Orbitals & Electrons (Test Mon. 2/10/25) Flashcards
s orbitals
Spherical, exits singularly (in sets of 1); holds 2 electrons
p orbitals
Dumbell/figure-8 shaped, come in sets of 3, hold a total of 6 electrons.
All 3 p orbitals are equal in size and energy, only differing in orientation and space.
d orbitals
Clover shapes (most), come in sets of 5, hold a total of 10 electrons
f orbitals
More complex in shape, come in sets of 7, hold a total of 14 electrons
Orbitals
Each atom contains many orbitals that overlap. Orbitals are areas where electrons are likely to be found, not rigid structures.
Node
All orbitals contain at least one node; an area found in the atom where an electron cannot be found
Electron configuration mnemonic
Electrons are added to orbitals starting at the top and working their way down.
The coefficient tells you the energy level and indicates energy of the electrons in orbital, distance from nucleus, and size of orbital.
The exponent is the number of electrons.
The variable stands for the type of orbital.
Ground state vs excited state
Ground- lowest possible energy level
Excited- One of more electrons at a higher energy level than expected
It’s impossible to contain orbitals that don’t exist or too many electrons in a given orbital.
Aufbau orbital diagram
Start at the lowest energy level and go up
Hund orbital diagram
When orbitals occupy equal energy, one electron enters each orbital before any orbital gets a 2nd one.
Pauli Exclusion orbital diagram
One orbital can have at most 2 electrons; when 2 electrons enter the same orbital, they must have opposite rotation
Magnetism: Paramagnetic vs Diamagnetic
Paramagnetic- At least one unpaired electron is attracted to a magnet
Diamagnetic- All electrons are paired, not attracted to a magnet
Isoelectronic
Having the same number of electrons
Periods
“Rows” of periodic table, number 1-7. Elements in the same period don’t necessarily have same properties
Groups
“Columns” numbered 1-8. Elements in same group are most similar because they have the same number of valence electrons
Valence electrons
The outermost shell of electrons of an atom. Are most important because they determine how elements will bond. To tell how many an element has, go from 1 to 8 on periodic table columns, skipping the big middle transition medals.
Charge
Elements will form charges to attain electron configuration of noble gas (they have a full valence shell, usually 8 valence electrons- octet rule).
Group 1- Alkali Metals
(H, Li, Na, K, Rb, Cs, Fr)
Valence electrons- 1
Reactivity- Extremely reactive, especially with water and oxygen
Physical properties- Soft, silver, shiny, low density
Group 2- Alkali-earth metals
(Be, Mg, Ca, Sr, Ba, Ra)
Valence electrons- 2
Reactivity- Very reactive, less than Alkali metals
Physical properties- Silver, higher densities than Alkali metals
Group 3-12- Transition metals
(Whole middle area)
Valence electrons- 1 or 2
Reactivity- Slightly less reactive than Alkaline-earth metals
Physical properties- Shiny, hood conductors, higher densities and melting points, variable changes
Group 3- Inner transition metals
(Whole bottom row separated)
Valence electrons- 1 or 2
Reactivity- Slightly less reactive than Alkali-earth metals
Physical properties- Lanthanides= shiny, some used in steel alloys. Actinides= radioactive
Group 17- Halogens
(F, Cl, Br, I, At, Ts)
Valence electrons- 7
Reactivity- Extremely reactive, especially with Alkali metals
Physical properties- Poor conductors, most diatomic, most gas at room temperature
Group 18- Noble gas
(He, Ne, Ar, Kr, Xe, Rn, Og)
Valence electrons- 8
Reactivity- Nonreactive, full valence shell
Physical properties- Colorless, odorless, gas at room temperature
Periodic table trends: Increasing energy levels
Down a group= energy levels increasing
Electrons in higher energy levels are further from the nucleus and so are less attracted to it
