Levels are the energy levels/periods/rows, Sublevels are s, p, d, and f, Orbitals are the positions that hold the electrons (# of orbitals half of # of maximum electrons)

attraction of opposite charge is highly favorable, meaning more energy is needed to break this stability

Ionic bonds are actually covalent bonds, it's just the electronegativity difference is so much that the electron is transferred completely and not shared

while there's attraction between all electrons and both nuclei, there is also repulsion between nuclei and between electrons, meaning covalent bonds form at the most ideal balance of these forces

X: internuclear distance in picometers Y: potential energy (Joules)

Up to the ideal/favorable point where the covalent bond fully forms, potential energy decreases due to the increasing nuclei-electron attraction, but afterwards, potential energy increases because the atoms are bonded together yet experience greater repulsion of like charges

Cl-Cl will have greater internuclear distance due to more shielding, and the shielding means the valence electrons are farther from both nuclei, so less attraction means greater potential energy

single-double-triple

energy required to separate 1 mole of a certain bond

the energy required to form one mole of a solid ionic compound from its gaseous ions

lattice energy = constant * (charge in Coulombs of ion #1)(charge in Coulombs of ion #2) over internuclear distance

Charge

the y axis is number of e-, and the x axis is the binding energy that decreases as the electrons move further away from the nucleus

As the energy levels and sublevels increase towards the right, the greatest gaps are between energy levels (EX: 1 to 2) and the smaller gaps are between sublevels (EX: s to p)

shows the enthalpy changes that occur in each step of separating/forming an ionic compound

5; however, some molecules don't require all steps to occur; EX: In MgS, both atoms are already single as opposed to Na + Cl2

convert any non-gaseous atoms into gaseous state; endothermic

separate any diatomic atoms into single atoms; endothermic

ionization energy is needed to remove the electron(s) from the cation; endothermic

the anion's electron affinity energy attracts the cation's electron(s); exothermic

the oppositely-charged gaseous ions now combine to form the final solid compound; exothermic

if atoms give away multiple electrons, those electrons have different ionization energies; so atoms that accept multiple electrons have different electron affinity energies specific to each

-more electronegative atoms

-atoms that currently have a positive formal charge

-atom that is most furthest down a period starting at period 3

of expected valence electrons - actual valence electrons

the uneven sharing of electrons in molecules due to asymmetrical structure and different electronegativities

structures in which the overall compound charge and position of the atoms stays the same, but the position of some electrons can move around due to not being localized

-more stable than the different resonance structures separate

-like a mix of all the resonance structures, so its bond length is between that of single and double bonds

because they have an s, p and d sublevels, not just s and p

molecular geometries can be determined thanks to the repulsions between electrons and the greatest minimizations of those repulsions

linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

180˚

120˚

109.5˚; to minimize repulsions, tetrahedral adopts a multi-plane structure which causes wider angles than what would've been in a planar structure

-120˚ between the equatorial

-90˚ between the axial and equatorial

90˚

lone pairs take up more space and thus have a wider range of repulsion, pushing the other electron pairs closer together

-trigonal planar: all electron pairs are bonded

-bent: 1 nonbonded electron pair

-tetrahedral: all bonded

-trigonal planar: 1 unbonded

-bent: 2 unbonded

the equatorial positions

-trigonal bipyramidal: all bonded

-seesaw: 1 unbonded

-t-shaped: 2 unbonded

-linear: 3 unbonded

-octahedral: all bonded

-square pyramidal: 1 unbonded

-square planar: 2 unbonded

-linear -trigonal planar -square planar -trigonal bipyramidal -octahedral -tetrahedral

as long as the atoms in the equatorial positions are the same and the atoms in the axial positions are the same, then nonpolar

if atoms on both side of central atom are the same, then nonpolar because of equal dipoles

perform this for every plane of the molecule

for each bond, draw the dipole arrow in the appropriate direction

assemble the arrows into a cycle; if the cycle can repeat endlessly, it is nonpolar

when an atom is bonding, the s and p orbitals mix together to form hybridized orbitals which correlate with an electron geometry

if atom does not have p orbitals; Ex: H

x and z: horizontal and perpendicular of each other y: vertical

90˚

sphere

pair of lobes (like an 8)

the different variations of the mixing of s and p orbitals (of equal energies) which all merge into one final hybridized structure

3; 3; 4

tetrahedral

2

some hybridized orbitals will not be bonded and will just be holding unbonded electrons

1s + any 2 of the 3 p orbitals (doesn't matter which 2 axes you choose)

3; trigonal planar

a type of covalent bond in which the orbitals of different atoms overlap

non-hybridized

pi bonds form perpendicular to sigma bonds and if the non-hybridized p-orbitals are aligned in the same direction

sigma bond

a sigma and pi bond

a sigma bond and 2 pi bonds

1 out of 3

2; linear