Liquids & Solids

1. Liquids & Solids Chapter 12

2. Properties of Liquids • Have definite volume & take shape of container • Lower kinetic energy than gases • Particles move about constantly, more orderly, more attractive forces between particles • Liquids are fluids (they flow)

3. Properties of Liquids • Higher density (x 1000s) than gases • Most have lower density than solids, (H2O exception)

4. Properties of Liquids • Transmit pressure equally in all directions, because they are relatively incompressible.

5. Properties of Liquids • Ability to diffuse more slowly than gases; will diffuse more rapidly with increase in temp.

6. Properties of Liquids • Surface tension increases with more attraction between molecules

7. Cohesion & Adhesion • Cohesion – force of attraction between particles of the same substance • Adhesion –force of attraction between particles of different substances

8. Properties of Liquids • Capillary action – the attraction of the surface of a liquid to the surface of a solid.

9. Vaporization • The process by which a liquid or solid changes to a gas. • Evaporation – particles escape from the surface of a nonboiling liquid and enter the gas state.

10. Vaporization • Boiling – change of a liquid to bubbles of vapor that appear throughout the liquid; not just at the surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure.

11. Freezing • When a liquid cools, average kinetic energy of the particles decreases. If low enough, the attractive forces between particles will pull the particles into an even more orderly arrangement. (solid)

12. Properties of Solids • Definite shape and volume • Intermolecular forces between particles are stronger, more closely packed. • Exhibit relatively fixed positions • More ordered arrangement than liquids or gases

13. Properties of solids • 2 types of solids • Crystalline – orderly, geometric arrangement • Definite melting point (temperature at which a solid becomes a liquid) • Amorphous – random arrangement • No definite melting point (glass, plastic, rubber) • Sometimes called supercooled liquids • Higher density than gases and most liquids • Higher incompressibility than liquids

14. Properties of Solids • Low rate of diffusion – while diffusion does occur in solids, it is millions of times slower than in liquids.

15. Crystalline Solids • Exist as • Single crystals • Groups of crystals fused together • Crystal structure – total 3-dimensional arrangement of particles of a crystal • Crystal lattice – a coordinate system of particles of a solid, consisting of multiple unit cells.

16. Crystalline solids • 7 types of symmetry

17. 4 Binding Forces in Crystal Structures • Ionic crystals – form hard, brittle crystals with high melting points and make good insulators – form between pos/neg ions. • Covalent network crystals – form large bonds between atoms that are covalently bonded to their nearest neighboring atoms – (Cx, SiO2, SiCx) – hard, brittle, rather high melting pts., non conductors or semi-conductors.

18. 4 Binding Forces in Crystal Structures • Metallic crystals – metal atoms are surrounded by a sea of valence electrons • Melting points vary greatly • Covalent molecular crystals – held together by weak intermolecular forces • Non-polar – H2, CH4, C6H6 – held together by weak London dispersion forces • Polar – H2O, NH3 – can be held together by dispersion forces, dipole-dipole forces, or even hydrogen bonding • Both have low melting points, easily vaporized, soft, and are good insulators

19. Amorphous Solids • - from the Greek word, “without shape” • Do not have a regular shape • Hold their shape for a long time, but DO flow!

20. Changes of State • Equilibrium – a dynamic (changing) condition in which two opposing changes occur at equal rates in a closed system. A closed system is one in which matter cannot enter or leave, but energy can. Equilibrium- individuals constantly coming & going, but total # people Stay the same. Not in equilibrium- in the morning more people come, as time passes, pool gets crowded.

21. Le Châtelier’s Principle • When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that minimizes the stress. A stress is typically a change in concentration, pressure, or temperature. • Ex. Suppose the temperature to a closed system is raised from 25oC to 50oC. • Liquid + heat energy vapor • Endothermic, shifts to the right in favor of the forward reaction.(vapor produced at higher rate, until equilibrium is attained.)

22. Ex. Holding mass and temperature constant in a closed system, yet increasing the volume causes a stress to the equilibrium of the system. What happens? • Concentration of the molecules in the system decreases because volume has increased. Because of this, fewer molecules strike the liquid surface and condense. Rate of evaporation is now higher than rate of condensation. • Liquid + heat energy vapor • Shifts to the right because of more evaporation

23. Equilibrium Shifts

24. Chemical Equilibrium(Ch.18) • A chemical reaction in which the products can react to re-form the reactants is called a reversible reaction. • A reversible reaction is in chemical equilibrium when the rate of its forward reaction equals the rate of its reverse reaction and the concentrations of its products and reactants remain unchanged.

25. Chemical reaction convention • A chemical reaction is written left  right (forward) and right  left (reverse).

26. Predicting the Direction of Shift • Changes in Pressure • A change in pressure affects only equilibrium systems in which gases are involved. • For changes in pressure to affect the system, the total number of moles of gas on the left side of the equation must be different from the total number of moles of gas on the right side of the equation.

27. Change in Pressure • Example • Let’s consider the synthesis of ammonia: • N2(g) + 3H2(g) 2NH3 • If pressure is increased on this system, equilibrium shifts to the right. This is because we have 4 moles of total gas on the left, and only 2 moles on the right. The system can reduce the number of molecules, and hence, the total pressure, by shifting the equilibrium in the direction of the lesser number of moles.

28. Remember: • Changes in partial pressure of a gas within a reaction WILL cause a shift in equilibrium. • Changes in pressure cause changes in concentration of gases within a reaction. • The addition or removal of some other gas has no effect even though it changes the total pressure.

29. Changes in Temperature • Reversible reactions are exothermic in one direction and endothermic in the other. • According to Le Châtelier’s principle, the addition of heat shifts the equilibrium so that heat is absorbed, which favors the endothermic reaction. The reverse is true for exothermic reaction. • HINT: Treat heat energy absorbed as a reactant, and heat energy released as a product, to determine shift.

30. Changes in Temperature • Example • 556 kJ + CaCO3(s) CaO(s) + CO2(g) • An increase in temperature caused by adding heat to the system would shift the equilibrium to the right. A decrease in temperature would shift it in the reverse direction.

31. Changes in Concentration • An increase in the concentration of a reactant is a stress on the equilibrium system. It causes an increase in collision frequency and generally an increase in reaction rate. • Ex. A + B C + D • An increase in the concentration of A shifts the equilibrium to the right. • A decrease in the concentration of A shifts to the left.

32. Possible changes of state

33. Phase – any part of a system that has uniform composition and properties.

34. Evaporation can be expressed as: • Liquid + heat energy  vapor • When a liquid changes to a vapor, it absorbs heat energy from its surroundings. • Condensation can be expressed as: • Vapor  liquid + heat energy • When a vapor condenses, it gives off energy to its surroundings. • Liquid-Vapor equilibrium is expressed: • Liquid + heat energy vapor • Double arrow indicates a reversible reaction

35. Phases of Matter • Phase- any part of a system that has uniform composition and properties • Equilibrium vapor pressure – the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature. • Volatile liquids – liquids that evaporate readily, have weak forces of attraction between their particles.

36. Which liquid is more volatile? • Oil, water, alcohol

37. Oil, water, alcohol • Which liquid is more evaporates most readily? • Which liquid exhibits a higher amount of kinetic energy at room temperature? • Which liquid has the weakest attractive forces between molecules? • Which liquid is least volatile?

38. Boiling • A liquid boils when it has absorbed enough energy to evaporate. • Boiling – the conversion of a liquid to a vapor within the liquid as well as at its surface. • Boiling point – the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure. • The lower the atmospheric pressure is, the lower the boiling point. (Remember, boiling water in the bell jar?)

39. The normal boiling point of each liquid shown occurs when its equilibrium vapor pressure equals 760 torr. Energy must be added continuously in order to keep a liquid boiling. The temperature of the liquid and its vapor at the boiling point remains constant despite the continuous addition of energy. The added energy is used to overcome the attractive forces between the molecules of the liquid during the phase change, and is stored in the vapor as PE. • s

40. Molar Enthalpy of Vaporization • Symbol - ΔHv • The amount of energy as heat that is needed to vaporize one mole of liquid at the liquid’s boiling point at constant pressure is called the liquid’s molar enthalpy of vaporization. • Measures the attraction between particles of a liquid. • The stronger the attraction is, the more energy is required to overcome it, resulting in a higher molar enthalpy of vaporization.

41. Freezing • Occurs when a substance loses enough heat energy to solidify. • Liquid  solid + energy (results in loss of energy) • The normal freezing point – the temperature at which the solid and liquid are in equilibrium at 1 atm, 760 torr, or 101.3 kPa, pressure. (for pure crystalline substances) • Freezing occurs at constant temperature, and the energy lost during freezing is the PE that was present in the liquid.

42. Melting • Also occurs at constant temperature • Solid + energy  liquid • For pure crystalline solids, the melting point and freezing point are the same. At equilibrium, melting rate = freezing rate

43. Molar Enthalpy of Fusion • The amount of energy as heat required to melt one mole of solid at the solid’s melting point is the solid’s molar enthalpy of fusion. • Symbol – ΔHf • Energy absorbed increases the solid’s PE as its particles are pulled apart, overcoming the attractive forces holding them together.

44. Sublimation & Deposition • At sufficiently low temperature and pressure conditions, a liquid cannot exist. • A solid substance exists, in this case, in equilibrium with its vapor instead of its liquid. • Solid + energy  vapor • Change of state from solid to gas-sublimation • Reverse process - deposition

45. Phase Diagram of Water • A graph of pressure vs. temperature that shows the conditions under which the phases of water exist. It also reveals how the states of a system change with changing temperature or pressure.

46. Triple point – the temp. and pressure at which the solid, liquid, and vapor of the substance can coexist at equilibrium.Critical point- the critical temperature & pressure of the substance.Critical temperature- (tc), the temp above which the substance cannot exist in the liquid state.Critical pressure-(pc), the lowest pressure at which the substance can exist as a liquid at the critical temperature.

47. Water – Most Abundant Liquid • 75% of Earth’s surface – oceans, rivers, lakes • Frozen in glaciers • 70%-90% of the mass of all living things • Chemical reactions of life take place in water

48. Structure of Water • Bent shape, polar molecule • Intermolecular – between molecules, hydrogen bonding • Intramolecular– between atoms within a molecule – covalent bonding

49. Rigid, open structure of solid ice crystals, has much space between them making water less dense as a solid. • Water molecules are as tightly packed as possible at 3.980C. • Water’s relatively high boiling pt. is due to strong hydrogen bonding between molecules.

50. Properties of Water • Room Temperature-transparent, odorless, tasteless, almost colorless. • Freezes at a pressure of 1 atm • Melts at a temperature of 00C • Molar heat of fusion of ice = 6.009 kJ/mol • Density of ice = 0.917 g/cm3 at 00C • Density of liquid water = .999 84 g/cm3 at 00C