AP CHAPTER 7 OUTLINE
ATOMIC AND ELECTRONIC STRUCTURE
III. Wave Properties of Matter and Wave Mechanics
A. Problem with Classical Physics
1. Classical Physics only works for large objects because small objects (e-) behave like waves instead of matter
a. DeBroglie's equation
Lambda = h/mv (h = plank's constant, m = mass, v = velocity, and lambda = wavelength)
b. Large objects have short wavelengths (small matter waves)
c. Small objects have long wavelengths (large matter waves)
2. Diffraction -- reinforcement or cancellation of wave intensity
a. Principal by which the electr microscope is based on
B. Properties of Waves
1. Traveling waves -- matter moves up and down as crest and trough move horizontally
2. Standing wave -- crest and nodes do not change position
a. Node -- points of zero amplitude and are fixed positons
b. Naturally leads to quantum numbers since the length of the string is a whole number multiple of half-wavelengths
L = n(lambda/2)
C. Electron Waves in Atoms
1. Wave mechanics -- focus on wae properties of matter
2. Quantum Mechanics -- synonomous name since wave mechanics predict quantized energy levels
a. First applied in 1926 by Erwin Schrodinger
3. Standing Waves in Atoms
4. Electrons can have different wave patterns (caled orbitals) which are expressed by by a wave function (psi)
a. Energy tells us energy of waveform, shapes of waveform tells us probability of finding the electron in that position!
b. Since atoms are 3D, we need three quantum numbers
5. Principal Quantum number (n) -- called shell (energy level)
a. n = 1 to infinite, large n value, more energy, farther from nucleus (larger wave)
6. Secondary Quantum number (l) -- subshells
a. l = 0 to (n-1), letter code used for l values, determines shape of orbital, also effect energy
b. 0 = s, 1 = p, 2 = d, 3 = f, 4 = g, 5 = h, etc.
7. Magnetic Quantum number(ml) -- orbitals
a. ml = -l to +l, orientation of orbitals to each other
b. s = 1, p = 3, d = 5, f = 7, etc.
IV. Electron Spin and the Pauli Exclusion Principle
A. Magnetic Properties of the Electron
1. Electron spin based on fact that electrons act like tiny magnets
a. Revolving electrical charge creates a small magnetic field
2. Spin Quantum number (ms) -- fourth quantum number for electron spin
a. ms = -1/2 or +1/2
B. The Pauli Exclusion Principle
1. No two electrons in the same atom can have identical values for all four of their quantum numbers
a. The maximum number of electrons in any orbital is two with opposite spins
C. Magnetic Properties of Atoms
1. Paramagnetic -- Weak attraction of atoms with unpaired electrons to a magnet
2. Diamagnetic -- Atoms in which all electrons are paired, are not attracted to magnets
V. Electronic Structures of Multielectron Atoms
A. Electron configuration (electronic structure) -- distribution of electrons among the orbitals
Ex. -- H 1s1
1. Orbital diagrams
Ex. --
B. The Aufbau Principle -- electrons fill the energy levels with lowest energy first
1. Hund's rule -- when electrons are placed in a set of orbitals of equal energy, they spread out as much as possible to give as few paired electrons as possible
VI. Electron Configurations and the Periodic Table
A. Use the structure of the periodic table to predict the filling of the subshells when we write the electronn configuration of an element
1. s-block = alkali and alkaline earth metals
2. p-block = group 13-18 (metals, nonmetals, halogens, noble gases)
3. d-block = transition metals
4. f-block = inner transition metals
B. Electronic Basis for the Periodic Recurrence of Properties
1. Properties are baed on the outer shell electrons
a. Inner (core) electrons are not involved in bonding, etc.
2. elements with similar properties have similar outer shell electrons configurations
C. Abbreviatied electron Configurations
1. Shorthand configuration (noble gas method)
ex. - Na [Ne]3s1 instead of 1s22s22p63s1
D. Valence Shell Elecron Configurations
1. Valence shell -- outer energy level where electrons are located for representative elements only
a. Only consist of s and p orbitals
b. Valence electrons -- electrons in valence shell
VII. Some Unexpected Electron Configurations
A. Filled subshells are more stable then unfilled subshells
1. Half-filled subshells are more stable then partially filled subshells
2. Filled subshells are more stable then half-filled subshells
3. electrons will sometimes jump orbitals to make filled or half-filled subshells to make the overall atom more stable
Ex. - Cr and Cu
VIII. Shapes of Atomic Orbitals
A. Heisenberg Uncertainty Principle -- can not measure a particles velocity and position at the same time
1. The statistical probability of finding the electron in a given position at a particular instant
2. Where probability is high, the amplitude of the wave is large
a. Electron cloud -- 3d space around an atom where you will find electrons in a given orbital
b. The closer to the nucleus, the greater the probability of finding the electron
3. high probability = high concentration of electrical charge, large electron density
c. Electron density -- how much of the electron's charge is packed into a given volume
B. Shapes and Sizes of s and p Orbitals
1. Shape and size based on 90% probability of finding electron
a. s - orbital = spherical, size increase by n, 2s and up have nodes (like wave amplitudes)
b. p - orbital = two lobes on opposite sides of the nucleus, nodal plane between them perpendicular to the nucleus, size also increases by n
C. Orientation of Orbitals in a p subshell
1. Three orbitals, 90 degrees from each other
a. Orbitals are on xyz coordinates
D. Shapes and Orientations of d orbitals in a d subshell
1. 5 orbitals, much more comlex
2. 4 of the orbitals have four lobes in different coordinate directions
3. 1 orbital has two lobes on z axis with donut ring around it in the x-y plane
IX. Variation of Atomic Properties with Electronic Structure
A. Effective Nuclear Charge -- net charge that outer shell electrons fells from the nucleus
1. Inner core electrons "shield" the nuclear charge from the outer shell electrons
2. ENC is difference between nuclear charge and the inner core charge
B. Sizes of Atoms and Ions
1. Angstrom (A) -- unit used for atomic length, 1 A = 1 x 10-10 m
a. 1 A = 100 pm 1 A = 0.1 nm
C. Periodic Variations in Atomic Size
1. Atomic radii increase going down the periodic table
a. Add one principal energy level in every period
2. Atomic radii decrease going across the periodic table
a. Increase in Effective Nuclear Charge due to more protons and electrons
3. Smaller changes in transition metals, since they are filling inner shells, not outer shells
D. Trends in the Sizes of Ions
1. Negative ions are always larger than atoms
a. Increase in electron-electron repulsion
2. Positive ions are always smaller than atoms
a. Decrease in electron-electron repulsion
E. Ionization Energy (IE) -- energy required to remove an electron from an isolated, gaseous atom or ion in its ground state
1. The amount of work that is required to pull an electron from the atom (how tightly the electrons are held by the atom)
a. First ionization energy -- energy needed to remove the first electron
b. Second ionization energy -- energy needed to remove the second electron, etc.
c. Successive ionization energies always increase because each subsequent electron is pulled away from an increasingly more positive ion, and that requires more energy
F. Periodic Trends in Ionization Energies
1. Ionization energy decreases as you go down the group
2. Ionization energy increases as you go across the group
G. Stability of the Noble Gas Configuration
1. Extremely difficult to break into the noble gas core
H. Electron Affinity (EA) -- PE change associated with the addition of an electron to a gaseous atom or ion in the ground state
1. Exothermic processs
2. EA decreases as you go down the periodic table
3. EA increase as you go across the periodic table
Outline based upon:
Brady, J. E., Holum, J. R., Russell, J. W. (2000). Chemistry: The Study of Matter and Its Changes. (3rd ed.). New York: John Wiley & Sons, Inc. pp. 275-314.
