Atomic Structure and the Periodic Table

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# Atomic Structure and the Periodic Table - PowerPoint PPT Presentation

Atomic Structure and the Periodic Table. The electronic structure of an atom determines its characteristics. Studying atoms by analyzing light emissions/ absorbtions. Spectroscopy: analysis of light emitted or absorbed from a sample Instrument used = spectrometer

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### Atomic Structure and the Periodic Table

Studying atoms by analyzing light emissions/absorbtions
• Spectroscopy: analysis of light emitted or absorbed from a sample
• Instrument used = spectrometer
• Light passes through a slit to become a narrow beam
• Beam is separated into different colors using a prism (or other device)
• Individual colors are recorded as spectral lines
• Light energy
• A wave of electric and magnetic fields
• Speed = 3.0 x 108 m/s
• Wavelength () = distance between adjacent peaks
• Unit = any length unit
• Frequency () = number of cycles per second
• Unit = hertz (Hz)
Relationship between properties of EM waves
• Wavelength x frequency = speed of light

·v = c

Calculate the frequency of light that has a wavelength of 6.0 x 107 m.

Calculate the wavelength of light that has a frequency of 3.7 x 1014 s-1

Visible Light
• Wavelengths from 700 nm (red) to 400 nm (violet)
• No other wavelengths are visible to humans
Quanta and Photons
• Quanta: discrete amounts
• Energy is quantized – restricted to discrete values
• Only quantum mechanics can explain electron behavior
• Analogy: Water flow
Another analogy for quanta
• A person walking up steps – his potential energy increases in a quantized manner
Photons
• Packets of electromagnetic energy
• Travel in waves
• Brighter light = more photons passing a point per second
• Higher energy photons have a higher frequency of radiation
• Planck constant

h = 6.63 x 10-34Js

E = hv

The energy of a photon is directly

proportional to its frequency

Deriving Planck’s constant

In a laboratory, the energy of a photon of blue light with a frequency of 6.4 x 1014 Hz was measured to have an energy of 4.2 x 1019 J.

Use Planck’s constant to show this:

E = (6.63 x 10-34 J·s) x (6.4 x 10141/s) = 4.2 x 1019 J

Evidence for photons
• Photoelectric effect – the ejection of electrons from a metal when exposed to EM radiation
• Each substance has its own “threshold” frequency of light needed to eject electrons
Determining the energy of a photon

E = hv

• Use Planck’s constant!
• What is the energy of a photon of radiation with a frequency of 5.2 x 1014 waves per second?
Another problem involving photon energy
• What is the energy of a photon of radiation with a wavelength of 486 nm?
Louis de Broglie – proposed that matter and radiation have properties of both waves and particles (Nobel Prize 1929)
• Calculate the wavelength of a hydrogen atom moving at 7.00 x 102 cm/sec
• = h
• m

m = mass

h = Planck’s constant

 = velocity

Hydrogen spectral lines

Balmer series:

n1 = 2 and n2 = 3, 4, …

Lyman series (UV lines):

n1= 1 and n2 = 2, 3, …

Atomic Spectra and Energy Levels
• Observe the hydrogen gas tube, use the prism to see the frequencies of EM radiation emitted
• Johann Balmer– noticed that the lines in the visible region of hydrogen’s spectrum fit this expression:

v= (3.29 x 1015 Hz) x 1 - 1

4

n2

n = 3, 4, …

v= RH x 1 - 1

n12

n22

RH = 3.29 x 1015 s-1

Rydberg Constant

• Energy of an electron in a hydrogen atom

-2.178 x 10-18 joule

E =

n2

n= principal quantum number

Differences in Energy Levels of the hydrogen atom

Use the Rydberg Equation

OR

Use the expression for each

energy level’s energy in the following equation:

E = Efinal – Einitial

Niels Bohr’s contribution

Assumed e- move in circular orbits about the nucleus

Only certain orbits of definite energies are permitted

An electron in a specific orbit has a specific energy that keeps it from spiraling into the nucleus

Energy is emitted or absorbed ONLY as the electron changes from one energy level to another – this energy is emitted or absorbed as a photon

Summary of spectral lines

When an e- makes a transition from one energy level to another, the difference in energy is carried away by a photon

Different excited hydrogen atoms undergo different energy transitions and contribute to different spectral lines

The Uncertainty Principle – Werner Heisenberg
• The dual nature of matter limits how precisely we can simultaneously measure location and momentum of small particles
• It is IMPOSSIBLE to know both the location and momentum at the same time
Atomic Orbitals – more than just principal energy levels
• Erwin Schrodinger (Austrian)
• Calculated the shape of the wave associated with any particle
• Schrodinger equation – found mathematical expressions for the shapes of the waves, called wavefunctions(psi) 
Born’s contribution
• Max Born (German)
• The probability of finding the electron in space is proportional to

2

Called the “probability density” or “electron density”

• s – high probability of e- being near or at nucleus

ELECTRON IS NEVER AT THE NUCLEUS IN THE FOLLOWING ORBITALS:

• p – 2 lobes separated by a nodal plane
• d – clover shaped
• f – flower shaped
• Each orbital can hold 2 electrons
• Orbitals in the same subshell have equal energies
Quantum numbers – like an “address” for an electron

n = principal quantum number

As n increases

* orbitals become larger

• electron is
• farther from nucleus more often
• higher in energy
• less tightly bound to nucleus
Quantum numbers
• l = angular momentum quantum number
• Values: 0 to n – 1
• Defines the shape of the orbital
Quantum numbers

Example: for d orbitals, m can be -2, -1, 0, 1, or 2

For p orbitals, m can be -1, 0, or 1

• ml=the magnetic quantum number
• Orientation of orbital in space

(i.e. pxpy or pz)

• Values: between – l and l, including 0
Quantum numbers
• ms = the spin number
• When looking at line spectra, scientists noticed that each line was really a closely-spaced pair of lines!
• Why? Each electron has a SPIN – it behaves as if it were a tiny sphere spinning upon its own axis
• Spin can be + ½ or -1/2
• Each represents the direction of the magnetic field the electron creates

Principal level 4

4p orbital

px orbital

spin up

n = 4, l = 1, ml = -1, ms = +1/2

Are these sets of quantum numbers valid?
• 3, 2, 0, -1/2
• 2, 2, 0, 1/2

NO!

Level 2

2d orbital – does

not exist!

YES!

Level 3

3d orbital

3dxz

Spin down

Electron configuration: rules
• Aufbau principle – electrons fill lowest energy levels first
• Pauli exclusion principle – only 2 electrons may occupy each orbital, must have opposite spins
• Hund’s rule – the lowest energy is attained when the number of electrons with the same spin is maximized

(because electrons repel

each other)

Energy level specifics

4s

• s and d orbitals are close in energy
• Example
• 4s electrons have slightly lower energy than 3d electrons
• The s electrons can penetrate to get closer to the nucleus, giving them slightly lower energy

3d

Noble Gas Configuration
• A shorter electron configuration
• Write the symbol for the noble gas BEFORE the element in brackets
• Write the remainder of the configuration
• Examples:
• Cl
• Cs
Special rules
• One electron can move from an s orbital to the d orbital that is closest in energy
• Only happens to create half or whole-filled d orbitals
• Examples: Cr, Cu