States of Matter: Solid, Liquids and Phase Changes Flashcards
crystalline solids
- well-ordered matter within the solid
- arrangement of atoms in the solid repeats itself
amorphous solids
don’t have extensive ordering of particles
types of crystalline solids
- molecular solids
- covalent network solids
- metallic solids
- ionic solids
molecular solids
molecules held together by intermolecular forces, with relatively low melting points (eg ice or benzoic acid)
covalent network solids
extended structures of atoms held together by covalent bonds with very high melting points
allotropes
different structural forms of an element
metallic solids
- metallic bonding between atoms
- metal atoms as cations in sea of delocalised electrons
- high electrical conductivity
- malleable, ductile
ionic solids
held together by electrostatic attraction between cations and anions, with high melting points
describe how most liquids are molecular
intermolecular forces keep particles close together but not strong enough to keep particles from moving past each other
define surface tension
amount of energy required to expand a liquid surface
how does surface tension work?
at top - intermolecular attractions from below mean that the net attraction is down, causing surface to contract.
in the middle - there are intermolecular attractions in all directions; no net attraction
how does strength of forces between particles in a liquid affect surface tension?
the stronger the forces between particles in a liquid, the greater the surface tension
define capillary action
the rising of a liquid in a narrow space against the pull of gravity
cohesive molecules
between molecules
adhesive forces
between molecules and container walls
how does water rise in a capillary tube?
- strong H-bonding interactions between the water and glass (SiO2) pull the water up
- gravity and cohesive forces pull the liquid down
diagram for water in a tube when adhesive>cohesive and vice versa
gas -> solid
deposition
solid -> gas
sublimation
sold -> liquid
fusion (melting)
liquid -> gas
vaporisation
gas to liquid
condensation
liquid to solid
freezing
attractive forces in solids, liquids, gases
solid: many attractive forces
liquid: some attractive forces
gas: no attractive forces
breaking attractive forces
requires (absorbs) energy: endothermic, so ∆H is positive
- adding heat increases temperature, increases KE, and overcomes attractive forces
forming attractive forces
releases energy: exothermic ∆H is negative
graph for phase changes
moving from one state of matter to another
how does ∆H of vaporisation relate to ∆H of fusion?
∆H(vap)>∆H(fus) for all substances
what type of process is a phase change?
a spontaneous process
spontaneous
∆G (change in free energy) is negative
equation for ∆G
∆G=∆H-T∆S
both ∆H and ∆S determine whether a process occurs spontaneously
∆H =
change in enthalpy due to phase change
∆S =
change in entropy due to phase change
entropy, S
the greater the degree of randomness or disorder in a system, the greater its entropy
∆S =
S (final) - S (initial)
for a reaction to be favourable, ∆S must be
positive
solid -> liquid -> gas requires
heat and increases entropy
describe gas-liquid equilibrium and draw a graph for it
as equilibrium between Rate(vap) and Rate(cond) is achieved, gas pressure reaches a constant value
vapour pressure
the pressure of gas in equilibrium (co-existing) with its liquid (or solid) at a specified temperature
vapour pressure depends on
temperature and the strength of intermolecular forces
effect of boiling point on vapour pressure
- the temperature at which vapour pressure of a liquid equals the external pressure
- enough kinetic energy to vaporise against external pressure; water can form vapour bubbles within
normal boiling point
boiling point at 1atm
effect of intermolecular forces on vapour pressure
greater intermolecular forces:
- fewer molecules escape liquid
- more KE required to overcome
why Hvap is greater than H fus
In the change from solid to liquid, the kinetic energy of the molecules must
increase only enough to partially offset the intermolecular attractions between molecules. In the change from
liquid to gas, the kinetic energy of the molecules must increase enough to overcome the intermolecular forces
does Xe have a lower or higher boiling point than Iodine
Xe would have a lower boiling point than iodine. Both are nonpolar with dispersion forces, but the forces
between xenon atoms would be weaker than those between iodine molecules since the iodine molecules are more
polarizable because of their larger size
