chapter 6 : shapes of molecules and intermolecular forces

how do lone pairs affect bond angles

lone pairs repel more strongly than bonded pairs so they push bonded pairs closer together which decreases the bond angle by approximately 2.5 degrees for each lone pair

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for a linear shapethere

carbon dioxide

180 degrees

2

2

0

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for a trigonometry planar shape

BF3

120 degrees

3

3

0

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for a tetrahedral shape

CH4

109.5 degrees

4

4

0

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for a pyramidal shape

NH3

107 degrees

3

3

1

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for a bent / non-linear shape

H2O

104.5 degrees

4

2

2

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for a triginal bipyramidal shape

PCL5

120 / 90 degrees

5

5

0

what is an example and what are the bond angles and how many electron regions, bonding regions and lone pairs there for an october deal shape

SF6

90 degrees

6

6

0

what is electronegativity

ability of an atom to attract electrons in a covalent bond

what 3 factors affect electronegativity

nuclear charge

shielding

atomic radius

how is electronegativity measured

using the pauling scale which compares elements to fluorine which is the most electronegative element

what three elements can have hydrogen bonds

fluorine oxygen and nitrogen

draw a hydrogen bond between two water molecules

what are the 3 types of intermolecular forces from strongest to weakest

hydrogen bonding

permanent dipoles

london forces

how does electronegativity change as you go across the periodic table

it increases as nuclear charge increases, shielding stays the same and atomic radius decreases

how does electronegativity change as you go down the periodic table

it decreases as atomic charge increases, shielding increases and atomic radius increases. atomic radius and shielding have a greater affect on electronegativity than atomic charge does though

what are the 4 most electronegative elements

the non-metals oxygen fluorine nitrogen chlorine

what are the 3 least electronegative elements

group 1 metals lithium sodium and potassium

how do ionic bonds form

if electronegativity is very differdnt between two covalently bonded atoms, one bonded atom will have a greater attraction to the shared pair of electrons and will pull it away from the other atom till it’s no longer shared

what is a non-polar bond

a covalent bond where the electrons are shared equally between the bonded atoms

in what 2 cases there be a non-polar bond

if the two atoms are the same

if the two atoms have the same or similar electronegativity

what is a polar covalent bond

a covalent bond where the elements have different electronegativity values

what is a dipole

a difference in charge between two atoms caused by a shift in electron density in a bond causing a separation in opposite charges

what is a permanent dipole

a dipole that doesn’t change

what is the relative repulsion of different types of bonds from weakest to strongest

bonded pair - bonded pair, bonded pair - lone pair, lone pair - lone pair

what is an intermolecular force

weak interactions between dipoles of different molecules

what are london forces

weak induced dipole-dipole bonds between all molecules whether polar or non-polar

how are induced dipoles formed (4)

the movement of electrons produce a changing dipole

an instantaneous dipole will exist (though it is constantly changing)

the instantaneous dipole will induce a dipole in another molecule

the induced dipole induces further dipoles in other molecules which then attract one another

what are 3 effects of more electrons in a molecule on london forces

larger instantaneous and induced dipoles

stronger attractive forces between molecules

greater induced dipole-dipole interactions

how does the number of electrons affect boiling point

the more electrons, the stronger the london forces so more energy is required to break the bonds to change state which means a greater boiling point

what is a simple molecular substance

simple molecules of small units containing a definite number of atoms with a definite molecular formula

what structure do simple molecular substances form as a solid

simple molecular lattices

what are two features of simple molecular lattices

molecules are held in place with weak intermolecular forces

atoms within each molecule are held together strongly by covalent bonds

what are the melting and boiling points like for simple molecular substances

low

what happens when simple molecular substances melt

weak intermolecular forces break but strong covalent bonds don’t

what happens when a simple molecular compound is added to a non-polar solvent like hexane

intermolecular forces form between the substance and solvent which weakens intermolecular forces in the simple lattice causing it to break apart and the substance to dissolve

are simple molecular substances soluble or insoluble in polar solvents and why

they are insoluble because the intermolecular forces (dipole-dipole bonds) in the polar solvents are too strong to break to dissolve

why are simple molecular substances not conductors

because there’s no charged particles that can move

why can polar substances sometimes dissolve in polar solvents

because the polar substance molecules and polar solvent molecules can attract each other and break the lattice apart

what is a hydrogen bond

a type of dipole-dipole bond between a molecule with an electronegative atom with a lone pair and a molecule with a hydrogen atom attached to an electronegative atom

what are 3 examples of electronegative atoms within each molecule a lone pair of electrons

oxygen nitrogen fluorine

how do you represent a hydrogen bond

a dashed line

why is solid ice less dense than liquid water

hydrogen bonds hold water molecules apart in an open lattice structure so water molecules are further apart in solids than liquids meaning ice is less dense

why does water have a relatively high melting and boiling point

they have hydrogen bonds as well as london forces so more energy is required to break these extra bonds which means a higher melting and boiling point as the lattice in ice needs to be broken for it to melt