Chem 105 Midterm 2

chemical bond

chemical bond

electrostatic attraction holding two atoms together; two nuclei attracted to the electrons between

forming bonds

releases energy

breaking bonds

requires energy

ionic bonds

complete transfer of electrons; metal and non metal; salt

covalent bond

electrons are shared between atoms; nonmetals

electronegativity

ability for an atom to steal an electron from another when bonding; large number indicates strength

Δx<2

covalent

Δx>2

ionic

energy

product of two molecules divided by their distance

crystal lattice

ordered 3D array of particles; high energy determined by higher charge, then smaller size

nonpolar covalent

equal sharing of electrons

polar covalent

unequal sharing of electrons

bond dipole

separation of electrical charge created when atoms with different electronegativities form a covalent bond

dipole moment

quantifies extent of charge separation; measured by applied electric field, increases as dipole increases

polar molecule

molecule with net dipole; asymmetric charge distribution - dipoles don't cancel

symmetric stretch

infrared inactive

asymmetric stretch

infrared active

bending mode

infrared active

metallic bond

communal sharing of electrons known as a "sea of electrons" - allows for malleability and the absorption of excess energy

binary molecular compound name

first nonmetal: #prefix+element name second nonmetal: #prefix+element name+-ide suffix

main group metal ionic name

cation charge implied - name of element; anion name is element + -ide

transition metal ionic name

roman numeral indicates charge of cation

polyatomic ion name

covalent bonds with charge; named cation + anion

hydrated salt

A salt that contains water of crystallization; name ionic salt and add [prefix] hydrate

hygroscopic salts

absorb water from air; forms desiccants, filters, etc

anhydrous salts

the salt crystals without water molecules in the compound

binary acid name

hydro(root)ic acid

oxoacid name

-ate turns to -ic -ite turns to -ous

lewis structures

prediction of valence electrons divided between bonding electrons and lone pairs

octet rule

atoms tend to lose, gain or share electrons in order to acquire a full set of eight valence electrons

double bond

covalent bond in which two pairs of electrons are shared between two atoms

triple bond

covalent bond in which two atoms share three pairs of electrons

bond order

the number of shared electron pairs between two atoms

bond length

the average distance between the nuclei of two bonded atoms; shorter = stronger, higher order bond; longer = longer, lower order bond

resonance structures

multiple arrangements of electrons; reality is an "average" of all structures

resonance rules

rearranged electrons not atoms, same net charge; no more structures than bonding sites

delocalization

spreading of electrons over more than one bond - greater spread (low kinetic) closer to the nucleus (low potential) reduces energy

formal charge

of valence electrons - ( # dots + # lines)

expanded valence shell

atoms in third row or below can accommodate more than eight electrons

VESPR theory

best arrangement of atoms is one that minimizes repulsion - keeps electrons as far away from each other as possible

steric number

number of atoms bonded to central atom + number of lone pairs on central atom

electron geometry

the geometrical arrangement of the electron groups

molecular geometry

the arrangement of bonded atoms, differs from electron geometry when lone pairs are present

linear

sn=2, 180

trigonal planar

sn=3, 120

tetrahedral

sn=4, 109.5

trigonal bipyramidal

sn=5, 90/120

octahedral

sn=6, 90

bond angle

angle between nuclei

bent

sn=3, 117, 2 bonds, 1 lone pair

trigonal pyramidal

sn=4, 107, 3 bonds, 1 lone pair

bent 2

sn=4, 104.5, 2 bonds, 2 lone pairs

seesaw

sn=5, <90/<120, 4 bonds, 1 lone pair

t-shaped

sn=5, <90, 3 bonds, 2 lone pairs

square pyramidal

sn=6, 90, 5 bonds, 1 lone

square planar

sn=6, <90, 4 bonds, 2 lone

functional group

subunit imparts characteristic behavior

alkanes

single bond between carbons

alkene

double bond between carbons

alkyne

triple bond between carbons

dark wedge

bond in front of plane

dashed wedge

bond behind the plane

line

parallel to plane

saturated hydrocarbon

max number of hydrogens per carbon in chain

unsaturated hydrocarbon

less than max number of hydrogens per carbon in the chain

isomers

same chemical formula, different structure

constitutional isomer

same molecular formula, different connectivity

geometric isomer (stereoisomer)

same covalent arrangements but differ in spatial arrangements

cis (z) isomer

two like groups on same side of c=c

trans (E) isomer

two like groups on opposite sides of c=c

conformers

different rotations about C-C bonds

enantiomers

isomers that are mirror images of each other

stereocenter

an atom at which the interchange of two groups produces a stereoisomer; carbon attached to 4 nonequivalent atoms/groups

amino acid

amine; carboxylic acid with R group in middle - chiral if R is not hydrogen (glycine)

valence bond theory

all bonds from from overlap of two half filled valence orbitals, bond = increased density

VBT advantages

applicable, explains shapes

VBT disadvantages

valence electrons belong to atom; doesn't explain magnetism, absorption/emission, or bond energy values

sigma bond

head-head overlap producing a single bond with electron density along internuclear axis, stronger bond that lowers energy

pi bond

side-side overlap of p orbitals with electron density above/below axis; planar node through bond axis; allows for multiple bonds to be formed

molecular orbital theory

wave functions of all electrons change when atoms approach each other, electrons do not belong to any one atom, orbitals extend over entire molecule

bonding orbital

a molecular orbital that can be occupied by two electrons of a covalent bond

antibonding orbital

a molecular orbital that is higher in energy than any of the atomic orbitals from which it was formed

pi bond node

one node parallel to atomic bond

pi* antibond node

two nodes perpendicular to atomic bond

proteins

amino acid structure is primary structure

peptide bond

carboxylic acid reacts with amine end - hydrolysis creates water

alpha helix

carbonyl O with H-N on amino acid four structures away

benzene

ring of six carbons with six molecular orbitals from combining PZ atomic orbitals

bands

many energy levels closely spaced together

metallic bonding

behaves as giant molecule, #MO as #atoms-bands

valence band

partially filled/filled band

conduction band

molecular orbitals higher in energy and distinctly separated from occupied valence band

metal

valence and conduction bands overlap, allowing electrons to easily delocalize

semiconductor

valence band full with small gap to conduction band; as group number increases, band gap increases

doping

chemically altering band gap by replacing atoms to increase conductivity

n-type doping

electron rich, populating conduction band

p-type doping

electron poor, less electrons in valence create "holes" that allow for movement of electrons

kinetic molecular theory

assumes gas molecules have insignificant volumes compared to their container volume, moves randomly and constantly, collide elastically, no intermolecular forces and that average kinetic energy is proportional to temperature

temperature

average energy of motion

effusion

gas escapes through a tiny opening, heavier gases diffuse more slowly