Week Ten Flashcards
Properties of liquids
strong inter molecular
free to move
cannot expand or contract significantly
Surface tension
measure of resistance of a liquid to an increase in its surface area
attraction of liquids on surface
net attractive forces that pulls them towards the interior of the liquid
middle of liquids - attraction
no net pull in any direction in the middle
Spherical shape of water
have less surface area so water is attracted to this shape
Forces that molecules in contact with surface have
adhesion
cohesion
adhesion forces
attract molecules in liquid to walls of container
cohesion forces
attract molecules in liquid to one another
Meniscus
curves surface of liquid, formed by adhesion and cohesion forces
strength of forces
adhesion forces are stronger than cohesive
Capillary action
upward movement of water against gravity
due to adhesion forces between water and oxygen
Viscosity
liquids resistance to flow
greater the viscosity, slower the liquid pours
What affects viscosity
combination of molecular shape and intermolecular forces
affected by temp
Viscosity vs temp
decreases as temp increase due to higher kinetic energy
Vapour pressure
the pressure that occurs when both the gas phase and liquid phase are happening at same time
vapour pressure in closed container
partial pressure increases as molecules enter gas phase
vapor pressure in open container
will continually lose molecules to gas phase until fully evaporated
Vapor pressure vs temp
at high temp, amount of molecules that can leave phases increases
liquid kinetic energy
have energy kinetic energy to move around within liquid phase
What affects vapour pressure
strength of intermolecular forces and temp
Vapor pressure - boiling liquid
when vapor pressure equals outside pressure, liquid will boil
Low pressure - vapor pressure - boiling
boiling point lower, pressure easily reached
high pressure - vapor pressure - boiling
boiling point increases,
Magnitude of forces in solids
forces range from large to small
can be bound by various attractive forces
intermolecular forces
molecular solids
covalent bonds
network solids
delocalised bonds
metallic solids
electrostatic bonds
ionic solids
Molecular solids
aggregates of molecules are bound by dispersion, dipolar, hydrogen bonding or combination
Network solids
have high melting points
held by covalent bonds
bonding pattern determines properties
Metallic solids
array of metal atoms embedded in mobile valence electrons
ductile and malleable
range of properties
properties determined by position on periodic table and number of valence electrons
Ionic solids
contain cation and anions strongly attracted to each other by electrostatic forces
determined by charges
temp during phase change
remains same
what must happen for temp to change during phase change
substance must be completely changed
phase changes - energy
require that energy is either supplied or removed from substance
Phase change
transition of substance from one phase to other
depends on temp, pressure, magnitude of bonds and intermolecular forces
molar enthalpy of melting
for solid to liquid
heat needed to melt 1 mole of substance at normal melting point
Molar enthalpy of vaporisation
for liquid to gas
heat needed to vaporise 1 mole of substance at normal boiling point
Molar enthalpy of sublimation
for solid to gas
heat needed to vaporise 1 mole of substance from solid phase
vaporisation
liquid to gas
condensation
gas to liquid
fusion
freezing
Sublimation
solid to gas
Deposition
gas to solid
Enthalpy of phase change
reverse processes have same magnitude but opposite signs
Supercritical Fluids
form upon compression of gases at high temp or heating liquid to very high temp at high pressure
Supercritical fluid properties
of both gas and liquid
Critical point
density of gas and liquid phase are equal and no phase boundary
critical temp and pressure
Phase diagrams
summarise phase behaviour of a substance
Boundary lines
between phases that separates regions where each phase is stable
triple point
where boundary lines meet at single point
x ray diffraction
determines the arrangement of atoms in crystalline structure
incoming x rays
atoms emit outgoing x rays
in phase and outer phase - direction
in different directions
constructive interference
cause diffracted beams in exposed crystals
bragg equation
relates wavelength of x rays to the direction and distance between the atom planes
n(wavelength) = 2dhklsin0
where dhkl = distance between planes
amorphous
when solids form rapidly and molecules are locked into irregular positions
amorphous solids
dont diffract xrays
defects in solids
alters properities of solid material
ceramics
very hard and high melting point
dont conduct electricity
contain metals with high oxidation stat and non metals with high negative oxidation state
hexagonal arrangment of spheres in solids
most dense packing
second layer of solids
sit in the dimples of first layer
third layer of solids
can be in same positions as first layer or in new positions
hexagonal closed packing
when 3rd layer is a repeat of first layer
cubic closed packing
when third layer is different to first 2 layers
crsytal shape
lattice
pattern of point
primitive cubic lattice
layers of atoms stacked directly above another so that all atoms lie along straight lines at right angles
body centred cubic lattice
simple cube with one entire atom in the center of cube
face centred cubic lattice
simple cube with atom in centre of each face of cube
ionic crystals
ions of opposite charge alternate with one another to maximis attraction and minimise repulsion
ionic solids - packing principle
closed packed structures formed by larger ions
smaller ions dill interstitial spaces
