lab 8 atomic structure
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Lab 9
Atomic Structure Emission Spectrum
Electron Configuration
HISTORY OF THE ATOMHISTORY OF THE ATOM
460 BC
Democritus develops the idea of atoms
he pounded up materials in his mortar
and pestle until he had reduced them to
smaller and smaller particles which he
called
ATOMSATOMS
(greek for indivisible)
Development of the Model of an Atom
400 BC Democritus – Particle Model
Ancient Greek philosopher Proposes that matter is composed of smallest
particles called atoms An idea with no evidence
Aristotle – Continuous Model Alternative idea Continuous model of matter- no smallest piece
HISTORY OF THE ATOMHISTORY OF THE ATOM
1808 John Dalton
suggested that all matter was made up of
tiny spheres that were able to bounce
around with perfect elasticity and called
them
ATOMSATOMS
Dalton proposes Atomic Theory in 1803 Based upon experimental evidence All matter composed of atoms Atoms of same element have same mass
and properties Atoms are neither created nor destroyed,
they are rather simply rearranged in chemical reactions Evidence – Lavoiser- Conservation of Mass
Compounds are composed of elements in simple whole number ratios Evidence – Proust
HISTORY OF THE ATOMHISTORY OF THE ATOM
1898 Joseph John Thompson
found that atoms could sometimes eject a
far smaller negative particle which he
called an
ELECTRONELECTRON
Thomson and the Discovery of Electrons
J. J. Thomson’s Experiment
Devised an experiment to find the ratio of the cathode ray particle’s mass (me) to the charge (e)
me /e = –5.686 x 10–12 kg C–1
Thomson’s Plum Pudding Model
Based upon the charge to mass ratio, the electron must be much smaller than the atom
Proposed negative electrons embedded in positive matrix of atom
HISTORY OF THE ATOMHISTORY OF THE ATOM
Thomson develops the idea that an atom was made up of
electrons scattered unevenly within an elastic sphere
surrounded by a soup of positive charge to balance the
electron's charge
1904
PLUM PUDDING
MODEL
Millikan’s Oil Drop Experiment
Measuring the Charge on an electron
Unstable Atoms and Radioactivity
Atoms are not indestructible Atoms are composed of smaller
particles Alpha particles – positively charged Beta particles – negatively charged Gamma rays – no charge
Radioactivity
Goldstein’s Discovery of Protons
Mass Spectrometer- Determining the Percent Abundance of Different Isotopes of same element
Mass Spectrometer
If a stream of positive ions having equal velocities is brought into a magnetic field, the lightest ions are deflected the most, making a tighter circle
Mass Spectrometry
EOS
A record of the separation of ions is called a mass spectrum
Isotopes of Neon Neon-20
10 protons 10 neutrons
Neon-21 10 protons 11 neutrons
Neon-22 10 protons 12 neutrons
HISTORY OF THE ATOMHISTORY OF THE ATOM
1910Ernest Rutherfordoversaw Geiger and Marsden carrying out
his famous experiment.
They fired Helium nuclei at a piece of gold
foil which was only a few atoms thick.
They found that although most of them
passed through. About 1 in 10,000 hit
Rutherford’s Gold Foil Experiment
Rutherford’s Gold Foil ExperimentRutherford’s Gold Foil Experiment
gold foil
helium nuclei
They found that while most of the helium nuclei passed
through the foil, a small number were deflected and, to
their surprise, some helium nuclei bounced straight back.
Rutherford’s Nuclear Model of the Atom
Most the alpha particles (helium nuclei) pass through the gold foil Atom mostly empty space Alpha particles did not hit anything
A very few deflected straight back Alpha particles deflected by a dense
positive nucleus
Visualizing the Pathway of Alpha Particles through a Gold Atom
ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION
Electromagnetic radiation.
Electromagnetic RadiationElectromagnetic RadiationElectromagnetic RadiationElectromagnetic Radiation
Most subatomic particles Most subatomic particles behave as PARTICLES and behave as PARTICLES and obey the physics of waves.obey the physics of waves.
Electromagnetic Radiation
The Electromagnetic Spectrum
Wave Model of Light
Wavelength measured in meters (m)
Frequency ƒ measured in waves per second (Hz)
Energy E measured in Joules (J)
Wavelength, Frequency and Energy
Wavelength inversely related to frequency
Increasing wavelength decreasing frequency Decreasing wavelength increasing frequency Wavelength x frequency = speed of light c = ƒ c = speed of light = 3.00 x 108 m/s
Frequency directly related to energy Increasing frequency increasing energy Energy = Planck’s constant x frequency E = h ƒ h = Planck’s constant = 6.63 x 10-34 J/Hz
wavelength Visible light
wavelength
Ultaviolet radiation
Amplitude
Node
Electromagnetic RadiationElectromagnetic RadiationElectromagnetic RadiationElectromagnetic Radiation
ElectroElectromagneticmagnetic SpectrumSpectrumElectroElectromagneticmagnetic SpectrumSpectrum
In increasing energy, RIn increasing energy, ROOYY GG BBIIVV
Sunlight viewed through a spectroscope
Prisms and diffraction grating
bend light
Red light with longer wavelengths bend less
Violet light with shorter wavelengths bend more
Separates white light into ROYGBIV Red – Orange – Yellow – Green – Blue – Indigo – Violet
long waves short waveslow frequency high frequencylow energy high energy
Photoelectric Effect
Photoelectric Effect
Bright red light shined on the photocell has no effect- regardless of intensity or time
Dim green light shined on the photocell causes electrons to be emitted and flow through the wire
Brighter green light emits more electrons per second- greater current
Explaining the Photoelectric Effect ONLY the photon or particle model can
be used to explain these results Energy is required to pull off negative
electrons attracted to positive protons in nucleus- breaking attraction
Each red photon does not have enough energy to pull off an electron – regardless of how long the light is shined
Each green photon has more energy and can pull off the electron when the green photon collides
Light Spectrum Lab!
Slit that Slit that allows light allows light insideinside
Line up the slit so Line up the slit so that it is parallel with that it is parallel with the spectrum tube the spectrum tube (light bulb)(light bulb)
The Emission Spectrum of Hydrogen- Discrete Bands of Colored Light
Excited Gases Excited Gases & Atomic & Atomic StructureStructure
Emission Spectra of Different Atoms: A Fingerprint to Identify
Rydberg and Balmer (1886)
Independently develop mathematical equations that fit the data for hydrogen emission spectrum
The electron had no yet been discovered
Neither had a model to explain the observed wavelengths
Just an equation that worked
Rydberg Equation
1/λ = RH [1/n12 - 1/n2
2]
RH = 1.09678 x 10-2 nm-1
Solve the equation for an electron
moving from level 4 to 2.
Bohr’s Model of the Atom (1910)
Assumed electrons orbit the nucleus in circular orbits
Proposed the energy of the orbit is proportional to the distance from the nucleus (increasing distance – increasing energy)
Assumed only certain allowable energies
Used angular momentum to calculate the allowable energy
Bohr’s Model of the Atom (1910)
When the atom absorbs energy Electron moves up to higher energy with
more potential energy farther away from nucleus
Unstable with higher PE Electron falls back down to lower levels
Energy released PE converted to KE as electron fall The color of light observed reflects the
energy released in the fall
Bohr’s Calculations
of the Energy
ΔE = -2.18 x 10-18 J (1/nf2 – 1/ni
2)
n = the energy level
ΔE = positive when electron climbs up levels
absorbing energy
increasing PE
ΔE = negative when e- falls down levels
releasing energy
decreasing PE
Niels BohrNiels Bohr
(1885-1962)(1885-1962)
ΔE = -2.18 x 10-18 J (1/nf2 – 1/ni
2)
Calculate the energy as an electron drops from level 6 down to level 2.
Calculate the frequency and wavelength of this photon.
ultraviolet infrared
Visualizing the Movement of the
Electron
Line Spectra of Line Spectra of Other ElementsOther Elements
Oops. Bohr’s equation does NOT predict these wavelengths.
Visualizing the “falling” e-
Where does the electron have more potential energy?
Electron is a wave - De Broglie
De Broglie Since light is both particle and wave,
perhaps so is matter both particle and wave Wavelength depends upon mass and
velocity Objects with large mass have negligible, so small we can ignore Electrons with very small mass have wave
properties that cannot be ignored
λ = h/p
Quantum Mechanics Heisenberg
Uncertainty principle Given the wavelike nature of electron Impossible to know both location and energy of
electron Can only calculate the probable location
Schrodinger Used calculus to “locate” the electron within
orbital Orbitals – regions of space representing the most
probable location of electron
E. SchrodingerE. Schrodinger1887-19611887-1961
W. HeisenbergW. Heisenberg1901-19761901-1976
Wave Functions: Calculating the Probability of locating an electron in a region of space
The Uncertainty Principle: Cannot both determine location and energy of electron
The Wave Function and Orbitals
The region near the nucleus is separated from the outer region by a spherical node - a spherical shell in which the electron probability is zero
EOS
Quantum Numbers
Values that emerge from the wave functions of Schrodinger
1st n = energy level 2nd ℓ = shape of orbital (s, p, d or f) 3rd mℓ = orientation (diff
versions) 4th ms
= magnetic spin of electron
Increasing Radius of s-orbital with higher values of n
s orbitals orbital p orbitalp orbital d orbitald orbital
f Orbitalsf Orbitalsf Orbitalsf Orbitals
The s-orbital
Spherical shaped orbital
ℓ = 0
mℓ = 0
Only one s-orbital in any energy level
The p-orbitalDouble-lobe shaped orbitalℓ = 1mℓ = -1 or 0 or +1Only three p-orbitals in any energy level- except for level one
Planes of zero probability
Model of d-orbital
Only electrons with opposite spins can be in the same orbital
Electron Configurations
Show the electrons in orbitals Box used to represent orbital Half arrow used to represent e- with opposite spins
Electrons are placed in orbitals of lowest energy first
Use sum of first two quantum numbers to determine which orbital fills first
1s
2s
3s3p
2p
1s
2s
3s3p
2p
Nickel Electron Configuration and quantum numbers
1s 2s 2p 3s 3p
1 0 0 ½ 2 0 0 ½ 2 1 -1 ½ 2 1 0 ½ 3 0 0 ½ 3 1 0 ½
2 1 0 ½ 3 1 -1 ½ 3 1 0 ½
4s 3d
4 01 0 ½ 3 2 -2 ½ 3 2 -1 ½ 3 2 0 ½
3 2 1 ½ 3 2 1 ½
1st # indicates energy level
n = 1 1st level
n = 2 2nd level
2nd # type of orbital
l = 0 is s-orbital
l = 1 is p-orbital
l = 2 is d-orbital