Chapter 7 - Atomic Structure and Periodicity Flashcards
Coulomb’s law
F=k((Q1 x Q2)/r^2)
Describes atomic radius size as a function of forces of attraction and repulsion between the nucleus and electrons and electrons and electrons, (shielding and different sized outer oribts of elements given same energy levels)
Hund’s rule
schoolbus rule, orbitals fill first half all the way through subshell until filled with half filled orbitals (in 2p, every orbital gets an “up” electron before any “down” electrons appear in any orbital)
Pauli Exclusion Principle
No two electrons in the same orbit, position, and orbital can have the same spin (arrow up and arrow down per box)
Aufbau Principle
Electrons fill in lowest energy level first (1s first, then 2s)
Electronegativity trend (give directions ex. left and down)
Right and up
Atomic size trend (give directions (ex. left and down are larger))
Left and down are larger
Does effective nuclear charge change as it moves down a group? Why?
No, as the same amount of valence electrons exist for each atom down a group. ( weirdly, as the electrons are still further away, down a group decreases electron affinity/binding energy even though the effective nuclear charge is technically the same (text for clarification if this is too confusing))
Order for determining an elements atomic size
- Energy level
a. Whether or not there are the same amount of full electron “shields” - Amount of protons
a. And effective nuclear charge acting on valence electrons - Amount of valence electrons
a. Electrons repel, protons in nucleus are less effective and attraction is more spread out thus making it bigger
4.
Metals are more prone to ______ electrons, whereas nonmetals usually _____ electrons. (gain, lose)
- lose
- gain
Cations are _____ than their original element, whereas anions are _____. Additionally, when a cation and anion are isoelectronic, the cation is substantially smaller. Why is this? (larger, smaller)
- smaller
- larger
because cations lose and electron, therefore decreasing electron-electron repulsion, whereas anions gain an electron, having the opposite effect. When isoelectronic, the cation has a much higher Zeff value (effective nuclear charge), increasing attraction and decreasing distance by Coulomb’s law.
Define isoelectronic
Having the same number of electrons (though often not protons)
In what direction does electron affinity increase? (ex. top left to bottom right)
up and to the right
Electronegativity
The tendency to pull on an electron in a bond
What changes going right on the periodic table? What changes moving down?
Effective nuclear charge increases right, energy level and thus shielding increases down
Where do metallic and nonmetallic character increase/decrease?
Metallic character increases inversely to nonmetallic character down and left
Principal quantum number represents
Shell/energy level (n)
Angular momentum quantum number represents:
Subshell/sublevel (n-1= l) (ex. p subshell, s subshell, d subshell)
Magnetic quantum number represents:
Orientations (ex. x-y, x-z, y-z)
Electron spin quantum number
Determines which way an electron spins (up or down)
define degenerate orbitals
oribitals which share the same energy level (the three orbitals in 2p are degenerate with each other)
Equation for change in energy (delta E)
Delta E = hv
h = 6.626 x 10^-34
v = frequency
de broglies equation
lambda = h/mv
h= 6.626 x 10^-34
m = mass in kg
Lyman series (H)
UV
122 nm, 103 nm, 97 nm, 95 nm, 94 nm
90-130
Balmer series (H)
visible
656 nm, 486 nm, 434 nm, 410 nm
400-700
Paschen series (H)
infrared
1875 nm, 1282 nm, 1094 nm
1000-1900
Radio wavelengths:
more than 1 m
Microwave wavelengths
between 1 mm to 1 m
Infrared wavelengths
700 nm to 1 mm
visible wavelengths
400-700 nm
x-ray wavelengths
.01-10 nm
gamma rays
smaller than .01 nm
uv wavelengths
10-400 nm
c=
c= lambda v
