Spectroscopy and Electron Configurations - PowerPoint Presentation

Slide1

Spectroscopy and Electron ConfigurationsSlide2

Light is an electromagnetic wave*.Slide3

Wave Characteristics

Frequency (

ν

)

: number of waves that pass a point in a given period of time

Total

energy

is proportional to amplitude and frequency of waves.

Because

speed of light (c)

is a constant (3 x 10

8

m/s), wavelength and frequency of electromagnetic waves are inversely proportional:

E = h

ν

; c =

νλSlide4

White light is composed of all colors which can be separated into a rainbow, or a spectrum, by passing the light through a prism

.

Each color light has a different wavelength, and, therefore, frequency.

ColorSlide5

5

Amplitude & WavelengthSlide6

Are there other “Colors”?Slide7

The Electromagnetic Spectrum

Visible light comprises only a small fraction of all the wavelengths of light – called the

electromagnetic spectrum.

Short wavelength (high frequency) light has high energy.

- Gamma ray light has the highest energy.

Long wavelength (low frequency) light has low energy.

-

Radiowave light has the lowest energy.Slide8

Electromagnetic SpectrumSlide9

Interactions of light and matter

Emission Transmission

Absorption Reflection

or ScatteringSlide10

What types of light spectra can

we observe?

A hot opaque body produces a

continuous spectrum

, a

complete rainbow of colors without any specific spectral lines

.Slide11

A hot, transparent gas produces an

emission line spectrum

a series of bright spectral lines against a dark background.Slide12

A cool, transparent gas in front of a source of a continuous

spectrum produces an

absorption line spectrum

- a series

of dark spectral lines among the colors of the continuous

spectrum.Slide13

Each chemical element produces its own unique set of spectral lines

.

Oxygen spectrum

Neon spectrumSlide14

Emission and absorption spectra

are inversely related.

Spectra of MercurySlide15

Identifying Elements with

Flame Tests

Na

K

Li

BaSlide16

Exciting Gas Atoms to Emit Light

with Electrical Energy

Hg

He

HSlide17

Analyzing the Hydrogen Emission Spectrum

Rydberg

found the spectrum of hydrogen could be described with an equation that involved an inverse square of integers.Slide18

Bohr Model of Hydrogen Atom

In the Bohr model, electrons:

- have

quantized

energies

.

- have orbits a fixed distance from

the nucleus.

e

-Slide19

The Bohr ModelSlide20

Interference: When Waves InteractSlide21

DiffractionSlide22

Tro, Chemistry: A Molecular Approach

22

2-Slit InterferenceSlide23

If electrons behave like particles, there should only be two bright spots on the target.

Electron DiffractionSlide24

Electron Diffraction

However, electrons actually present an interference pattern, demonstrating they behave like waves.Slide25

The Bohr Model

Integer number of de Broglie wavelengths must fit in the circumference of orbit.Slide26

26

Electron Transitions

To transition to a higher energy state, the electron must

absorb

energy equal to the energy difference between the final and initial states.

Electrons in high energy states are unstable. They will transition to lower energy states and

emit

light.Slide27

Principal Energy Levels in Hydrogen

The wavelengths of lines in the emission spectrum of hydrogen can be predicted by calculating the difference in energy between any two states.Slide28

Bohr Model of H AtomsSlide29

29

Hydrogen Energy Transitions

For an electron in energy state

n

, there are (

n

– 1) energy states to which it can transition. Therefore, it can generate

(

n – 1) lines.Slide30

Chemical Fingerprints

Every atom, ion, and molecule has a unique spectral “fingerprint.”

We can identify the chemicals in gas by their fingerprints in the spectrum.

With additional physics, we can figure out abundances of the chemicals, and much more.Slide31

Other spectroscopy

Many spectroscopic techniques rely on these electronic transitions used with different sources of light.

Energy can also be absorbed and emitted in other “modes” including vibration and rotation, leading to other types of spectra.

OHSlide32

Uncertainty Principle

Heisenberg stated that the product of the uncertainties in both the position and speed of a particle was inversely proportional to its mass.

x

= position,

Δ

x

= uncertainty in position

v = velocity,

Δv = uncertainty in velocitym = massThe more accurately you know the position of a small particle, like an electron, the less you know about its speed.

and vice-versaSlide33

Wave Function,

y

Calculations show that the size, shape and orientation in space of an orbital are determined by three integer terms in the wave function.

added to quantize the energy of the electron

These integers are called

quantum numbers.

principal quantum number,

n

angular momentum quantum number, l

magnetic quantum number,

m

l

spin quantum number,

m

sSlide34

Principal Quantum Number,

n

characterizes the energy of the electron in a particular orbital

corresponds to Bohr’s energy level

n

can be any integer.

The larger the value of

n

, the more energy the orbital has.Energies are defined as being negative.An electron would have E = 0 when it just escapes the atom.The larger the value of n, the larger the orbital.As n gets larger, the amount of energy between orbitals gets smaller. for an electron in H

E

n

= -2.18 x 10

-18

J

1

n

2Slide35

l

= 0, the

s

orbital

Each principal energy state has 1

s

orbital.

lowest energy orbital in a principal energy statesphericalnumber of nodes = (n – 1)Slide36

p

orbitalsSlide37

d

orbitalsSlide38

f

orbitals