line spectra and the bohr model
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Line Spectra and the Bohr Model
Flame Test Colors
• Color’s are emitted when atoms are given energy
• The actual color seen is a mixture of many colors
Atomic Emission Spectrum
Atoms Give Off Colors of Light
• AES: The set of frequencies (colors of light) emitted by atoms of an element
• Each element’s atomic emission spectrum is unique (Fingerprint of the element)
• Line Spectrumhttp://www.colorado.edu/physics/2000/quantumzone/index.html
Observing the Spectrum• Spectrascope: Acts like a prism to
split light into its actual colors– Our eye sees the colors given off all
mixed together– 2 colors could look the same to us,
but have vastly different “line spectra”
Bohr’s Model of the Atom• Atom has distinct energy levels,
(Orbits) starting with n=1 then n=2, n=3…
• Ground State: lowest energy level• When excited, it jumps to a higher
state (excited state)
• When it goes back down, it emits energy (light)– ‘Step ladder’
Small orbit = low energy stateLarge orbit = high energy state
Bohr’s Line Spectra• Energy of light given off is due to how
far the electron is ‘falling’ through levels– Not all of it is visible
• Different jumps give different wavelengths
• Grouped in “series”– Lyman series: Emits light in the UV region– Balmer series: Emits light in the visible
spectrum– Paschen series: Emits light in the IR
region
Figure 4.16 – Prentice Hall Chemistry
Glowing• Other ways electrons emit light:
– Fluorescence • Atoms in the molecule absorb certain
wavelengths of light and emit it back out• Demo: Needs UV light to absorb/emit• The reaction is FAST
– Photoluminescence (foh-tuh-loo-muh-nes-uh ns)
• Atoms absorb wavelengths of light and emit – however the reaction is SLOW
• Charging your glowing object– Chemiluminescence
• Glowsticks
Electrons act like waves?• Louis de Broglie
– Radiant Energy waves (light) behave like a particle…. Could the opposite be true?
• Particles act like waves?– Predicted that all moving matter has
wave characteristics based on it’s velocity and mass
– Works for all matter traveling at speed v, only relevant to electrons however
=hmv
De Broglie Waves- What is the wavelength of an electron
moving with a speed of 5.97 x 106 m/s? - (The mass of the electron is 9.11 x 10–
28 g)
- Basis for electronmicroscope
Modern Quantum Theory of Electrons
Probability and Orbitals
Werner Heisenberg• Heisenberg’s Uncertainty Principle
- It is impossible to know both exact speed and location of an electron
Bohr wasn’t right…• Bohr’s equations only
worked for hydrogen atoms – (1 electron)
• Too simple!
2 important Bohr ideas:1. Electrons exist only in certain discrete energy levels (described by their quantum #’s)2. energy is involved in moving an electron from one level to another
Quantum TheoryWhat do people think….
• Niels Bohr:“Fundamentally incomprehensible”
• Richard Feynman (nobel prize winner)“If you think you understand quantum
mechanics, you don’t understand quantum mechanics”
Describing Electrons using Orbitals• Atomic Orbital:
– A region around the nucleus of an atom where an electron with a given energy is likely to be found
– Different from ORBIT (the circular ring)
Orbitals• Unlike Bohr’s model, the orbital
makes no attempt to describe the electron’s path
• Recall Bohr’s levels (quantum #’s: n=1, n=2, …)
• Orbital Theory does the same but has 3 quantum #’s– N, l, ml
The 3 Quantum #’s1) Principal Quantum Number N•Can have a positive integer value n=1, 2, 3, etc•As n increases, the orbital gets larger and increases in energy
2) Azimuthal Quantum Number L (subshell)•Can have integral values from 0 to n-1 for each value of n•Defines the shape of the orbital
3) Magnetic Quantum Number Ml (orbital)•Has integral value from -L to L
Value of L 0 1 2 3
Letter used
S P D F
Sounds very confusing…• Where do you live?
– State, City, Street Name, House #• Electrons are identified the same
way..– Principle energy level, sublevel,
orbital
1. Principle energy levels (1,2,3…)2. Sublevel (s, p, d, f)3. Orbitals
Values of N, L, and MlTextbook problems 6.49-6.54
Within the sublevels…• S has 1 orbital
• P has 3 orbitals
• D has 5 orbitals
• F has 7 orbitals
– 2 electrons can fit in each orbital (box)
Why?
(a) Predict the number of subshells in the fourth shell, that is, for n = 4. (b) Give the label for each of these subshells. (c) How many orbitals are in each of these subshells?
• Solution • (a) There are four subshells in the fourth shell,
corresponding to the four possible values of l (0, 1, 2, and 3).
• (b) These subshells are labeled 4s, 4p, 4d, and 4f. • (c)
– There is one 4s orbital (when l = 0, there is only one possible value of ml: 0).
– There are three 4p orbitals (when l = 1, there are three possible values of ml: 1, 0, and –1).
– There are five 4d orbitals (when l = 2, there are five allowed values of ml: 2, 1, 0, –1, –2).
– There are seven 4f orbitals (when l = 3, there are seven permitted values of ml: 3, 2, 1, 0, –1, –2, –3).
SAMPLE EXERCISE 6.6 Subshells
Principle Energy Level• How many sublevels are there in
each energy level?– N=1 has 1 sublevel (s)
• 1s– N=2 has 2 sublevels (s and p)
• 2s & 2p– N=3 has 3 sublevels (s, p, and d)
• 3s, 3p, & 3d– N=4 has 4 sublevels (s, p, d, and f)
• 4s, 4p, 4d, & 4f
Practice Exercise 2(a) What is the designation for the subshell with n = 5 and l = 1? (b) How many orbitals are in this subshell? (c) Indicate the values of ml for each of these orbitals.
Answers: (a) 5p; (b) 3; (c) 1, 0, –1
Radial Probability Functions• The probability that an electron
(with a certain energy) will be a distance “r” from the nucleus
Node
Shapes of orbitals• S orbital
– “Sphere”
• P orbital– “Dumbbell”
• D orbital– “Flower”
What happens to the sphere as “n”
gets bigger?
Quiz on Friday!!• Shapes of orbitals
– Know what they are– Draw the s, p, and d orbitals
• Quantum #’s questions– Be able to determine & write them
Check-Up Questions• Give the numerical values of n and
l for each:– A) 2p– B) 2s– C) 4f– D) 5d
• Solution:– A) n=2, l=1– B) n=2, l=0– C) n=4, l=3– D) n=5, l=2
Sample Test Questions…1) The __________ quantum number defines the shape of an
orbital.• a) spin• b) magnetic• c) principal• d) azimuthal• e) psi
2) The n = 1 shell contains ___ orientations of p orbitals. All the other shells contain ___ orientations of p orbitals.
• a) 3, 6-9• b) 0,3• c) 6,2• d) 3,3• e) 2,6
3) In a px orbital, the subscripts x denotes the __________ of the orbital.
• a) energy• b) spin of the electrons• c) probability of the shell• d) size of the orbital• e) axis along which the orbital is aligned
Fillin’ ‘em up with electrons• Aufbau Principle
– Add one electron at a time to the lowest energy level available
• The 4th Quantum # (sorry)– Ms
describes the “spin” of an electron– Possible values: +½ or –½
• Pauli Exclusion Principle– No two electrons in an atom can have
the same set of four quantum numbers (n, l, ml, and ms)
– This means no more than 2 electrons per orbital
Using Pauli…
What is the set of quantum #’s for the last electron?
N=2, L=0, Ml=0, Ms= +½
Electron Configurations• The way electrons are distributed
among orbitals• Ground state: most stable
electron configuration• If not for Pauli, they would all
crowd the 1s… but they can’t
• The levels will fill up by lowest energy first (electrons are lazy)
Energy of each level
EN
ER
GY
1s
2s
3p3p 3p
2p2p 2p
3s
4s
3d3d 3d 3d 3d
4p4p 4p
Sneaky!
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