1. Physiological Chemistry Chapter 5 States of Matter: Liquids and Solids

2. Changes in State Changes in state are considered to be physical changes During a change of physical state many other physical properties may also change This chapter focuses on the important differences in physical properties among –gases –liquids –solids

3. Comparison of Physical Properties of Gases, Liquids, and Solids

4. Changes of State A change of state is the process whereby a substance changes from one physical state to another (e.g., melting, boiling) –these changes can be effected by changing temperature and/or pressure endothermic: heat is absorbed, which increases the kinetic energy of the particles –melting, evaporation and boiling, sublimation exothermic: heat is released, which reduces the kinetic energy of the particles –freezing, condensation, deposition

5. Changes of State

6. The Liquid State Compressibility: liquids are practically incompressible Viscosity: a measure of a liquid’s resistance to flow –a function of both attractive forces between molecules and molecular geometry –complex or polar molecules tend to have higher viscosity than simpler or nonpolar molecules –viscosity decreases as temperature increases

7. Surface Tension Surface tension: is a measure of the attractive forces exerted among molecules at the surface of a liquid –Surface molecules are surrounded and attracted by fewer liquid molecules than those below –Net attractive forces on surface molecules pull them downward Surfactant: is a substance added which decreases the surface tension (soap)

8. Vapor Pressure of a Liquid Place water in a sealed container –Both liquid water and water vapor will exist in the container How does this happen below the boiling point? –the gas particles that escape the liquid phase are called the vapor of the liquid –they would not normally exist in the gas phase at the temperature/pressure at which evaporation is occurring Kinetic molecular theory: liquid molecules are in continuous motion, with their average kinetic energy directly proportional to the Kelvin temperature

9. energy + H 2 O(l)  H 2 O(g) Temperature Dependence of Vapor Pressure As temperature increases, the fraction of molecules having the activation energy necessary for evaporation increases (shaded area) due to their higher kinetic energy Even though at cold temperatures some molecules can be converted to vapor, molecules with higher kinetic energy have more propensity to escape from the liquid phase

10. Evaporation Evaporation is the process by which particles on the surface of a liquid acquire enough energy to overcome the attractive forces within the liquid, escaping the liquid phase and entering the gas phase –As liquid particles evaporate: the volume of the liquid decreases – if the liquid is in an open container it eventually disappears! the temperature of the liquid decreases (as long as the container does not exchange heat with the environment) – evaporation has a COOLING effect! –as particles evaporate they absorb heat from the liquid

11. Boiling When the vapor pressure is high enough that it equals the atmospheric pressure, bubbles begin to form throughout the volume of the liquid and it begins to boil –boiling is the process of evaporation in which the liquid is transformed into a gas within the body of the liquid through bubble formation –the boiling point of the liquid is the temperature at which its vapor pressure equals the surrounding atmospheric pressure at high altitudes, where the atmospheric pressure is lower than sea level, water boils below 100°C inside a pressure cooker, where the pressure is much higher than atmospheric, water boils above 100°C

12. Boiling Point Boiling point, and melting point, is dependent on intermolecular forces –the stronger the intermolecular (attractive) forces that exist within the liquid, the more difficult for a liquid to evaporate, thus it has a LOWER vapor pressure, and HIGHER boiling point liquids that evaporate easily (due to weak intermolecular forces) have HIGHER vapor pressures (and LOWER boiling points) and are said to be volatile –polar molecules have higher b.p. than nonpolar molecules of similar molar mass

13. Intermolecular Forces Physical properties of matter are explained in terms of their intermolecular forces Different substances melt or boil at different temperatures because the strength of the intermolecular forces that hold particles together within matter varies among different substances There are three major types of intermolecular forces that affect and determine the behavior of matter Understanding the nature of these forces is of fundamental importance in understanding the physical and chemical properties of matter, including the multitude of complex biological molecules that are responsible for life (e.g., proteins, carbohydrates, and RNA/DNA)

14. London Dispersion Forces These are the weakest of the intermolecular forces They are not fixed but rather develop momentarily and intermittently between molecules as they approach each other and their electron clouds become briefly distorted and instantaneously polarized The strength of London forces depends on the ease of distortion of the electron cloud, which increases with size and molar mass

15. Dipole-Dipole Interactions These forces develop between polar molecules when the negative end of one molecule is attracted to the positive end of another The greater the polarity of the molecules, the stronger the attraction between them, and the higher the melting and boiling point of the substance

16. Hydrogen Bonding Hydrogen bonding: –is a special type of dipole-dipole attraction –is a very strong intermolecular attraction causing higher than expected b.p. and m.p. Requirement for hydrogen bonding: –a hydrogen atom directly bonded to O, N, or F atom qualifies a molecule to both donate and accept H in H-bonding interactions

17. Hydrogen Bonds Based on its molar mass, the calculated boiling point of water should be around −80°C, making life on Earth impossible! The strength and number of hydrogen bonds in water is what makes ice float on water, another reason why life on Earth flourished and was able to survive the ice ages.

18. Examples of Hydrogen Bonding Hydrogen bonding has an extremely important influence on the behavior of many biological systems Water forms FOUR hydrogen bonds in the solid state, but on average forms less than four in the liquid state Water molecules in the solid state are perfectly arranged in a tetrahedral fashion with respect to one another, making ice less dense than liquid water – thus ice floats on water!

19. Hydrogen Bonds

22. The Solid State Particles are closely packed due to attractive forces strong enough to resist motion Properties of solids: –fixed shape and volume –incompressible –m.p. depends on strength of attractive force between particles –a solid may be crystalline (ordered array of particles) or amorphous (disordered arrangement of particles)

23. Types of Crystalline Solids 1.Ionic solids –made up of positive and negative ions –high m.p. and b.p. –hard and brittle –a common example is NaCl 2.Covalent solids –held together entirely by covalent bonds –high m.p. and b.p. –extremely hard –an example is diamond

24. Types of Crystalline Solids 3.Molecular Solids –made up of molecules held together by intermolecular attractive forces –usually soft with low m.p. –volatile and poor electrical conductors –a common example is ice

25. Types of Crystalline Solids 4.Metallic Solids –made up of metal atoms held together by “metallic bonds” –these bonds are formed by the overlap of metal atomic orbitals –there are regions of high electron density, which are very mobile and move freely from atom to atom –this results in high conductivity –examples include Ag and Cu

26. Sublimation of Solids Sublimation: process of conversion of molecules in the solid state directly to the gaseous state An example is dry ice (solid carbon dioxide), which converts directly to a gas at atmospheric pressure Solid water (i.e., ice) slowly sublimes, which is why snow flurries don’t last for long even on a dry cold day

27. Sublimation and deposition of iodine (I 2 )