Atomic Structure, Periodicity & Bonding

Atomic Number (Z)

number of protons in the nucleus → determines identity of the element

Mass Number (A)

number of protons and neutrons

Isotopes

atoms of the same element with same number of protons but different number of neutrons

orbitals

regions where there is a high probability of finding electrons

s-Orbitals

1 orbital per energy level, spherical shape, holds 2 electrons

p-Orbitals

appear from n=2 onward, 3 orbitals (px,py,pz), dumbbell shaped, hold 6 electrons total

d-Orbitals

appear from n=3 onward, 5 orbitals, hold 10 electrons total

f-Orbitals

appear from n=4 onward, 7 orbitals, hold 14 electrons total

Aufbau principle

electrons occupy the lowest-energy orbitals first

Pauli Principle

maximum of 2 electrons per orbital and electrons must have different spin

Hund`s Rule

electrons fill equal-energy orbitals singly before pairing → minimizes electron repulsion

Atomic Radius across a Period

decreases, more protons increase nuclear attraction

Atomic Radius down a Group

increases, additional electron shells increase size

First Ionization Energy

energy required to remove one mole of electrons from one mole of gaseous atoms

Ionization Energy across a Period

increases because of stronger nuclear attraction

Ionization Energy down a Group

decreases because of more shielding and larger distance from nucleus

Electronegativity

Ability of an atom to attract bonding electrons

Electronegativity across a Period

increases because number of protons increases → nuclear charge becomes stronger

Electronegativity down a Group

decreases because more shells are added → atomic radius and shielding increases

Isoelectronic Species

Species with the same number of electrons, radius decreases as proton number increases

s-Block

groups 1 and 2, very reactive metals, form +1/+2 ions

p-Block

groups 13-18, metals, metalloids, non-metals

d-Block

group 3-12, transition metals, form at least one ion with a partially filled d-subshell

properties of transition metals (d-Block)

variable oxidation states, colored compounds, catalytic activity, complex ion formation

Binary Ionic Compounds

metal + non-metal ending in “-ide”

roman numerals with transition metals

roman numerals show oxidation state because transition metals can have different charges

Transition Metal Ion Formation

transition metals lose 4s electrons before 3d electrons because unfilled s orbitals need lower energy but filled they are higher energy than d orbitals

Ionic Bonding

Electrostatic attraction between oppositely charged ions, occurs between metals and non-metals

Properties of Ionic Compounds

high melting and boiling points (because of strong electrostatic attraction), conduct electricity when molten or aqueous (in a solid compund ions are fixed in position → do not conduct elecricity), brittle crystalline structures (layes of ions can shift so ions with same charge may line up next to each other → same charges strongly repel → repulsion causes crystal to crack or shatter)

Lattice Enthalpy

energy released when gaseous ions come togehter to form an ionic solid, bonding is stronger with higher ionic charge (because attraction is stronger) and smaller ionic radius (also stronger attraction because ions can get closer together)

Covalent Bonding

electrostatic attraction between positively charged nuclei and shared electrons because non-metals dont wanna lose electrons

Bond Polarity

polar bonds form when atoms have different electronegativities → one atom attracts electrons more → one side is slightly positive and the other slightly negative

Metallic Bonding

positive metal ions in a sea of delocalized electrons → explains electrical conductivity, thermal conductivity, malleability, ductility

single bond

weakest but longest bond, 1 shared electron pair

double bond

stronger and shorter than single bonds, 2 electron pairs are shared

triple bond

strongest and shortest bond, 3 electron pairs are shared, high electron density between the nuclei

Alloys

mixtures of metals or metal+non-metal

different atom sizes distort the lattice and prevent layers from sliding easily → increased strength

VSEPR Theory

=Valence Shell Electron Pair Repulsion Theory states that electron pairs repel each other and molecules arrange in away to minimize their repulsion

Bond pair

bonding electron pair → shared by 2 atoms, spread out between the nuclei, occupies less space

Lone pair

belongs to only one atom → stays closer to central atom, has higher electron density, occupies more space → repel other electron pairs more strongly

Order of repulsion

• lone pair - lone pair repulsion is strongest

• lone pair - bond pair repulsion is intermediate

• bond pair - bond pair repulsion is weakest

lone pairs repel bonding pairs…

because lone pairs occupy more space and push bond pairs stronger together

Intermolecular Forces

forces of attraction between molecules, weaker than covalent bonds or ionic bonds because they do not involve full electron sharing or full ionic attraction between atoms → still determine how strongly molecules attract each other

physical properties affected by intermolecular forces

boiling point, melting point, solubility, viscosity

London Dispersion Forces (LDF)

weakest imf, caused by temporary dipoles (through electron movement), present in all molecules

strength increases with number of electrons, molar mass and surface area

Dipole-Dipole Forces

intermolecular attractions between polar molecules (because polar molecules have a slightly positive and a slightly negative side)

Hydrogen Bonding

strong type of dipole-dipole attraction, occurs only when Hydrogen bonds directly to Nitrogen, Oxygen or Fluorine

Ion-Dipole Forces

strong attractions between ions and the partial charges of polar molecules, help ionic compounds dissolve in polar solvents like water

solubility rule

“like dissolves in like” → polar substances dissolve in polar solvents, non-polar substances dissolve in non-polar solvents

Polyatomic Ions

one Ion made of multiple atoms → treated as one (always stay together)