atomic structure and the periodic table n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Atomic Structure and the Periodic Table PowerPoint Presentation
Download Presentation
Atomic Structure and the Periodic Table

Loading in 2 Seconds...

play fullscreen
1 / 36

Atomic Structure and the Periodic Table - PowerPoint PPT Presentation


  • 98 Views
  • Uploaded on

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

loader
I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
capcha
Download Presentation

PowerPoint Slideshow about 'Atomic Structure and the Periodic Table' - reilly


An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript
studying atoms by analyzing light emissions absorbtions
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
electromagnetic radiation
Electromagnetic radiation
  • 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
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
Visible Light
  • Wavelengths from 700 nm (red) to 400 nm (violet)
  • No other wavelengths are visible to humans
quanta and photons
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
Another analogy for quanta
  • A person walking up steps – his potential energy increases in a quantized manner
photons
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
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
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
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
Another problem involving photon energy
  • What is the energy of a photon of radiation with a wavelength of 486 nm?
slide14
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
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
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, …

rydberg equation works for all lines in hydrogen s spectrum
Rydberg equation: works for all lines in hydrogen’s spectrum

v= RH x 1 - 1

n12

n22

RH = 3.29 x 1015 s-1

Rydberg Constant

energy associated with electrons in each principal energy level
Energy associated with electrons in each principal energy level
  • 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
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
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
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 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
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
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”

atomic orbital the wavefunction for an electron in an atom
Atomic Orbital – the wavefunction for an electron in an atom
  • 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
more about orbitals
More about orbitals
  • Each orbital can hold 2 electrons
  • Orbitals in the same subshell have equal energies
quantum numbers like an address for an electron
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
Quantum numbers
  • l = angular momentum quantum number
    • Values: 0 to n – 1
    • Defines the shape of the orbital
quantum numbers1
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 numbers2
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
describe the electron that has the following quantum numbers
Describe the electron that has the following quantum numbers:

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
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
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
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
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
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