GenChem Finals

MOLECULAR POLARITY

Result from imbalance distribution of electric charge in the molecule

Due to the formation of partial positive and partial negative ends

Determined by bond polarity and shape

abt the entire molecule

BOND POLARITY

Determined by the electronegativity difference of the two bonded nonmetals

Symbolized by polar arrow

Molecules that contain polar bond are not always polar

only abt the bond

electronegativity difference

mostly ionic bonds - 1.7 or higher

polar covalent - greater than 0.4 to less than 1.7

mostly covalent - 0.4 or less

nonpolar covalent - 0

ionic bonds

polar but we don’t use polar arrows and partial charges

depicting polar bonds

unequal electron sharing can be depicted by a polar arrow which points to the more electronegative element and partial charges

δ+ +→ δ-

polar covalent bond

much more common than either pure ionic or pure covalent bonds

shared electron pair isn’t shared equally, but remains closer to one atom than the other

unequal electron sharing causes the more electronegative atom of the bond to be partially negative and the less electronegative atom to be partially positive

electron density shifts to more electronegative element

nonpolar covalent bond

equal sharing of electrons

no partial charges

lewis structures of molecules

same as the lewis structure of covalent compounds

need to be able to discern the central atom, bonding grps, and non-bonding grps

converting a molecular formula → lewis structure

1. Draw the skeletal formula (central atom and surrounding atoms)

2. Get total # of valence electrons and deduct the # used to form single bonds from all of the atoms

3. Divide remaining electrons by 2 and distribute those lone pairs to the surrounding atoms first, then the central atom

4. is the central atom doesn’t have full octet, change a lone pair of a surrounding atom into another bonding pair for the central atom to form a multiple bond

5. add brackets and charges in the case of polyatomic ions

6. for acids, H atoms are never bonded to the central atom, only the oxygen atom

central atom

less electronegative atom

appears only once in the formula (?)

hydrogen can never be a central atom

single bonds

are never between 2 surrounding atoms

multiple bonds

if there aren’t enough electrons for the central atom to attain an octet, a multiple bond is present

molecules that don’t follow the octet rule

molecules w/ electron deficient atoms

expanded valence shells

molecules w/ odd electron atoms

molecules w/ electron deficient atoms

molecules that have less than 8 valence electrons

B and Be are electron deficient most times

expanded valence shells

central atom w/ more than 8 valence electrons

can be atoms from period 3 and below (nonmetals?)

molecules w/ odd electron atoms

the lone electron / radical electron makes molecules very unstable and short-lived

happen due to d-orbitals

molecular shapes

the function of a molecule depends on its shape

Valence Shell Electron Pair Repulsion Theory (VSEPR Theory) and Molecular Shape (?)

each grp of valence electrons around a central atom is located as far as possible from others to minimize repulsions

a “grp” of electrons is only the # of electrons that occupies a localized region around an atom

a single bond, double bond, triple bond, lone pair, or single electron all count as a single grp each

electron grps (2 categories) are only those around the central atom (total # = bonding grps + nonbonding grps)

for multiple central atoms, one will be specified

bonding grps

single bonds, double bonds, or triple bonds

each are counted as 1 bonding grp

non-bonding grps

lone pair (around the central atom only) and lone electron

each are counted as 1 bonding grp

electron grp arrangement

depends in the total number of electron grps

defined by both bonding and non-bonding electron grps

• linear

• trigonal planar

• tetrahedral

• trigonal bipyramidal

• octahedral

Does not always = the shape of the molecule

linear

2 electron grps around the central atom

180 degree ideal bond angle

trigonal planar

3 electron grps around the central atom

forms a triangle and lies flat on a plane

120 degrees ideal bonding angle

tetrahedral

4 electron grps around the central atom

109.5 degrees ideal bond angle

trigonal bipyramidal

5 electron grps around the central atom

triangle at the middle and 2 pyramids formed (1 top, 1 bottom)

120 degree ideal bond angle between 2 equatorial sides

90 degree ideal bond angle between an axial and equatorial side

has two different positions for electron groups, and two ideal bond angles

Equatorial-equatorial repulsions are weaker than axial-equatorial repulsions

Where possible, lone pairs in a five electron-group system occupy equatorial positions

axial

positioned vertically

equatorial

positioned at the middle (horizontal)

octahedral

6 electron grps around the central atom

forms 8 sides

90 degrees ideal bond angle always

molecular shape

3D arrangement of nuclei joined by the bonding grps

defined by the relative positions of the nuclei

classified using AX_mE_n

AX_mE_n

A - central atom

X - surrounding bonding grps

m - # of bonding grps

E - non-bonding valence electron grp

n - # of non-bonding grps

bond angle

angle formed by the nuclei of 2 surrounding atoms w/ the nucleus of the central atom

ideal and real

ideal bond angles

determined by basic geometry alone

similar electron grps around the central atom → real = ideal

real bond angles

deviate in many cases

• more bonds

• lone pair

• presence of more electronegative atom

more bonds

more bonds → more repulsion → bond angle affected

real ≠ ideal bond angle

angle of multiple bond and other bonding pair → greater than ideal

angle of 2 bonding grps → less than ideal

lone pair

freedom of lone pair to move → more repulsion

angle of lone pair and bonding grp is greater than the ideal bond angle

angle of 2 bonding grp is less than the ideal bond angle

2 lone pairs may neutralize each other

presence of more electronegative atom

attracts bonding grps to itself

possible shapes

as long as there’s no lone pair, the name of the shape = EGA name

Shapes of Linear

only 1 shape total

• Linear (AX₂)

Shapes of Trigonal planar

2 shapes total

• Trigonal planar (AX₃)

• Bent or v-shaped (AX₂E)

Shapes of Tetrahedral

3 shapes total

• Tetrahedral (AX₄)

• Trigonal Pyramidal (AX₃E)

• Bent or v-shaped (AX₂E₂)

Shapes of Trigonal Bipyramidal

4 shapes total

lone pairs are always equatorial

• Trigonal Bipyramidal (AX₅)

• Seesaw (AX₄E)

• T-shaped (AX₃E₂)

• Linear (AX₂E₃)

Shapes of Octahedral

3 shapes total

• Octahedral (AX₆)

• Square Pyramidal (AX₅E)

• Square Planar (AX₄E₂)

pi bonds (π)

single bond in a double bond

middle bond of a triple bond

sigma bond (σ)

single bonds

chemical bonds

aka intramolecular forces

Intermolecular forces

forces of attraction that exist between molecules in a compound

aka Van der Waals Forces

aka secondary bonds

responsible for physical properties of matter (eg boiling pt)

weaker than chemical bonds

illustrated w/ dotted lines

Characteristics of IMF

weaker than primary bonds (ionic & covalent bonds)

strength of IMF affects the melting & boiling point as well as solubility of substances

causes the compound to exist in certain states of matter at a given temperature

weak IMF = gases at room temperature (ex: Helium gas)

very strong IMF = solids

Polar molecule

unequal distribution of charge as a result of its polar bonds and shape

aka dipole molecule bcs it has 2 poles (partial positive and partial negative)

Nonpolar molecule

equal distribution of charge which may be due to presence of nonpolar bonds and its shape

does not contain dipole

Polar molecules ex

Diatomic molecules which contain polar bonds

Shapes of these molecules are bent, trigonal pyramidal, seesaw, T-shaped, square pyramidal

contain one or more polar bonds and the individual bond dipoles do not cancel

Nonpolar molecules ex

Diatomic molecules which contain nonpolar bonds

Shapes of these molecules may be linear, trigonal planar, trigonal bipyramidal, square planar, tetrahedral, octahedral

Molecules that may contain polar bonds but are ALL cancelled

Molecules with carbon and hydrogen atoms in their chemical formula(hydrocarbons) or molecules with long chain of hydrocarbons

Types of IMF

dipole-dipole forces

hydrogen bonding

ion dipole force

london forces

dipole-dipole forces

between polar molecules

creates a higher boiling pt

polarizability

More electrons are further from the nucleus → Nucleus has less pull on far off electrons

more electrons and More negative charge = more polarizability

hydrogen bonding

special kind of dipole-dipole forces formed when a hydrogen atom bonded with highly electronegative such as N, O & F

resulting bond is highly polar with hydrogen (δ +)

ion dipole forces

interaction between ion and a polar molecule

ionic compound + dipole

Example: sodium chloride dissolved in water

London forces

exists among non-polar molecules

aka London Dispersion or Dispersion Forces

weak force and usually found among identical molecules such as O2 , N2 and Cl2 since there is no attraction among 2 opposite charges

all interacting molecules have london dispersion force

possible as long as 2 particles have a pt of contact

dependent on the size of atom or molecule: Larger atoms or molecules = electrons→ easier to induce them into temporary dipoles

increase with increasing molecular mass

IMF strength

Ion dipole > hydrogen bond > dipole-dipole > london dispersion