2009 Chemistry I: Modern Chemistry Holt, Rinehart, & Winston
Chapter 4: Arrangement of Electrons in Atoms, pp 97-122
Vocabulary
1 angular momentum quantum #; 2 Aufbau principle; 3 continuous spectrum; 4 electromagnetic radiation;
5 electromagnetic spectrum; 6 electron configuration; 7 emission-line spectrum; 8 excited state; 9 frequency;
10 ground state; 11 Heisenberg uncertainty principle; 12 Hund’s rule; 13 magnetic quantum # 14 noble gas;
15 noble gas configuration; 16 orbital; 17 Pauli exclusion principle; 18 photoelectric effect; 19 photon;
20 principal quantum #; 21 quantum; 22 quantum #; 23 quantum theory; 24 spin quantum #; 25 wavelength
1 value indicating the shape of the orbital; l = 0 & all positive integers < n -1
2 an electron occupies the lowest energy orbital that can receive it
3 the emission of a continuous range of frequencies of electromagnetic radiation
4 a form of energy that exhibits wavelike behavior as it travels through space
5 all the forms of electromagnetic radiation, collectively
6 the arrangement of electrons in an atom
7 a diagram indicating the degree to which a substance emits radiant energy with respect to wavelength
8 a state in which an atom has a higher potential energy than it has in its ground state
9 the # of waves that pass a given point in a specific time, usually 1 second; Hz = 1 wave/sec
10 the lowest energy state of an atom
11 ‘tis impossible to know simultaneously both the position & velocity on an electron or any particle
12 orbitals of equal energy are each occupied by 1 e- before any receives a 2nd e- & all have same spin#
13 value indicating the orientation of an orbital around the nucleus; m = whole #s, - l to + l, including 0
14 Group 18 elements (He, Ne, Ar, Kr, Xe, Ra)
15 Notation in which only the outer main energy level electron occupation is identified; K’s = [Ar]4s1
16 a 3-D region around the nucleus indicating the probable location of an electron
17 no 2 electrons in the same atom can have the same set of 4 quantum #s; only 2 e- fit into each orbital
18 emission of e- from materials when light of certain frequencies shines on the surface of the material
18 the emission of electrons from metal when light shines on that metal
19 a particle of electromagnetic radiation having 0 mass & carrying a quantum of energy
20 value indicating the main energy level occupied by an electron; n = 1, 2, 3, 4....
21 the minimum quantity of energy that can be gained or lost by an atom
22 values that specify properties of atomic orbitals & the properties of electrons in orbitals
23 the study & mathematical description of the structure & behavior of atoms & subatomic particles
given that all energy comes in tiny, indivisible bundles
24 value indicating 1 of 2 fundamental spin states of an electron in an orbital; s = + 1/2
25 the distance between corresponding points on adjacent waves (crest to crest; trough to trough)
Main Ideas
Q: What prevents the negative electrons from being drawn into the positive nucleus?
A: We need a better atomic model, one based on the absorption & emission of light by matter.
pp 97-103
I. New Atomic Model
A. Properties of Light
1. Electromagnetic radiation
a. Exhibits wavelike behavior as it travels thru space
1. Repetitive nature characterized by wavelength () & frequency in Hz (1/sec)
 

cwhere c = speed of light
b. Travels at a constant speed of 3.00 x 108 meters/sec thru a vacuum
c. Spectrum, from highest  to lowest:  rays, X rays, UV, visible, IR, microwave & radio
B. Photoelectric Effect
1. Emission of electrons when light shines on metal
a. Noted by Max Planck (1900) that when metals are heated, light is emitted.
b. Planck proposed energy is emitted in quanta, or small, specific packets of energy
c. Planck concluded that E = h , where E = energy & h = Planck’s constant = 6.626 x 10-34 Js
2. Requires a minimum frequency regardless of the light’s intensity
a. This could not be explained by simple wave dynamics.
3. Albert Einstein proposed light had a dual nature: wave & particle (dubbed a “photon”)
a. E photon = h 
b. The minimum energy needed to eject electrons corresponded to a minimum frequency.
c. Because different metals bind their electrons to varying degrees, each has own minimum .
C. Hydrogen Atom Emission-line Spectrum
1. When current is passed thru vacuum tube of H gas at low pressure, the tube has a pinkish glow.
2. When a narrow beam of this was shined thru a prism, it separated into specific wavelengths.
a. Balmer series: 4 specific colors of the visible spectrum
b. Lyman series: 5 specific  in the UV spectrum
c. Paschen series: 3 specific  in the IR spectrum
3. Classical theories, predicting a continuous spectrum, failed to explain these results.
4. Neils Bohr (1913) developed atomic model linking emission spectrum to electrons in orbits.
a. Orbits were discrete paths of a definite, fixed energy.
b. Ground state = lowest energy = electron closest to the nucleus
c. Energy levels were fixed, like rungs on a ladder.
d. When energy is added to raise an electron to higher level, absorption occurs.
e. When electron falls from excited state back to lower level, a photon is emitted = emission.
f. Energy of absorption = emission = E photonh
  All of this is great; however, it did not apply to other atoms nor explain their chemical behavior.
Q: If light had a dual wave/particle nature, did electrons themselves (and anything else) have that too?
A: We need even a BETTER model, one dealing with quanta of energy.
pp 104-110
II. Quantum Model of the Atom
A. Electrons as Waves
1. Louis de Broglie (1924) proposed electrons have wavelike properties.
a. Electron beams can be diffracted, or bent, passing by the edge of X or thru a small opening.
b. Electron beams can interfere with each other; interference occurs when waves overlap.
B. Heisenberg Uncertainty Principle
1. Werner Heisenberg (1927) realized photons of light detects the presence of electrons, but the
energy of photons & electrons are so close that the photon itself knocks the electron away.
2. He concluded it was impossible to know simultaneously both position & energy of electrons.
C. Schrödinger Wave Equation
1. Erwin Schrödinger (1926) developed an equation treating electrons as waves.
2. Whereas Bohr assumed quantization, Schrödinger’s equations reconciled data with theory.
3. He & Heisenberg laid the foundation for the modern Quantum theory.
D. Atomic Orbitals & Quantum Numbers
1. The 1st 3 Quantum numbers result directly from solutions to the Schrödinger equation.
a. Principle quantum # = n = main energy level occupied by the electron (1, 2, 3, 4...)
 The total # of orbitals = n2
b. Angular momentum quantum # = l = shape of orbital = 0, n-1 [0 = s, 1 = p; 2 = d; 3 = f]
c. Magnetic quantum # = m = orientation of orbital = 0, + l [s has 1, p has 3, d has 5, f has 7]
2. The 4th Quantum number describes its spin state, either clockwise or counterclockwise.
a. Spin quantum # = s = either + ½ or – ½
b. A single orbital holds a max of 2 electrons, but they must have opposite spin states.
Q: How do we know where the electrons exist in any given atom?
A: We follow a few simple rules, with noted quantum # relationships above, & figure it out!
pp 111-122
III. Electron Configurations
A. Rules
1. Aufbau principle
a. Electrons occupy the lowest-energy orbital possible.
b. Orbitals, in order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
2. Pauli exclusion principle
a. No 2 electrons in the same atom can have the exact same set of 4 quantum numbers.
b. So, you can only have 2 electrons per orbital/shape/orientation & these are opposite spins.
3. Hund’s rule
a. Orbitals of equal energy keep unpaired electrons as long as possible, all with parallel spins.
b. This means, for the p orbital, px, py, & pz, that each gets 1 electron before any gets 2.
c. Same applies for d orbitals, dx2-y2, dxy, dyz, dxz, & dz2.
d. This is also true for the f orbitals, of course.
B. Representing Electron Configuration (3 Methods)
1. Orbital Notation
a. Unoccupied orbitals are represented by a line, ___, with orbitals name written under it. ___
b. An orbital containing 1 electron is shown as ___, & 2 electrons as ___. H= ____; He= ___
2. Electron Configuration Notation
a. The # of electrons in a sublevel is shown as a superscript to the sublevel designation.
b. For example, H = 1s1; He = 1s2; Na = 1s2, 2s2, 2p6, 3s1, etc...
 Elements of the 2nd Period
1. The highest-occupied energy level = e- containing main E level with the highest n value
2. The inner-shell electrons = those that are not in the highest-occupied energy level
3. B-Ne: 1s2, 2s2, 2px1.. 2px12py1.. 2px12py12pz1.. 2px22py12pz1.. 2px22py22pz1..2px22py22pz2 or 2p6
4. Group 18 (Group VIII) elements have 8 electrons in outermost shell, i.e., a full octet <not He>
3. Noble-Gas Notation
a. After Ne’s full octet, all the electrons of the 3rd Period (Na-Ar) have their 1st 10 e- the same.
b. Rather than rewrite Ne’s electron configuration each time, it is abbreviated by [Ne]...
 [Ne] = 1s2, 2s2, 2p6
 Na = [Ne]3s1; Mg = [Ne]3s2; Al = [Ne]3s23px1, ... Ar = [Ne] 3s23p6
c. Elements of the 4th Period begin with 4s then go to 3d (transition elements) & then to 4p
 Exceptions!!!! Cr = [Ar]4s13d5 and Cu = [Ar]4s13d10 ... and we don’t know why!
d. Elements of the 6th period: La adds an e- to 5d but with Ce-Lu, electrons add to the f orbital.
 Because 5d & 4f are so close in energy, deviations from “rules” are numerous.
IV. Blocks
A. The periodic table, subdivided into ‘blocks’, aids in determining configurations. (See below.)
B. You never need to memorize an element’s configuration, except Cr & Cu, just follow the table!
(Really, it’s that easy.)
s
d
p
f