chp 12 and 15 Flashcards
importance of shape of molecule
affects vapour pressure, melting point, boiling point and solubility. shape determines how it interacts with other molecules
VSEPR
uses knowledge of valence shell to predict shape
based on principle that negatively charged electron pairs in valence shell repel each other. meaning that these electrons are arranged as far away from each other as possible
electron pairs in the single covalent bonds repel each other
lone pairs are treated the same as electron pairs in covalent bonds because they still take up space int he atom
double and triple bonds are treated as a single bond
tetrahedral (methane)
electron pairs in the single covalent bonds repel each other
repulsion forces the bonds as far apart from each other as possible, so hydrogens are arranged in a tetrahedral shape at an angle of 109.5º
pyramidal (ammonia)
lone pairs are treated the same as electron pairs in covalent bonds because they still take up space int he atom
3 hydrogen atoms for a pyramidal shape with the nitrogen. the ,one pair occupies more space because it isn’t covalently bonded to a and is therefore closer to the nitrogen. this means that the 3 covalent bonds are pushed closer together making the angle less than 109.5º
trigonal planar (water)
2 lone pairs and 2 covalent bonds
4 electron pairs repel each other and 2 hydrogen atoms for a v shape with the oxygen. because of the two lone pairs the 2 covalent bonds are pushed closer together
angle around central atom less than 109.5º
120º
linear (HF)
fluorine has 3 lone pairs and 1 covalent bond
four pairs form a tetrahedral arrangement due to mutual repulsion
results in a linear model
180º
BH3
BeH2
3 electron pairs, no lone pairs
trigonal planar
2 electron pairs, no lone pairs
linear
strength of intermolecular forces
covalent molecular forces have lower melting and boiling points because forces BETWEEN molecules are weaker. these forces are disturbed when covalent substance change state
vapour pressure
pressure that gaseous molecules exert on the closed container walls when the rates of evaporation and condensation become equal
liquids with stronger intermolecular forces have lower vapour pressure molecules are held together tightly making it harder for them to escape from the surface of the liquid
increasing temp increase vapour pressure. molecules have a higher average kinetic energy that enables them to overcome intermolecular forces. as temperature increases, more particles have enough energy to escape from liquid to gas. increasing pressure
boiling point of a liquid
temperature at which the liquids vapour pressure reaches the atmospheric pressure of the surroundings
electronegativity and polarity
in intermolecular forces, the electrostatic attraction is between positive and negative charges in the molecules
charges result because of uneven electron distribution within molecules
electronegativity is the tendency of an atom in a covalent bond to attract electrons
in a covalent bond atoms are competing for electrons being shared between them
non polar diatomic molecules
when atoms have identical electronegativities, then electrons are shared equally between them. non polar because there is no charges on either end
even distribution of valence electrons between atoms
if electronegativities between atoms are very similar, then it can be considered non-polar
electron density
measure of the probability of an electron being present at a particular location within an atom
polar diatomic molecules
2 atoms with different electronegativities, electrons stay closer to the most electronegative atom because they are more strongly attracted to that atom
polar: imbalanced electron distribution. has two oppositely charged poles
all diatomic with different atoms are polar to some extent, greater difference in electronegativities the greater the polarity
delta
partial charge, adding up partial charges = 0 because charge are due to uneven sharing of electrons, so molecules are still neutral
negative: has larger share of electrons, excess of electrons
positive: lost some of its share of electrons
dipole
separation of positive and negative charges
polyatomic molecules polarity
symmetrical balanced dipoles= non-polar
methane: carbon is more electronegative but individual dipoles cancel each other out. results in a molecule with no overall dipole
in non-symmetrical molecules, individual dipoles don’t cancel each other, creating a net dipole making the molecule polar
dipole dipole forces
only in polar molecules
results from attraction between positive and negative ends of polar molecules
weak since partial charge are small
more polar the molecule, the stronger the force
stronger force the higher the melting and boiling points because force bond molecules together in solid or liquid. higher energy is required to to break stronger bonds
hydrogen bonding
only occurs in molecules in which a hydrogen atom is covalently bonded to a N,O or F
NOF are small and very electronegative so they strongly attract electrons in a covalent bond. creates a large partial positive charge on hydrogen atom, which is attracted to lone pairs of electrons in neighbouring NOF atoms
hydrogen in covalent bond with NOF is is exposed due to lack of electrons and is highly attracted to lone pairs of NOF in other molecules
strong intermolecular bond, x10 than dipole dipole, but 1/10 of ionic
dispersion forces
forces of attraction in non polar substances (existence of intermolecular forces, without them nothing will hold molecules together and they would be gases)
caused by temporary (instantaneous) dipoles in molecules that result in random movement of the electrons surrounding molecule
strength increases as size of molecules increase. larger molecules have more electrons so it is easier to create temporary dipoles. larger molecules have higher melting and boiling points
molecules that form long chains have higher dispersion forces, more contact area to interact with it neighbouring molecules to form stronger forces
how to increase polarity of a molecule
larger difference in electronegativities
high degree of asymmetry in the shape of the molecule causing an imbalance in the bond dipole
requirement of hydrogen bonding
- hydrogen atom covalent bonded to an NOF atom
- lone pair of electron of NOF atom of neighbouring molecules
H bonds are caused by NOF’s small radius and high electronegativities and single electron of hydrogen atom
larger radii reduce concentration of negative charge around them.
hydrogen bonding in water
ice floats because it is less dense
in ice water molecules are held in a crystal lattice. in which molecules are held further part then in liquid water
when ice melts crystal lattice collapses and molecules pack together more tightly
crystal has greater volume, 4 hydrogen bonds to 4 neighbouring molecules
electrostatic attraction
strong in water because of large electronegativity difference meaning larger partial charge
water properties
can dissolve many substances easily high specific heat capacity high latent heat of vaporisation expands on freezing react with CO2 to make glucose and oxygen
water shape
lone pairs of water have greater repulsion than electron pairs in covalent bonds so they push the hydrogen atoms closer together
going down hydrides (except water)
melting and boiling points increase
Intermolecular forces are stronger (dispersion)
dispersion forces increase with strength as mass increases
water vs hydrides
hydrogen bonds in water give it a much higher boiling and melting point
high surface tension of water
water molecules at the surface are not completely surrounded by other water molecules, so they form hydrogen bonds everywhere BUT above.
molecules in body of water have no net force but surface level molecules have a downwards force. water molecules at surface are so attracted to each other by strong hydrogen bonds
surface tension meaning
measure of the resistance of a liquid to increasing its surface area
heat capacity of water
measure of a substances capacity to absorb and store heat energy, how much energy a substance absorbs as temp increases
SHC: measure of how much energy is required to increase temp by 1º of 1 gram of substance
SHC of water is relatively high because of hydrogen binds. these bonds can absorbs and store large amount of heat energy before they break
latent heat
when a something reaches its boiling/melting point the temp remains constant until all of substance changes state even though further energy is being absorbed
energy still being absorbed while state is changing is latent heat
energy absorbed by a fixed amount of substance as it changes state at its melting (fusion)/boiling (vaporisation)point
depends on strength of intermolecular forces
LHF of water is 6.0KJ/mol. energy required to disrupt ice lattice by breaking some hydrogen bonds
LHV of water is 44KJ/mol.
dissolution
particles of solute separate
particles of solvent separate
particles of solute and solvent are attracted to each other
forces:
holding solute, holding solvent (H Bonds), form between solute and solvent
like dissolves like
H bonds between water i such stronger that dispersion forces in non polar. attraction between water molecules cannot be overcome and water doesn’t separate to form a solution
water dissolving covalent
molecular compounds that form H bonds:
strength of intermolecular forces are similar, so they readily interact with each other
the more polar, the more likely it will dissolve
molecular compounds that ionise:
highly polar molecular compounds form ions when dissolved in water
polar covalent bonds break producing ions, covalent bond forms between ion and water molecule (makes hydronium in HCl), ion dipole attraction between ions and water
water dissolving ions
ion dipole: attraction between ion and polar molecule
if ion dipole attraction are strong enough, the water molecules can pull out ions from crystal lattice and into the solution
ionic compound–>hydrated ions (dissociation)
energy required to separate ions from lattice must be less than energy released when ions are hydrated
solubility
maximum amount of solute that will dissolve in a given amount of solvent
increasing temp increases solubility, at higher temps solute and solvent have more energy to overcome forces of attraction that hold particles together in the solid
saturated: no more solute can be dissolved at a certain temp
unsaturated: contains less solute needed to make solution saturated
supersaturated: contains more dissolved solute than a saturated solution
crystallisation
when an unsaturated solution becomes saturated and crystals form.
1. cooling: may reduce the solubility of a dissolved solute to the point where not all of the substance available is soluble. slow cooling = larger crystals
2. evaporating the solvent from the solution
falser rate of evaporation = smaller crystals
gases in water
Solubility decreases with rising temp
deoxygenates water so water must be cooled before released into environment
