Liquids and Solids

1. Liquids and Solids Ch 11

2. 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/cm3, some up to 20 g/cm3 • Most organic liquids & solids have densities ranging from 0.7 – 2.0 g/cm3 • Gas densities are usually between 10-2 and 10-4g/c,3

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

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

5. Intermolecular Forces • 3 Types to be aware of: • Dipole-Dipole Forces • London Dispersion Forces • Hydrogen Bonding

6. 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-) r2

7. 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-) r2

8. 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.

9. 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.

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

11. London Forces of Attraction • What’s up with I2 and Br2? • Polarizabilityof 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

12. London Forces of Attraction Aklanes – CnH2n+2 Called normal alkanes, n-alkanes, because the vary in a regular way. (homologous series) • 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.

13. 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).

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

15. 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

16. 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.

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

18. Viscosity Low viscosity High 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.

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

20. 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

21. Evaporation • Factors that affect evaporation • Surface area of the liquid – the greater the surface area, the greater the evaporation

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

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

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

25. 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.

26. 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

27. Heat of Vaporization • ΔHvap- 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) • ΔHvap= -ΔHcond

28. 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 ΔHvap. • Polar substances have higher ΔHvap than similar shape nonpolar substances • Increasing London forces increases ΔHvap

29. 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. Heat of Vaporization