Unit 1: Liquids and Solids

Unit 1: Liquids and Solids

Kinetic-Molecular Description of Liquids and Solids • Solids and liquids are condensed states. • The atoms, ions, or molecules in solids and liquids are much closer to one another than in gases. • Solids and liquids are highly incompressible. • Liquids and gases are fluids. • They easily flow. • The intermolecular attractionsin liquids and solids are strong.

Table 13-1, p. 448

Kinetic-Molecular Description of Liquids and Solids • Schematic representation of the three common states of matter.

Kinetic-Molecular Description of Liquids and Solids • If we compare the strengths of interactions among particles and the degree of ordering of particles, we see that Gases< Liquids < Solids • Miscible liquids are soluble in each other. • Examples of miscible liquids: • Water dissolves in alcohol. • Gasoline dissolves in motor oil. • The natural diffusion rate of liquids is slower than gases

Kinetic-Molecular Description of Liquids and Solids • Immiscible liquids are insoluble in each other. • Two examples of immiscible liquids: • Water does not dissolve in oil. • Water does not dissolve in cyclohexane.

Kinetic-Molecular Description of Liquids and Solids • Solid particles do not readily diffuse into other solids • However, analysis of 2 different blocks of solids e.g. Cu and Pb that have been pressed together for a period of years show that each block contains some atoms of the other element  solids do diffuse but very slowly and if pressure is applied. Fig. 13-2, p. 449

Intermolecular Attractions and Phase Changes • Intermolecular forces(IMF) refer to the forces between individual particles (atoms, molecules, ions) of a substance • These forces are quite weak relative to intramolecular forces i.e. covalent and ionic bonds within compounds • If it were not for IMF, condensed phases would not exist • IMF influence physical properties

Intermolecular Attractions and Phase Changes The important intermolecular attractions from strongest attraction to the weakest attraction are: • Ion-ion interactions (ionic bond) • The force of attractionbetween 2 oppositely charged ions is directly proportional to the charges on the ions (say q+ and q-)  F α q+ x q- • Thus ionic substances containing multiple charged ions e.g. Mg2+, Al3+, etc. have stronger forces of attraction ( & thus higher m.p. and b.p.) than those with singly charged ions • The force of attractionbetween 2 oppositely charged ions is also inversely proportional to the square distance b/w the ions  F α 1/d2

Intermolecular Attractions and Phase Changes • Ion-ion interactions (ionic bond) • So for a series of similarly charged ions, the closer approach of smaller ions results in stronger interionic attractive forces  Higher m.p. and b.p Smaller the ions  stronger the ionic bond  higher m.p and b.p.

Intermolecular Attractions and Phase Changes • Example 1: Arrange the following ionic compounds in the expected order of increasing melting and boiling points. NaF, CaO,CaF2 You do it! What important points must you consider?

Intermolecular Attractions and Phase Changes 2. Dipole-dipole interactions • Occur between polar covalent molecules because of the attraction of the δ+ and δ- atoms of another molecule

Intermolecular Attractions and Phase Changes 2. Dipole-dipole interactions e.g. BrF (polar molecule). Each polar molecule is shaded with regions of high e-s density (red) and regions of high positive charge (blue). Attractive forces are shown as blue arrows and repulsive forces as red arrows Molecules tend to arrange themselves to maximize attractions by bringing regions of opposite charge together while minimizing repulsions by separating regions of like charge

Intermolecular Attractions and Phase Changes 2. Dipole-dipole interactions • An increase in temp  increase in translational, rotational & vibrational motion of molecules  random orientation of molecules relative to each other • Consequently, dipole-dipole interactions become less important as temp  • All these factors make compounds having only dipole-dipole interactions more volatile than ionic compounds

Intermolecular Attractions and Phase Changes 3. London Forces are very weak. • Also known as: • Instantaneous dipole-induced dipole interactions • Dispersion forces • London dispersion forces • They are the weakest of the intermolecular forces. • They exist in ALL molecules • This is the only attractive force in nonpolar molecules.

Intermolecular Attractions and Phase Changes 3. London Forces are very weak • In a group of Ar atoms, the temporary dipole in one atom induces other atomic dipoles. • Each atom’s e- cloud is attracted by the nucleus of the other atom or is repelled by the other atoms’ e-s cloud

Intermolecular Attractions and Phase Changes • Similar effects occur in a group of I2 molecules.

Intermolecular Attractions and Phase Changes Polarizability increases with increasing numbers of e- and therefore with increasing sizes of molecules Therefore, dispersion forces are generally stronger for larger molecules For molecules that are large or quite polarizable the total effect of the dispersion forces can even be higher than dipole-dipole interactions or H-bonding

Intermolecular Attractions and Phase Changes • NOTE: • The term “Van der Waals forces” usually refers to: • Dipole - dipole interactions • London forces

Intermolecular Attractions and Phase Changes 4. Hydrogen bonding • They are NOT really chemical bonds • Are a special case of a very strong dipole-dipole interaction. • Criteria for strong H-bonding: • A hydrogen bond donor: polar covalent molecule containing H attached to either one of the three small, highly electronegative elements – F, O or N. • A hydrogen bond acceptor: highly electronegative elements – F, O or N.

Intermolecular Attractions and Phase Changes 4. Hydrogen bonding • Consider H2O molecules • Each O atom can form two H-bonds and • Each H atom can form one H-bond δ+ δ- δ+ δ+ δ- δ+

Intermolecular Attractions and Phase Changes Hydrogen bonding in (b) methanol, CH3 OH and (c) ammonia, NH3 The H-bonds are due to electrostatic attraction between the δ+ charged H of one molecule to the δ- charge O or N of another. Fig. 13-4, p. 452

Intermolecular Attractions and Phase Changes 4. Hydrogen bonding • Typical H- bond energies range from 15 – 20kJ /mol • This is four – five times greater than that of dipole-dipole interactions • As a result, H-bonds exert a considerable influence on the properties of substances • E.g. H-bonds are responsible for the unusually high b.p. and m.p. of water, methanol and ammonia

Intermolecular Attractions and Phase Changes • The unusually high bps of NH3, H2O and HF are due to H-bonding • CH4 is non-polar  only weak London forces • As molecular mass increases (e.g Grp 4) bp increases because of increased dispersion forces. Boiling points of some hydrides as a function of molecular weight

So to summarize….. Note: ALL molecules contain dispersion forces

The Liquid State - Properties Viscosity • Viscosity is the resistanceto flow. • For example, compare how water pours out of a glass (low viscosity) compared to molasses, syrup or honey (high viscosity). • Oil for your car is bought based on this property. • 10W30 or 5W30 describes the viscosity of the oil at high and low temperatures.

The Liquid State - Properties Viscosity • For a liquid to flow, molecules must be able to slide past each other. • The stronger the intermolecular forces (IM)  the more viscous the liquid Substances that have a great ability to form H-bonds usually have high viscosities

The Liquid State - Properties Viscosity • Increasing the size and surface area of molecules  increased viscosity due to increased dispersion forces Pentane, C2H5 Viscosity = 0.215 centipoise at 25oC dodecane, C12H26 Viscosity = 1.38 centipoise at 25oC

The Liquid State - Properties Viscosity • As temp , molecules move more rapidly, their kinetic energies are able to overcome IM forces  a decrease in viscosity • An example of viscosity of two liquids.

The Liquid State - Properties Surface Tension • Molecules below the surface of a liquid are influenced by IM attractions from all directions • Those on the surface are attracted unevenly; are only attracted toward the interior  pulls the surface toward the center • Surface tension is a measure of the unequal attractions that occur at the surface of a liquid. • It is a measure of the forces that must be overcome to expand the surface area of a liquid

The Liquid State - Properties Surface Tension Fig. 13-9, p. 457

The Liquid State - Properties Surface Tension • Coating glass with silicone polymer greatly reduces adhesion of water to the glass • The left side of each glass has been treated with Rain-X which contains a silicone polymer. • Water on the treated side forms droplets that are easily swept away. p. 457

The Liquid State - Properties Droplets of mercury of glass • The small droplets are almost spherical, whereas larger ones are flattened due to the effects of gravity • This shows that surface tension has more influence on the shape of the small (lighter droplets) Surface Tension p. 457

The Liquid State • Floating paper clip demonstration of surface tension.

The Liquid State Capillary Action • Capillary action is the ability of a liquid to rise (or fall) in a glass tube or other container

The Liquid State • Cohesive forces are the forces that hold liquids together. • Adhesive forces are the forces between a liquid and another surface. • Capillary rise implies that the: • Adhesive forces > cohesive forces • Capillary fall implies that the: • Cohesive forces > adhesive forces

Mercury Water The Liquid State • Water exhibits a capillary rise. • Mercury exhibits a capillary fall.

The Liquid State • Capillary action also affects the meniscus of liquids. • Capillary action helps plant roots take up water and dissolved nutrients from soil • Roots, like glass, exhibit strong adhesive forces for water.

The Liquid State Evaporation • Evaporation is the process in which molecules escape from the surface of a liquid and become a gas. • Evaporation is temperature dependent.

The Liquid State Evaporation Liquid continuously evapourates from an open vessel Equilibrium between liquid and vapour is established in a closed container in which molecules return to the liquid at the same rate as they leave it. A bottle in which liquid-vapour equilibrium has been established. The droplets have condensed Fig. 13-10ab, p. 458

The Liquid State Distribution of kinetic energies of molecules in a liquid at different temperatures. At the lower temperature, a smaller fraction of the molecules have the same energy required to escape from the liquid, so evapouration is slower.

The Liquid State Vapor Pressure • DEFINITION: Vapor pressure is the pressure exerted by a liquid’s vapour on its surface at equilibrium. • Because rate of evapouration increases in increasing temperature  Vapour pressure of liquids always increases with increasing temperature

The Liquid State • Vapor Pressure (torr) and boiling point for three liquids at different temperatures. 0oC20oC30oC normal boiling point diethyl ether 85 442 647 36oC ethanol 12 44 74 78oC water 5 18 32 100oC • Easily vapourized liquids are said to be volatile • They have relatively high vapour pressures

The Liquid State Vapor Pressure as a function of temperature. Notice that the plot is not linear Each substance exists as a liquid for temp & presure to the left of its curve Each substance exists as a gas for temp & pres to the right of its curve The normal boiling point of a liquid is the temp at which its vapour pressure = 1 atm (760 torr) What are the intermolecular forces in each of these compounds? You do it!

Increasing temperature Higher vapour pressure The Liquid State Vapor Pressure Stronger attractive forces Lower vapour pressure Higher boiling point

The Liquid State Boiling Points and Distillation • The boiling point is the temperature at which the liquid’s vapor pressure is equal to the applied pressure. • The normalboiling point is the boiling point when the pressure is exactly 1 atm (760 torr). • E.g. water boils at 100 0C at 1 atm • If the applied pressure is lower than 1 atm, e.g. on a mountain water boils at a lower temp • Takes longer to cook food on a mountain because the temp of boiling water is lower

The Liquid State Distillation • Distillation is a process in which a mixture or solution is separated into its components on the basis of the differences in boiling points of the components. • Different liquids have different cohesive forces  different vapour pressures  boil at different temp • Distillation is another vapour pressure phenomenon.

The Liquid State Lab setup for distillation • During distillation of an impure liquid, nonvolatile substances remain in the flask • The liquid is vapourized and condensed before being collected in the receiving flask.

The Liquid State Heat Transfer Involving Liquids • The specific heat (J/g .oC) or molar heat capacity (J/mol .oC) of a liquid is the amount of heat that must be added to a stated mass of liquid to raise its temp by 1 oCwith no change in phase • It is given the symbol “C”

The Liquid State Heat Transfer Involving Liquid • Example : How much heat is released by 200. g of H2O as it cools from 85.0oC to 40.0oC? The specific heat of water is 4.184 J/goC.

Unit 1: Liquids and Solids