Liquids and Solids Ch 11

Comparison of Liquids and Solids to Gases  Liquids & solids are much more dense than gases  Inorganic liquids and solids have densities ranging from 1 – 8 g/cm 3, some up to 20 g/cm 3  Most organic liquids & solids have densities ranging from 0.7 – 2.0 g/cm 3  Gas densities are usually between and g/c, 3

Comparison of Liquids and Solids to Gases  Gases expand to fill all available space & must be kept in enclosed containers  Liquids fills any container from the bottom up to a level dictated on by the mass of the liquid present  Liquids conform to the shape of their container  Solids maintain shape without a container

Comparison of Liquids and Solids to Gases  Gases lack significant attractive forces  Liquids and Solids- significant attractive forces!

Intermolecular Forces  3 Types to be aware of:  Dipole-Dipole Forces  London Dispersion Forces  Hydrogen Bonding

Dipole-Dipole Forces  Molecular compounds share electrons in a covalent bond, usually not equally!  e - congregate at 1 end of the molecule, giving it polarity, creating a dipole.  Polar molecules are attracted to each other.  Attractive forces are represented by the equation:  Force = (δ+)(δ-) r 2

Dipole-Dipole Forces  For gases to become a liquid – the attractive forces must overcome the KE of the moving gas molecule.  Decreasing the distance b/w molecules increases the attractive force.  Increasing P on a gas forces the molecules closer together  Cooling a gas reduces its avg KE  Force = (δ+)(δ-) r 2

Dipole-Dipole Forces  Boiling Pt (Condensation Pt) – indicator of the attractive forces b/w molecules  Measure of how much KE has to be increased so it overcomes the attractive forces in the liquid.  Low BP – low attractive forces  High BP – higher attractive forces  Highly polar molecules have higher BPs.

London Forces of Attraction  Explain how nonpolar gases develop the forces necessary to condense into liquids.  Nonpolar atoms & molecules may become momentarily polar when an unsymmetrical distribution of their e - results in instantaneous dipoles.  Sometimes called dispersion forces, instantaneous dipole forces, or induced dipole forces.

London Forces of Attraction  Very weak attractive forces, leading to very low BPs. The halogens, like the noble gases, don’t have permanent dipoles, but… Iodine is a solid and bromine is a liquid at room T.

London Forces of Attraction  What’s up with I 2 and Br 2 ?  Polarizability of e - clouds!  The ease with which the e - cloud around an atom or molecule can be deformed into a dipole.  Small atoms/molecules have e - clouds held tightly to nucleus- low polarizability.  Large atoms/molecules, w/ loosely held e - have high polarizability

London Forces of Attraction  These forces can explain the behavior or many molecules.  The more e - in a molecule, the more opportunity to form instantaneous dipoles- so… increase in attractive forces means higher BPs. Aklanes – C n H 2n+2 Called normal alkanes, n- alkanes, because the vary in a regular way. (homologous series)

Hydrogen Bonding  Extraordinarily large dipole-dipole forces attributed to the large electronegativity difference between H and the other atom on the next molecule (F, N, or O).

Physical Properties of Liquids  Surface tension  Viscosity  Evaporation  Vapor Pressure  Boiling Pt.  Heat of Vaporizaiton

Surface Tension  Caused by an increase in the attractive forces b/w molecules at the surface of a liquid compared to the forces b/w molecules in the center (bulk) of the liquid.  Causes fluids to minimize their surface area…  Small droplets form spheres

Surface Tension  Look at a molecule on the interior…  The solvent molecule is surrounded by other solvent molecules on all sides.  Look at a molecule on the surface…  Some of the molecules surrounding the the solvent molecules have been removed so the surface molecules will compensate by attracting neighboring molecules more strongly to reduce added potential energy.  Causes surface molecules to be closer to each other.

Surface Tension  Cohesive forces – attractions b/w identical molecules in the liquid  Adhesive forces – attractions b/w different molecules, like a liquid and a flat surface If cohesive forces are stronger than adhesive forces… If adhesive forces are stronger than cohesive forces…

Viscosity  A liquid’s resistance to flow.  Attractive forces are responsible for viscosity.  Molecules move more freely in solutions with low attractive forces  Liquid alkanes have lower viscosities because they only have London forces  Water is more viscous because it has hydrogen bonding  Syrup is very viscous because all the bulky sugar molecules have lots of –OH groups, which hydrogen bond to the water in the mixture. Low viscosity High viscosity

Viscosity  Decreases as the liquid’s T is increased.  Molecules have higher KE, weakens intermolecular forces (IMFs).

Evaporation  The process in which a liquid in an open container is slowly converted into a gas at the surface of the liquid.  Some liquids evaporate more rapidly than others.  Reverse of condensation, must have enough sufficient KE to escape the attractive forces

Evaporation  Factors that affect evaporation  Surface area of the liquid – the greater the surface area, the greater the evaporation

Evaporation  Factors that affect evaporation  Temperature – Increasing the T increases the # molecules with enough KE to escape as a gas

Evaporation  Boiling – when T is increased enough, boiling occurs.  Molecules do not have to reach the surface to enter the gas phase.

Vapor Pressure  Pressure that develops in the gas phase above a liquid when the liquid is placed in a closed container. Dynamic equilibrium – occurs when the rate the liquid evaporates equals the rate the gas condenses

Vapor Pressure  Rate a liquid evaporates – dependent on T  Rate a gas condenses – dependent on the frequency the gas molecules collide with the liquid “wall” of the container.  Therefore – vapor pressure depends only one the nature of the liquid (attractive forces) & the temperature (KE)  If T increases, Vapor Pressure increases.

Boiling Point  Boiling occurs when the vapor pressure of the liquid is equal to atmospheric pressure  Normal boiling point- refers to the boiling point when atmospheric pressure is 760 mHg

Heat of Vaporization  ΔH vap - the energy needed to convert 1 gram of liquid into 1 gram of gas at a temperature equal to the normal boiling point of the liquid.  Units are J/g or J/mol (if using molar heat of vaporization)  ΔH vap = -ΔH cond

Heat of Vaporization  There are differences in the heats of vaporization that can be related to the IMFs  For similar-size molecules, hydrogen- bonded substances have largest ΔH vap.  Polar substances have higher ΔH vap than similar shape nonpolar substances  Increasing London forces increases ΔH vap

Heat of Vaporization CompoundFormulaΔH vap (kJ/mol) Attractive force WaterH2OH2O+43.9H bonding AmmoniaNH H bonding Hydrogen fluoride HF+30.2H bonding Hydrogen chloride HCl+15.6Dipole-dipole Hydrogen sulfide H2SH2S+18.8Dipole-dipole FluorineF2F2 +5.9London ChlorineCl London BromineBr London MethaneCH London EthaneC2H6C2H London PropaneC3H8C3H London The amount of heat needed to vaporize a liquid is very large. This explains why water can be quickly raised to its boiling point, but a long time is needed to boil away all the water.