Post on 20-Jan-2018
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Bohr’s Model of the Atom • 1913
Science is not a one-man showThe work of other scientists help
lay a foundation for new discoveries
Bohr set out to explain
• The relationship between
- the emission lines observed in excited element’s spectrum and
- the element’s atomic structure
Red, green, blue and violet spectral lines
Classically, light is thought to be a wave• Light or electromagnetic
radiation is a wave in which
- electric field oscillates perpendicular to magnetic field through space
Wavelength (symbol: , called lambda)
= distance between any 2 adjacent identical points of a wave (i.e. between 2 crests or 2 troughs)
Frequency of a wave = # of wavelengths of that wave passing a fixed point in one second
Frequency symbol: , called nu or new Unit: 1/second or Hertz (Hz)
The shorter the wavelength, the higher the frequencyThe longer the wavelength, the lower the frequency
Math. relationship b/t wavelength and frequency
where c = 3.00 x 108 m/s = speed of light in a
vacuum
c = x
Or
c is fixed:
• If increases, will decrease
• If decreases, will increase
Rethinking light: Light also behave as a particle1. Blackbody radiation Widely accepted: - A heated object emits
light- The hotter it gets, higher
energy re-radiated e.g. UV, X-ray, Gamma rays but no object does this
E photon =h
- Planck proposed: light is not continuous as previously thought but discrete bundles of energy, called quanta
h = Planck’s constant = 6.626×10 −34 J s)⋅
Rethinking light: Light also behave as a particle
2. The photoelectric effect:
• Certain frequencies of light were able to knock electrons off the surface of metals
• Other frequencies regardless of the intensity of light, were unable to do so.
Bohr’s simple assumption
By now Bohr has inherited the ideas that- Light is quantized- Spectroscopy is a powerful tool to infer about
atomic structure
- Adding his own assumption: What if some aspects of atomic structure, such as electron orbits and energies, could only take on certain values?
What is spectroscopy?
• Analysis of the way matter absorbs or releases radiation
• Spectroscopy is used to:a) identify elements or compoundsAtomic Spectra: Fingerprints of Elements
Spectroscopy is used to:• b) obtain information about bonding and
structures of compounds
IR spectrum
Spectroscopy is used to:• c) Quantitatively determine the concentrations
of substances present.
A simple spectroscope
Two main types of atomic spectra:Emission vs. Absorption Spectra
“Bright-line Spectra” is atomic emission spectra in the visible range.
Absorption vs. Emission Spectra of Hydrogen
Note that the spectral lines are on the same wavelength.
Emission spectrum:As excited electrons relax to lower energy levels, light is emitted. This release of energy explains for the color lines on the emission spectrum of hydrogen atom.Many more lines but they’re in IR and UV regions and thus are not visible.
Absorption spectrum:Observed when we see white light passing through a glass tube containing Hydrogen gas. The Hydrogen atom absorb very specific wavelength of radiation resulting in missing lines in the continuous spectrum.
Advantages/Uses of Spectroscopy• a) can definitively distinguish between substances
with very similar physical and chemical properties.
eg. Members of the alkali metals family
K Na
Advantages/Uses of Spectroscopy• b) Sample is distant • Eg. Applications in Astronomy
Astronomers have made the first direct detection and chemical analysis of an atmosphere of a planet that exists outside our solar system.
Spectrum of the hot gases in a nearby star-forming region, the Omega Nebula (M17)
Advantages/Uses of Spectroscopy• b) Sample quantity is limited • eg. in Forensics “Crystal meth”
Back to Bohr’s model
Emission spectrum of heated hydrogen gas
• Bohr’s postulate #1:An electron can only sit in specific energy levels.
Energy levels of the e- in the H atom can be calculated as (R = Rydberg constant expressed in J = 2.18 x 10-18 J)
n = Principal Quantum number;
- Can only be integer 1, 2, 3…∞- indicates the orbit.
Energy also in J
A note about units of energy
• Joule is too big of a unit when dealing with e-• New energy unit: The electron-volt (eV) • Define: • 1eV ≈ 1.6 x10-19 J= the kinetic energy gained by an electron when accelerated through 1 volt of potential difference
Can also be expressed as
When n =1, lowest possible energy or ground state energy of a hydrogen electron = -13.6 eV
Where E expressed in J Where E expressed in eV
Why the negative sign?
Because:- the energy of an electron in orbit is compared with the energy of an electron that is free from its nucleus (n=∞) This free e- has an energy = 0
- e- in orbit is more stable than e- infinitely far from nucleus, the energy of an electron in orbit is always negative.
Where E expressed in J
Where E expressed in eV
Energy level in Hydrogen atom
How to get this value?
Note that because of the negative sign, E1 is the lowest in energy (instead of the highest)
• Absorption of Energy = Excitation
The quantum leap
If the incoming energy is exactly equal to the difference in energy between two orbits, the electron leaps to excited state
If the incoming energy is a little less or greater, it passes through the atom unaffected
When moving within these allowed energy states (stationary states) the electron does NOT emit (release) energy.
• Emission of Energy = Relaxation
The quantum leap
• As excited e- relaxes to lower energy level, it releases the energy difference between the two orbits in the form of electromagnetic radiation.
• This explains spectral lines of excited atoms and the origin of light itself.
Energy of emitted photon = h = Efinal - Einitial
Bohr’s postulate #2: Electron transitions
Excitation RelaxationElectron moves to
higher E level
lower E level
Energy is
Absorbed Released
Form of Energy
Heat, light, electrical
Electro-magnetic radiation
Relaxation from higher orbit to
Energy emitted Name of the series
n=1 UV Lymann=2 Visible Balmern=3 IR Paschenn=4 IR Brackett
What happens to a sample of hydrogen gas (in a discharge tube) when an electric current is run through
it?
• Electrical energy is absorbed and excites the electrons
• Relaxation follows
• H2 emits a purplish light is emitted.
• Color light perceived by the eyes
https://youtu.be/955snB6HLB4
Emission in real life
• Flame test helps identify the metals
• Fireworks• LED light vs sodium vapor
street lamps
Emission in real life
Coop. Learning
When E= -13.6/ n2 in eV In Joules1eV = 1.6 x10-19 J
n=1 n=2
n=3
n=4
n=5
n=6
n= infinity
When E= -13.6/ n2 in eV
In Joules1eV = 1.6 x10-19 J
n=1
E = -13.6 eV E= -13.6 x 1.6x 10-19 = 2.176x 10-18J
n=2 E= -13.6/4 eV E= -5.44 x 10-19 J
n=3 E= - 1.51 eV E= -2.42 x 10-19 J
n=4 E= -0.85 eV E= -1.36 x 10-19 J
n=5 E= -0.54 eV E= -8.70 x 10-20 J
n=6 E= -0.38 eV E= -6.04 x 10-20 J
n= infinity E= 0 E= 0 J