Chemistry - Exam 2

(Chapters 5, 6, 7)

metals

lose electrons

non-metals

gain electrons

a single line represents…(—)

a covalent bond

electronegativity

• ability of atom to hold onto its electrons and not want to lose them

when is a partially negative charge made

δ−

• when electrons spend more time with more electronegative atom it will want to pull those electrons to it and hold onto it more

• the more electronegative atom is pulling those electrons closer to it so it is become more negative

• because more electrons = negative

electron deficient

• happens when a more electronegative atom pulls electrons close to it, becoming partially negative, and causes the other atom to be partially positive

  • this atom has been deprived and therefore electron deficient

polarized

• unequal sharing of electrons

• one more positive end and one more negative one

what causes an electron dipole

polarization

• because one side is pulling the electrons and the other is losing theirs

where does the point of the electron dipole indicate

the more electronegative atom

dipole moment

• direction of charge separation

• measures how unevenly charge is distributed in a bond or molecule

electrostatic potential diagrams

show electron density

relationship between dipole and polarity

a higher dipole equals greater polarity

the ____the difference in electronegativity, the ____the magnitude of partial positive and negative charges will be and, hence, the ____the dipole moment of the bond.

greater

why is it partial negative or positive

because they are not transferring the electrons completely but just unevenly sharing them

formal charge

• bonding electrons are shared equally and if broke each would take one

  • why you count individually and not by 2s

• sum of formal charges should equal some of overall charge on ion

smaller formal charge means…?

more stability

where should negative formal charges be

on the more electronegative element d

delocalized electrons

electrons that are not bound to a single atom or bond can switch around (think resonance)

isomer

• same molecular formula different structures

• different atomic arranges

  • the actual elements

resonance structures

• different electron arrangements

VSEPR Theory

• how to predict 3D shape of a molecule

• electrons will arrange themselves far apart as possible to minimize repulsion

• electron domains

• electrons repel each other because they are both negative and causes them to want to be as far apart as they can

electron domains

• lone pairs and bonding electrons

  • doesn’t matter if a single, double, or triple bond

    • still equals one

the larger the electron domain the more or less repulsion?

more repulsion

steric number

another name for electron domain

which electron domain has the greatest distance

tetrahedral (steric number 4)

lone pairs

• have more volume and take up more space than bonds

  • have more space to spread out because not confined to a bond and causes the other electron pairs to be pushed farther apart

• reside in equatorial positions

shorter bonds

• bond strength increases

  • more attraction to the nucleus

equatorial

• on the equator

• 120 degrees

axial

• vertical line 

• 90 degrees

the higher the electronegativity…

• the more polar the bond is because there is an increase in the separation of the charges because it is being pulled more strongly towards the more electronegative atom and leaving a bigger gap between the two

central atom has lowest or highest EN

lowest electronegativity

finding average formal charge

add up all of them and then divide the amount of possible scenarios

bond order

• the number of lines

• double bond = 2

• triple bond = 3

less formal charge…

more stable

bond order

• count all the bonds in that structure

  • then divide by the amount of bonding pairs (how many peripheral atoms you have)

    • if there is resonance structures

• average is when you divide them and then total is just counting the bonds

total formal charge v. average formal charge

total = just counting and adding them together

average = counting them and then dividing by the amount of peripheral atoms

put negative formal charge on what?

more electronegative atom

the best resonance structure is influencing the structure the…

most 

if all four are different around a tetrahedral

• fix one atom at the top

  • rotate

• should be 3 isomers

resonance structures

• same connectivity

  • same atoms are bonded to each other

• same number of valence electrons

  • only difference is location of where the valence electrons are

best resonance structure characteristics

• has lowest amounts of formal charge

• negative formal charge on more electronegative atom

why is negative formal charge on the more electronegative atom

• because formal charges mean an increase of repulsion and destabilization

  • and the more electronegative atom is stable and can take that already since it has a strong hold on its electrons

non zero dipole

polar

factors that affect polarity of a molecule

1. the geometry

2. polarity of bonds involved

if all peripheral atoms are the same…

non-polar

• because pulling equally in opposite directions of each other

lone pairs on top will create…

a polar molecule

• because it is pushing all the other atoms down to make room for itself

  • think bent shape

• more separation of charges

take average when…

equivalent resonance structures

if unequal resonance structures use…

the best one

when dipoles are in opposite directions of one another the molecule is…

nonpolar

when will an atom become partially negative/positive when it wasn’t before

• if its formal charge is more negative than the other atom

valence bond theory

• a covalent bond is the overlap of half-filled atomic orbitals that make a pair of electrons that are shared between the two bonded atoms

• and then they overlap

According to valence bond theory, a covalent bond results when…?

1. an orbital on one atom overlaps an orbital on a second atom

2. single electrons in each orbital combine to form an electron pair

hybrid orbital theory

• the number of electron domains equal the number of hybrid orbitals

what forms when two single electrons overlap

sigma bond

what type of electrons (paired or unpaired) make bonds

single

• because you covalently sharing them so you can only give one

pi bonds result from

unhybridized orbitals

lone pairs always reside in

hybridized orbitals

bonding

• interaction of atomic orbitals to form molecules

degenerate

same energy