find ν using c as a constant.

C. The energy in a photon of light is directly proportional to the frequency of the light.

•  1. frequency, energy •  2. can find the energy of a single

quantum (photon) of radiation at any given frequency.

C. The energy in a photon of light is directly proportional to the frequency of the light.

•  3. proportionality constant that relates the two is called Planck's constant (h).

•  4. formula:

E = h x v

example: a spectral line has frequency of 3.5x10 12 hertz. What is the energy of a photon of radiation of this frequency? E = h x v (h=6.6262x10-34J/sec)

E = (3.5x1012Hz) (6.6262x10-34J • sec)

E = (2.3x10-21J)

V. Electron energy levels in Bohr's Model

A. There are certain different orbits in which an electron can travel around a nucleus.

1. each circular orbit (or shell) is at a fixed distance from the nucleus

V. Electron energy levels in Bohr's Model

2. the greater the radius of that shell, the greater the energy of the electron in that shell.

3. these electron orbits are known as energy levels

B. When electrons absorb energy firm an outside source, they jump from lower to higher energy levels. àwhen they fall back to their original levels , energy is emitted (light); the same amount as was absorbed.

B. Bohr Model

12

3456 ❚ Energy of photon

depends on the difference in energy levels❚ Bohr’s calculated

energies matched the IR, visible, and UV lines for the H atom

Close to nucleus ���low energy

higher energy

C. In energy atom in its normal state, all electrons are in the lowest energy levels available (energetically stable)

VI. Atoms and Radiation

•  A. When all of the lowest energy levels are occupied, the atom is in the ground state (unexcited).

VI. Atoms and Radiation

•  B. When electron moves to higher energy level, atom is in the excited state, and is energetically unstable.

VI. Atoms and Radiation

C. Bright line spectrum of an element represents the energy levels in its atoms.

✁problems with Bohr's Model:

✁ only explained some of the lines in the bright line spectrum

✁ really only worked for hydrogen ✁ need sublevels and electron cloud

model to account for all of the lines.

VII. The Modern Model of the Atom

A. Mechanics 1. Classical Mechanics - Newton's Laws of

Motion (Newtonian Mechanics)

Describes the behavior of visible objects traveling at ordinary velocities. Bohr’s basis for his model, but couldn’t explain why electrons would stay at on energy level or another. When looking at H-spectral lines, noticed more one (several closely spaced).

VII. The Modern Model of the Atom 2. Quantum Mechanics – (wave

mechanics) Describes the behavior of extremely small particles traveling at velocities at or near the speed of light

a. Louis de Broglie - particles could have properties of waves

Planks quanta gave wave properties, deBroglie said electron streams are like waves of light and have properties of both particles and waves (matter behaves as waves)

• b. Schrodinger - described the behavior of electrons in terms of quantized energy changes "quantum mechanics"

Describe a wave equation used to determine the probability of finding an electron in any give place or orbital

Schrodinger’s Cat

Schrodinger Cat part II

Radial Distribution Curve Orbital

c. Heisenberg - uncertainty principle

!Region of space where there is a probability of finding an electron is called an orbital

"The more precisely the POSITION is determined, the less precisely ���the MOMENTUM is known"

B. Principal Energy Levels 1. Energy Levels •  Bohr - High Energy

(outer level)

Low Energy

1 2 3

4

Principal Quantum Numbers (N) Number of electrons

2818

32

Corresponds to energy level

2. Sublevels Principal Quantum Numbers (N) Sublevel

Present

1 1s 2 2s2p 3 3s3p3d 4 4s4p4d4f

•  Orbital -

–  Region of space where an electron is probably found

•  Electron spin –  An orbital can hold 2 electrons that spin in

opposite directions.

Electrons are represented by arrows

Rules: 1. Pauli Exclusion Principle

– Each orbital can hold TWO electrons with opposite spins.

– No two electrons in an atom can have the same 4 quantum numbers.

– Each e- has a unique “address”:

2. Aufbau Principle Electrons fill the

lowest energy orbitals first.

Electrons to be added must be placed in unfilled orbitals of lowest energy for stable configuration.

3. Hund’s Rule – Within a sublevel, place one e-

per orbital before pairing them. – “Empty Bus Seat Rule”

RIGHT WRONG

Energy Level Diagram

orbital

Orbital - Place where electrons are probably found

Electrons have “up” and “down” spin

c

Shapes of electron orbitals ���

The s orbital

The p orbitals

py

px

The p orbitals

pz

The d orbitals

The d orbitals

The f orbitals

f

Click here for orbital viewer View the grand table

Orbital Shapes Video

s and p orbital shapes

C. Electron Configurations

1s2 = Helium 1s22s1 = Lithium 1s22s22p63s23p6

4s23d104p6 = Krypton

Energy Level

Sub Level # of Electrons

O

8e-

•  Orbital Diagram

•  Electron Configuration

1s2 2s2 2p4

Electron Configuration Notation

1s 2s 2p

•  Shorthand Configuration

S 16e-

Valence Electrons Core Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

Notation•  Longhand Configuration

© 1998 by Harcourt Brace & Company

s p

d (n-1)

f (n-2)

1 2 3 4 5 6 7

6 7

Periodic Patterns

Periodic Patterns

•  Period # –  energy level (subtract for d & f)

•  A/B Group # –  total # of valence e-

•  Column within sublevel block –  # of e- in sublevel

s-block 1st Period

1s1 1st column of s-block

1 2 3 4 5 6 7

Periodic Patterns•  Example - Hydrogen

1

2

3

4

5

6

7

Periodic Patterns

•  Shorthand Configuration – Core e-: Go up one row and over to the

Noble Gas. – Valence e-: On the next row, fill in the #

of e- in each sublevel.

[Ar]

1 2 3 4 5 6 7

4s2 3d10 4p2

Periodic Patterns•  Example - Germanium

•  Full energy level

1 2 3 4 5 6 7

•  Full sublevel (s, p, d, f)

•  Half-full sublevel

Stability

1 2 3 4 5 6 7

Stability•  Ion Formation

–  Atoms gain or lose electrons to become more stable.

–  Isoelectronic with the Noble Gases.

1 2 3 4 5 6 7

Stability•  Ion Formation

–  Atoms gain or lose electrons to become more stable.

–  Isoelectronic with the Noble Gases.

Feeling overwhelmed?

Try a few! Mg = Fe = Ru = Ir = Ca+2 =

Cl-1 =