Chem 1A03 - Atomic Structure and Theory (Term 1)

Electromagnetic spectrum

• Light is is a form of energy; referred to as electromagnetic radiation

• C = 𝝀𝝂

  • speed of light is constant (c)

1nm conversion rate

= 10^-9 m

Wavelength

distance from peak to peak (m)

Frequency

how many peaks in a given time (s^-1 or Hz)

How is wavelength and frequency related

• High frequency = shorter wavelength

• Low frequency = longer wavelenght

Blackbody radiation

Blackbody Radiation – all objects emit electromagnetic radiation due to thermal energy (except

at 0 K). As a blackbody is heated, it will eventually glow, producing visible light, a phenomenon

known as incandescence.

Stefan-Boltzman law

States that the total intensity emitted radiation of a blackbody is σT4.

Wilhelm Wien

Wilhelm Wien found the wavelength emitted with maximum intensity, λmax is proportional

to 1/T (Wien’s Law) The higher the temperature, the shorter the λmax.

Ultraviolet Catastrophe

• Using classical wave theory, Rayleigh and Jeans derived that I(λ) would be proportional to 1/λ4.

• So, hot bodies should emit intense UV and X-Ray radiation.

• Even the human body would produce light and glow in the dark.

• Classical wave theory could not predict observed spectra – the ultraviolet catastrophe.

Quantum Theory

Energy is Quantized (discontinuous)

Photon

a discrete packet (quantum) of energy

Voltmeter

measures how fast the e- are going (KE)

Ammeter

measures how many e- are ejected and travelling (current)

Red light - 700nm

red light does not eject electrons and no current flow

e red light < e threshold

Ammeter and Voltmeter show no reading

increasing the intnesity of the red light does not chnage the readings

Green light - 500nm

green light ejects photons and current flows

E green light > e threshold

• Ammeter: reading increases

• Voltmeter: Reading increases

Purple light

• ejects electrons and currents flows

• e purple light>e threshold

• ammeter: reading (no change compared to green light)

• Voltmeter: reading increases compared to green

Doubling the intensity of the purple light

Doubling the intensity of the purple light means you double the number of ejected electrons and current

E purple light > E threshold

Ammeter: Reading (2x the previous reading

voltmeter: Reading (no chnage from less intense light)

The wavelength of light needed to eject electrons from a metal is 85.8 nm.

Are electrons ejected if light of 90.0 nm shines on the metal sample?

A. Yes

B. No

Now. The energy has to be the same or higher to generate a photoelectron

Longer wavelength = less energy

Each packet of light interacts with 1 electron, if you have a sufficent packet of light to overcome the attraction of the elcton to the nucleus of the atom and bulk of the material, then it gets ejected

To ionize you have to meet that energy threshold

The development of an atom: Daltons theory

single indivisible entities unique to each element

cannot be altered destroyed

combinations form more complex atoms

however; there were negatively charged particles called electrons that were not accounted for

Plum Pudding Model

negatively charged particles , corpuscles, electrons, were embedded in a sea of positive charges

• However, when we fire alpha particles at gold foil, they get depleted.

The nuclear model

highly concentrated postive charges in a very small volume contained most of the atoms mass and electrons orbiting around it.

Problem with the nuclear model

any charged particle that is moving in a circular obrbit continuously will eventually loe energy and will spiral into the nucleas

Bohrs idea: To combine the model of an atom with Planck’s ideas of quantized energies

Electronic transition in hydrogen: Absorption

Electronic transition in hydrogen: emission

Quantifying the electronic transitions in hydrogen

DELTA E> 0 : Absorption

DELTA E < 0: Emission

Thompsons electron diffraction experiemnt

diffraction occurs when radiation and object spacing are similar. (e.g. X-

rays,  = 100 pm, array of Al atoms 200 pm spacing)

Wave-particle duality

• waves, like light, display particle like properties

• Partciles, like lectrons, display wave like properties

  • Small particles like electrons have long wavelength

Heseinberg Uncertainty principle

there is a uncertainty of the position of an electron in an atom

• one can olny speak of the probailty of finding the electron in a volume within the atom.

Principal quantum number

• n

• Describes orbital energy and distance from the nucleus

• the larger the n vale, the further it is from the nucleus

• E in the orbital is at a higher energy level

• n describes a principal shell

Obrital angular quantum number, l

• describes orbital shape

• l - 0, 1, 2, 3, n-1

• S=0

• p=1

• d=2

• f=3

• Number of l values is the number of subshells in a given n shell

• cannot be negative

Magentic quantum number, ml

• describes orbital orientation

• ml = -l….+l

• ml has (2l + 1) values

• 
  

Electron spin quantum numeber, ms

• Electrons have spin, generating a magnetic field

• values are + or - ½

• pair of electron with opposite spins has no magnetic field

Diamagnetic vs paramagnetic

• Atoms or ions with all spins paired are diamagentic

  • most atoms are like this, the effect is so weak that it goes unoticed

• 1or more unpaired electrons is paramagnetic

  • stronger than dimagnetic so it ovverides it.

Orbital shapes; s and f

Orbital shapes: d

Aufbau Principles

Occupy orbitals so that the atoms energy is minimized; occupy orbitals of lower energy

Paule-exclusion principle

ever e- has a unique set of quantum numbers

Hunds rule

orbitals of the same energy level are occupied singly before they are paired