AP®︎ Chemistry: Unit 2 Flashcards

Increases across a period and decreases down a group because of charge and distance effects.

Outermost electrons that participate in bonding and determine many chemical properties.

Electrons are shared nearly equally between atoms with similar electronegativities.

Electrons are shared unequally between atoms with different electronegativities.

Delta positive and delta negative charges caused by unequal electron distribution in a bond.

A separation of charge in a bond caused by unequal electron sharing.

Polar bonds have some ionic character, so bonding ranges continuously from covalent to ionic.

Electrostatic attraction between cations and anions, usually formed after electron transfer.

Conductive, malleable, ductile, and lustrous because mobile electrons move through the solid.

Metal plus nonmetal is usually ionic; two nonmetals are usually covalent; metals are metallic.

About 0 to 0.4 is nonpolar covalent, 0.4 to 1.7 polar covalent, above 1.7 often ionic.

Observed properties classify bonding more reliably than electronegativity difference alone.

A bond treated as effectively nonpolar because the electronegativity difference is very small.

Electrostatic attraction increases with larger charges and shorter distances between oppositely charged particles or ions.

Cations are positively charged ions and anions are negatively charged ions, typically formed by electron loss and gain.

A curve showing how energy changes as bonded atoms move closer together or farther apart.

The internuclear distance where potential energy is lowest and the bonded atoms are most stable.

The energy required to separate bonded atoms from their equilibrium distance to infinite separation.

Stronger bonds have larger bond energies and lower minimum potential energies than weaker bonds.

Very close atoms repel strongly, intermediate distances attract and bond, and far-apart atoms have near-zero interaction.

The distance between the nuclei of two bonded or interacting atoms.

Electrostatic attraction between ions is stronger for larger charges and shorter distances, so smaller and more highly charged ions interact more strongly.

A network of alternating positive and negative ions, not separate discrete molecules.

A repeating three-dimensional arrangement of ions in an ionic solid that maximizes attractions and minimizes repulsions.

A solid made of cations and anions held together by strong electrostatic attractions in a crystal lattice; ionic compounds are typically brittle, high-melting solids that conduct electricity when molten or dissolved.

A mixture of elements, at least one a metal, that forms a metallic solid with altered properties.

An interstitial alloy of iron and carbon in which carbon occupies spaces in the iron lattice.

A substitutional alloy, usually copper and zinc, in which zinc replaces some copper atoms.

Small atoms fill gaps in one type, while similar-sized atoms replace lattice atoms in the other.

Electrostatic attraction between metal cations and mobile, delocalized valence electrons spread throughout the solid.

A drawing that shows valence electrons, lone pairs, and bonding pairs in a molecule or ion.

A model where electrons are assigned to specific atoms or shared between specific atoms.

Lone pairs stay on one atom; bonding pairs are shared between atoms in a bond.

Most atoms seek eight electrons; H has 2, Be 4, B 6, and some larger atoms exceed 8.

A central atom in period 3 or higher can hold more than eight valence electrons.

Double and triple bonds form when atoms must share more than one electron pair.

Count valence electrons, choose a central atom, add single bonds, complete octets, then adjust bonds or lone pairs.

Add all valence electrons from each atom, adjusting for ionic charge by adding or subtracting electrons.

Usually place the least electronegative atom in the center, but hydrogen is never central.

Show transferred electrons, then place each ion in brackets with its charge outside.

Draw the covalent structure, adjust electrons for charge, and enclose the ion in brackets.

The actual structure is the average of all resonance structures, not any one drawing.

Electrons spread over multiple atoms rather than confined to one bond or atom.

Average bond order equals total bond lines divided by equivalent bond positions.

Assigned charge on an atom assuming bonding electrons are shared equally.

Valence electrons minus lone electrons minus number of bonds around the atom.

Prefer full octets, minimal formal charges, and negative charge on more electronegative atoms.

The total of all assigned charges must equal the molecule's or ion's overall charge.

Equivalent forms contribute equally; nonequivalent forms are judged by octets and formal charges.

A double-headed arrow between resonance structures shows alternative representations of one species.

They can fail for odd-electron species and cannot fully show delocalization or fractional bonds.

Two or more equivalent Lewis diagrams represent one delocalized electron arrangement.

The most stable valid structure has full octets and the most favorable formal charges.

Molecules or ions with an odd total number of valence electrons cannot give every atom a full octet.

Electron pairs arrange around a central atom to minimize Coulombic repulsion.

Draw the Lewis structure, count electron domains, find electron geometry, then ignore lone pairs for shape.

Regions of electron density around a central atom, including bonds and lone pairs.

A formula where A is central atom, X is bonded atoms, and E is lone pairs.

Electron geometry includes all domains; molecular geometry describes only atom positions.

Linear 180°, trigonal planar 120°, bent <120° or <109.5°, tetrahedral 109.5°, trigonal pyramidal ~107°, trigonal bipyramidal 90°/120°/180°, seesaw distorted, T-shaped ~90°, octahedral 90°, square pyramidal <90°, square planar 90°.

Nonbonding pairs repel more strongly than bonding pairs and compress bond angles.

Six electron domains arranged symmetrically with 90° angles between adjacent positions.

Mixing of valence orbitals to form equivalent orbitals that match electron-domain geometry.

sp gives linear 180°, sp2 gives trigonal planar 120°, sp3 gives tetrahedral 109.5°.

Sigma bonds form by head-on overlap; pi bonds form by sideways overlap of parallel p orbitals.

Head-on overlap is stronger than sideways overlap, so sigma bonds have greater bond energy.

Delocalized bonding can make multiple bonds equivalent, giving the same intermediate bond length.

Count one sigma in every bond, then add one pi for each double and two for each triple.

Five electron domains form three equatorial and two axial positions, with lone pairs equatorial first.

A molecule is polar when polar bond dipoles do not cancel because of its geometry.

Pi bonds restrict rotation, allowing cis-trans isomers with different spatial arrangements across double bonds.

Unequal electron sharing in a bond caused by a difference in electronegativity.

Bond order is the number of shared electron pairs between atoms; higher bond order gives shorter, stronger bonds, while lower bond order gives longer, weaker bonds.

Single bonds involve 2 shared electrons and are 1σ, longest, and weakest; double bonds involve 4 shared electrons and are 1σ + 1π, intermediate in length and strength; triple bonds involve 6 shared electrons and are 1σ + 2π, shortest and strongest.

For similar bond orders, larger bonded atoms generally form longer covalent bonds because their nuclei are farther apart.