Chem chapter 5

bond polarity

electrons are not shared equally between atoms in a covalent bond, resulting in a partial positive charge and a partial negative charge

electronegativity

ability of an atom to attract bonding electrons to itself

bond polarity and electronegativity connection

larger EN difference = higher bond polarity

negative end of bond polarity goes towards the more EN atom

electronegativity trend

increases across a period and decreases down a group

• Fluorine is highest, francium is lowest

• noble gases not assigned EN values

• opposite of atomic size trend

pure covalent bond

difference in EN is 0- equal sharing of atoms in a bond

nonpolar covalent

EN difference between 0.1-0.4

polar covalent

EN difference is between 0.5-1.9 - unequal sharing of electrons between atoms in a bond

ionic bond EN

EN difference greater than 2

dipole moment

measure of bond polarity

• a material with a positive and negative end

• directly proportional to the size of partial charges and the distance between them

dipole moment equation

percent ionic character

the percentage of a bond’s measured dipole moment compared to what it would be if electrons were completely transferred

• indicates the degree to which the electron is transferred

percent ionic character formula

measured dipole/dipole if e- was transferred *100%

resonance purpose

used when 2+ valid Lewis structures can be drawn for the same compound

formal charge definition

an electron bookkeeping system that allows us to discriminate between alternative Lewis structure

• atoms may end up with more or fewer electrons than the valence electrons they brought in order to fill octets, resulting in a charge

resonance theory

suggests that there is some degree of delocalization of the electrons

• delocalization of charge helps stabilize the molecule

resonance structures

when there is more than 1 Lewis structure for a molecule that only differ in the position of electrons

formal charge rules

1. sum of all formal charges in a neutral molecule must equal 0

2. sum of formal charges in an ion must = the charge of the ion

3. small (or zero) formal charges on individual atoms are better than large ones

4. when formal charge cannot be avoided, negative formal charge should lie on the most electronegative atom

formal charge formula

# of valence electrons - [# of nonbonding electrons + 1/2(# of bonding electrons)]

SHORTCUT: how many valence electrons atom should have - how many it actually has

group number - (number of lines + number of dots)

exceptions to octet rule

odd electron species (free radicals), incomplete octets, and expanded octets

odd electron species (free radicals)

molecules or ions with an odd number of electrons- incomplete octets, molecules, or ions with fewer than 8 electrons

incomplete octets

elements (specifically metalloids and hydrogen) whose tendency is not to have a complete octet

expanded octets

molecules with more than 8 electrons around an atom - 3rd row nonmetals and below can do this

expanded octets and Lewis structures

when a high charge separation is detected, like a high positive charge (ex. +2) on a central atom sitting next to negative charges on terminal atoms

• take lone pair from the negative atom and double bond it to the central atom, reducing positive charge by 1 on the central atom and increasing negative charge by 1 on terminal atom

bond energy

the amount of energy in gas state it takes to break one mole of a bond in a compound

• greater bond energy = stronger/more stable bond

trends in bond energy

in general, more electrons shared = stronger bond

• must compare bonds between like atoms (i.e triple bonded carbon, double bonded carbon, and single bonded carbon)

in general, the shorter the bond, the stronger the bond

• must compare similar types of bonds (Br-F > Br-Cl > Br-Br)

bonds get weaker down a column and stronger across a period

• larger atomic radius = weaker bonds + vice versa

bond length

the distance between the nuclei of bonded atoms

• we use the average bond length for similar bonds across many compounds

trends in bond length

in general, more electrons shared = shorter bond

• must be compared between like atoms (c triple bond greater than c double bond greater than c single bond)

generally decreased across a period and increases down a column

as bonds get longer, they get weaker

structure of compounds depends on

• skeletal arrangement of atoms

• kind of bonding between atoms (ionic, polar covalent, or covalent)

• shape of molecule

molecular geometries

• describe the shape of a molecule with terms that relate to geometric figures

• have characteristic “corners” that indicate the positions of the surrounding atoms around a central atom in the center of the geometric figure

  • have characteristic angles called bond angles

valence shell electron pair repulsion theory

electron groups around the central atom will be most stable when they are as far apart as possible- the resulting geometric arrangement will allow us to predict shapes and bond angles in the molecule

electron groups

each lone pair of electrons constitutes one electron group, and each bond constitutes one electron group regardless of single, double, or triple status

linear geometry

two electron groups

bond angle 180 degrees

trigonal planar geometry

three electron groups

bond angle 120 degrees

tetrahedral geometry

four electron groups

bond angle 109.5 degrees

trigonal bipyramidal geometry

5 electron groups

positions above and below central atom are axial positions

• bond angle between axial and equatorial positions is 90

positions in the same base plane as the central atom are equatorial positions

• bond angle between equatorial positions is 120

octahedral geometry

six electron groups

figure has 8 sides- explains octa- prefix

all positions are equivalent

• bond angle 90 degrees

lone pair groups

occupy space on central atom but not “seen” as points on molecular geometry

• take up more space on central atom because their electron density is exclusively on central atom

  • affects bond angles and makes bonding-pair angles smaller than expected

relative sizes of repulsive force interactions

lone pair to lone pair > lone pair to bonding pair > bonding pair to bonding pair

two types of tetrahedral electron geometry

trigonal pyramidal and bent

trigonal pyramidal molecular geometry

4 electron groups around the central atom, one of which is a lone pair

bond angle of 107 degrees

bent molecular geometry

4 electron groups and 2 are lone pairs

bond angle of 104.5 degrees

bond angle distortion

as more lone pairs are added, bond angles get progressively smaller

3 types of trigonal bipyramidal geometry

seesaw, t-shaped, linear

seesaw molecular geometry

5 electron groups, 1 of which is a lone pair

90 and 120 degree bond angles

t-shaped molecular geometry

5 electron groups, 2 of which are lone pairs

linear molecular geometry

5 electron groups, 3 of which are lone pairs

2 types of octahedral geometry

square pyramidal and square planar

square pyramidal molecular geometry

6 electron groups, 1 of which is a lone pair

bond angles between axial and equatorial positions are less than 90 degrees

square planar molecular geometry

6 electron groups, 2 of which are lone pairs

bond angles between equatorial positions are 90 degrees

representing 3D shapes on paper

• the central atom is put in the plane of the paper

• use a straight line for atoms in the same plane of central atom

• use a solid wedge for atoms in front of the plane

• use a hatched wedge for atoms behind the plane

illustrations of molecular geometries using 3D

geometries with multiple central atoms

for multi-center molecules, each central atom has a designated shape

determining molecular polarity

1. draw Lewis structure for molecule and determine its molecular geometry

2. determine if the molecule contains polar bonds

3. determine if polar bonds add together to form a net dipole moment

molecular polarity and water solubility

water is a polar molecule, so other polar molecules and ionic compounds will dissolve in water

nonpolar molecular geometries

linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

polar molecular geometries

bent, trigonal pyramidal, seesaw, t-shaped, square pyramidal

dipole moment trend

generally, the more electrons two atoms share and the larger the atoms are, the larger the dipole moment