Electron Configurations Flashcards
Electromagnetic Spectrum
The range of all types of electromagnetic radiation (light)
Speed of light
C = 3.0 x 10^8 m/s
Wavelength
The distance between two peaks of a wave, specifically electromagnetic radiation. (𝝺, measured in m)
frequency
How often a peak occurs while a wave passes a specific point. (f, measured in Hz or s^-1)
Equation for the wavelength and frequency of light
C = f * 𝝺
Energy of light
Energy = 1/𝝺
continuous spectrum
How we experience light, it appears as a continuous series of colors with no gaps.
emission spectrum
The spectrum of light produced when a gaseous element is exposed to high voltage (or heat) and electrons are energized.
absorption spectrum
The spectrum of light produced when a gaseous element that has been heated and cools, and electrons return to lower energy.
Line spectra
The emission spectra specific to an element, where only certain wavelengths are emitted.
Quantization
Electromagnetic radiation comes in discrete packets or quanta. Quanta are represented only by integers, (think of steps vs a ramp)
Photon
a quantum of electromagnetic radiation energy
Energy of a photon
E = h x f
Planck’s constant
h = 6.63 x 10^-34 j/s
Energy levels
fixed distances from the nucleus of an atom where electrons may be found
Energy of a orbital
En = -Rh (1/n^2)
Rydberg constant
Rh = 2.18 x 10^-18 j
Quantum numbers
All of the values used to describe the energy of an electron
Principal quantum number
n, the distance and energy of an orbit
Ground state
the most stable formation for an atom, where electrons are at their lowest energy level
excited state
when an electron moves up one or more orbitals
Number of electrons in n=1
2
Number of electrons in n=2
8
Number of electrons in n=3
18
Number of electrons in n=4
32
Heisenberg uncertainty principle
states that it is impossible to precisely know the location and speed of an electron simultaneously
Schrodinger’s wave function
used to describe the electrons in an atom based on a probability density function. Electron’s have a high likelihood of occupying a region in space called an atomic orbital.
Atomic orbital
An area where electrons have a high likelihood of occupying a region in space
S orbital
sphere shaped, can hold 2 electrons
p orbital
dumbbell shaped, there are 3 p orbitals per energy level, each can hold 2 electrons for a total of 6.
d orbital
flower shaped, there are 5 orbitals per energy level, each can hold 2 electrons for a total of 10.
Degenerate orbitals
Orbitals with the same energy are degenerate (same principal quantum number and orbital shape).
Orbital diagrams
Electrons represented by arrows, in orbitals represented by boxes.
Electron confirgurations
A way to represent the location of electrons in an atom, either as a diagram or written out.
Pauli exclusion principle
Electrons also must have opposite spins if they are in the same orbital so one arrow is up and the other down to represent this
Hund’s rule
every degenerate orbital in a sublevel is singly occupied and all electrons in singly occupied orbitals have the same spin
Aufbau principle
electrons are added to an atom’s lowest energy orbital first. You will only need to know up to 4p!
Condensed electron configuration
A short hand representing inner core electrons with the noble gas from the row above
Exceptions to the Aufbau rule
Transition metals will fill or half fill 3d before 4s. Cu: 4s13d10 and Cr: 4s13d5
Period
the rows of the periodic table
Group
column of the periodic table
Isoelectronic species
Atoms with identical electron configurations, will be a series of cations, an element and anions.
Metal
generally have three or fewer valence electrons, which are delocalized
Nonmetal
generally have four or more valence electrons, which are NOT delocalized
Metalliod
show hybrid metallic and non-metallic character. The touch the zig-zag line shown on the table.
