Intermolecular forces, liquids

Difference in state of matter is due too..?

distance between particles

• solids locked in place

• slight disorder, particles move (slower)

• gases - total disorder , much empty space

state of substance depends on…?

• kinetic energy of particles ( temp dependent) - higher KE → faster → particles overcome attractive forces ( keeps particles separate)

• The strength of attraction between particles

two sides of tug a war kinda 

intermolecular force s

attraction between 2 separate molecules

• controls physical properties

intramolecular forces

bonds between atoms in a molecule ( intra=within molecule)

intramolecular is stronger than intermolecular forces

another name for intermolecular forces

van der waal forces

ions vs dipoles

ion = charge

dipole = partial charge ( overall atom has neutral charge)

ion-dipole forces 

attraction between an ion (Na+) and a polar molecule ( a molecule with a charge of 0, but uneven electron distrubtion (probs due to electronegativity) leading to dipole moment 

ions full charge is attractred to part of the partial charge on the dipole

one of the strongest IMF forces cause we’re using a full charge 

need an ionic compound

dipole - dipole forces

attraction of polar molecuels with permanent dipole moments, the partial negative ends of polar molecules are attracted to the partial positive ends of other polar molecules ( not as powerful as ion dipole)

higher dipole moment leads to higher boiling point and higher melting point ( harder to seperate the molecules )

melting point/boiling point

amount of energy u have to put in to separate the molecules

hydrogen bonding

especially strong form of dipole-dipole interaction

• occurs when hydrogen is bonded to N, O or F - due to their high electronegativity

• when they bond, the electronegative atoms, create an insanely polar bond with hydrogens 1 electron; esenittaly leaving hydrogens nucellus exposed

why does water form such a strong hydrogen bond 

• cause each water molecule can donate 2 Hs and accept 2 Hs, making them fit together like puzzle pieces; forming a strong network 

• 4 hydrogen bonds per oxygen leads to a full tetrahedral molecule, when these freeze they form a chain of tetrahedral molecules, creating an open and gap filled structure; which is why ice is less dense than water 

London Dispersion Forces

• for an instant in time, electrons can be asymmetrical arranged around the nucleaus such that the atom is polarized → instantaneous dipole → which induces a dipole in a neighbouring atom

• LDF are present in all molecules regardless of polarity

• affected by shape: longer molecules have stronger LDF, than spherical molecules cause of their increased surface area

• also affected by molecular weight; more electrons means easier for an instatnatous dipole to occur

• scale with weight, alter with form

Viscosity

• resistance to flow units: Pas ( equivalent to kgm^-1s^-1)

• higher viscosity = stronger intermolecular forces

• higher viscosity = molecular shape is easier to entangle

what creates surface tension

• at the surface of a liquid there is an imbalance of intermolecular forces ( no forces above surface) -

• this creates a tight surface as forces are pulling down and away from top( surface tension)

• it is caused by cohesive forces ( binds molecules to like molecules) 

surface tension units

• units Jm^-2 - surface tension is the energy cost of creating more surface area ( work required to create more surface area)

examples of strong surface tension

• water droplets or bugs walking on water ( strong hydrogen bonds)

interfacial behaviour - cohesive forces

describes how 2 phases interact

• cohesive forces : binds like molecules to one another

• visually liquid beads up

• ( IMFs between like molecules have to be strong)

what creates a concave surface

when adhesive forces are greater than cohesive forces - molecule binds to surface as much as possible 

what creates a convex surface

cohesive forces are greater than adhesive forces, liquid does NOT bond to surface bonds internally instead

interfacial behaviour - adhesive forces

• Adhesive Forces: binds molecules to the surface ( liquid can form strong IMFS with surface)

• liquid spreads out

how to tell if cohesion of adhesion wins ( might not be needed for exam)

• like sticks to like

• polar liquid sticks to polar surfaces

• non polar liquids stick to non polar surfaces

• a mix of both? cohesion wins

capillary action

• ability of a liquid to rise or fall in a tube without external forces

• due to a mix of adhesive and cohesion forces

• adhesive forces draw liquid along sides of tube ( surface)

• cohesion forces between liquid molecule, pulls remaining molecules along with first molecules

adhesive forces are the catalyst but cohesion forces carry everything else along too

evaporation 

• occurs when molecules near surface of liquid get enough energy to overcome IMF and transition from liquid to gas 

• ease of evaporation dictates both boiling point and vapour pressure 

evaporation in open system vs closed system

• open system: molecules evaporate and are removed from system ( float away)

• closed system: molecules evaporate but condense back at the same rate, system is in equilibrium ( this effects vapour pressure )

vapour pressure

pressure exerted by temporary gas state of solid or liquid when it is in equilibrium with its original form

• in a closed system when rate of evaporation = rate of condensation pressure of gas at this point is vapour pressure 

• vapour pressure increase with temp → cause temp raises KE → molecules have more energy to escape 

• also increase with weaker IMFS ( weaker IMF → easier for molecules to escape IMFs)

• also increase with surface area → more molecules on surface→ more molecules that can escape for given time 

when is boiling point reached

when vapour pressure = external pressure

• when these two are equal , bubbles of vapor can form inside liquid not just on surface, → liquid boils

normal boiling point?

when liquids vapour pressure is 1 (atm) atmosphere = 760mm Hg

normal = standard atmospheric pressure

what happens to boiling point at higher altitudes?

• air pressure decreases as altitude increases,

• at higher alitidue liquids boil at lower temp

• however this means it takes longer for food to cook, because the liquid is cooler so it has less energy and cooking/heating things up takes longer 

volatile liquids

• evaporate easily

• have higher vapour pressure

Phase Change Diagram 

plots state of matter as function of pressure and temp 

sublimation: solid turns directly into gas

deposition: gas changes directly into a solid

Freezing: Liquid turns to solid

Melting: solid turns to liquid

Boiling: liquid turns to gas

condensing : gas turns to liquid 

triple point

all three phases present

super critical fluid

• has properties between liquid and gas

• ex: SC CO₂ Used to extract caffeine from coffee

what determines density

density = mass/volume , liquid water has more density than solid ( due to the shape of the solid having gaps)

liquid crystals

• molecule is above melting point ( should be liquid) but exhibits solid characteristics

• a true liquid is isotropic - molecules point in every possible direction ; not the case for liquid crystals

• its an opaque liquid with crystalline ordering but also ability to flow

• usually long and rod like - after second melting point are isotropic ( randomly oriented - become ‘true liquids’)

nematic liquid crystals 

crystals ordered along the long axis of molecule ( one direction)

smectic liquid crystal

• ordered along long axis AND another dimension ( sheets break it up?)

cholesteric liquid crystals

• ordered along long axis and in twisted layers?