Unit 2: Molecular and Ionic Structure and Properties

"I think I did well on that test" I got a 76 :/

intramolecular forces

intramolecular forces

attractions within a molecule

intermolecular forces

attractions between molecules

ionic bonding

transfer of electrons from a metal to a nonmetal

coulomb’s law

an increase in charge or decrease in size = increased bond energy

covalent bonding

sharing of electrons between non metals

polar molecule

electrons are not shared equally; large change in electronegativity

nonpolar molecule

electrons shared equally; small change in electronegativity

repulsion; the molecules are too close to form a bond

where the molecules are actually bonded together (internuclear distance at this point = bond length)

where the molecules are too far apart to form a bond

bond energy

bond energy

strongest bond

lowest potential energy = ?

smaller

bond length is shorter for _______ molecules

increased bond energy, decreased bond length

increased bond order = ?

bond order

the number of bonding pairs of electrons between two atoms

from sp to sp³

sp³

what is the hybridization of oxygen?

pi bonds

formed by the lateral overlap of two atomic orbitals; e-shared around atoms

sigma bonds

a result of the head-to-head overlapping of atomic orbitals; e- shared between atoms

8 sigma, 3 pi

sigma bonds and pi bonds?

2 sp³ orbitals

yes, from sp to sp²

bond polarity

difference in electronegativity of two elements

more electronegative atom

dipole arrow points towards ______

increased bond energy

delocalized electrons = ____________

single bond

sigma bond

double bond

sigma bond + pi bond

triple bond

sigma bond + 2 pi bonds

Octet rule

All atoms end up with 8 electrons around them (except for hydrogen and boron)

Steps to draw Lewis Structures

1. count the valence electrons of all the atoms

2. determine the central atom

3. remove an electron from each atom for a bond (S—O bond removes one electron from S and one from O)

4. add the remaining electrons as lone pairs to create an octet around each atom

   1. if octet has not been achieved add multiple bonds

   2. if all atoms have achieved an octet and not all valence electrons have been assigned, add the remaining electrons to the central atom

most electronegative

the central atom is the ______

expanded octet

when a atom has more than 8 electrons

the electrons of pi bonds are delocalized

in resonance structures

formal charge

a hypothetical charge assigned to an atom in a molecule. It's based on the assumption that electrons in all chemical bonds are shared equally between atoms, regardless of their relative electronegativity.

= valence electrons - (lone electrons + ½ bond electrons)

better structures

molecules with lower formal charges are _____

negative

a more electronegative atom will have a more _______ formal charge

sign

adjacent atoms should not have formal charges with the same ________

charge on chemical species

sum of formal charges =

closer to 0

a more favorable Lewis Structure will have a formal charge ___________

0

6 valence electrons - (4 unbonded electrons + 4 bond electrons / 2) = 0

formal charge of oxygen in H2O

0

6 valence electrons - (4 unbonded electrons + 4 bond electrons/2) = 0

formal charge of oxygen in CO2

0

4 valence electrons - (0 unbonded electrons + 8 bond electrons/2) = 0

formal charge of carbon in CO2

0

1 valence electron - (0 unbonded electrons + 1 bond electron/2) = 0

formal charge of hydrogen in H2O

-1

6 valence electrons - (6 unbonded electrons + 2 bond electrons/2) = -1

formal charge of oxygen furthest to the left in NO3

-1

charge of resonance structure

the one on the left, because it has a charge of 0

which resonance structure is more favorable?

negative

bonds and lone pairs are ________

repel

electrons (-) ____ each other

minimize repulsion

because bonds and lone pairs are negative, they arrange themselves to ________________

steric number

number of electron domains

2

sp steric number

3

sp² steric number

4

sp³ steric number

Linear

180

Linear bond angle

2

Linear (# of electron domains)

Linear

Linear Basic Geometry

trigonal planar

120

trigonal planar bond angle

3

trigonal planar # of electron domains

trigonal planar

Trigonal planar basic geometry

bent

< 120

bent bond angle

3

bent (# of electron domains)

trigonal planar

bent basic geometry

Tetrahedral

109

tetrahedral bond angle

4

Tetrahedral (# of electron domains)

tetrahedral

tetrahedral basic geometry

trigonal pyramidal

<109

trigonal pyramidal bond angle

4

trigonal pyramidal (# of electron domains)

tetrahedral

trigonal pyramidal basic geometry

bent (tetrahedral)

<< 109

bent (tetrahedral) bond angle

4

bent (tetrahedral) (# of electron domains)

tetrahedral

bent (tetrahedral) basic geometry

trigonal bipyramidal

120, 90

trigonal bipyramidal bond angle

Sawhorse or Seesaw (trigonal bipyramidal)

< 90, < 120

Sawhorse or Seesaw (trigonal bipyramidal) bond angle

T-shape (trigonal bipyramidal)

< 90

t-shape (trigonal bipyramidal) bond angle

linear (trigonal bipyramidal)

180

linear (trigonal bipyramidal) bond angle

octahedral

90

octahedral bond angle

square pyramidal

< 90, < 90

square pyramidal bond angle

square planar

90

square planar bond angle

T-shape (octahedral)

< 90

T shape (octahedral) bond angle

linear (octahedral)

180

linear octahedral bond angle

polar molecule

a molecule with an uneven distribution of charges; asymmetrical; bonds don’t cancel out; lone pairs of electrons disrupt electrical symmetry

nonpolar molecule

a compound with an even distribution of charges; symmetrical; identical atoms

polar; the lone pairs on the oxygen atom make the molecule asymmetrical

polar or nonpolar?

nonpolar, because it has an even distribution of electron density, and its dipoles cancel each other out. 

polar or nonpolar?

polar

any molecule with lone pairs of electrons on the central atom