Properties of Liquids and Intermolecular Bonds

1. Chapter 13 Liquids and Solids by Christopher Hamaker Chapter 13

2. Properties of Liquids • Unlike gases, liquids do not respond dramatically to temperature and pressure changes. • We can study the liquid state and make five general observations. • Liquids have a variable shape, but a fixed volume. • Liquids take the shape of their container. • Liquids usually flow readily. • However, not all liquids flow at the same rate. • Liquids do not compress or expand significantly. • The volume of a liquid varies very little as the temperature and pressure change. Chapter 13

3. Properties of Liquids, Continued 4. Liquids have a high density compared to gases. • Liquids are about 1000 times more dense than gases. 5. Liquids that are soluble mix homogeneously. • Liquids diffuse more slowly than gases, but eventually form a homogeneous mixture. Chapter 13

4. Intermolecular Bond Concept • An intermolecular bond is an attraction between molecules, whereas an intramolecular bond is between atoms in a molecule. • Some properties of liquids, such as vapor pressure, viscosity, and surface tension, are determined by the strength of attraction between molecules. • Intermolecular bonds are much weaker than intramolecular bonds. Chapter 13

5. Intermolecular Bonds • Recall that a polar molecule has positive and negative charges concentrated in different regions due to unequal sharing of electrons in bonds. • This uneven distribution of electrons in a molecule is called a dipole. • Intermolecular attractions result from temporary or permanent dipoles in molecules. • There are three intermolecular forces: • Dispersion forces • Dipole forces • Hydrogen bonds Chapter 13

6. Dispersion Forces Dispersion forces, or London forces, are the result of a temporary dipole. Electrons are constantly shifting, and a region may become temporarily electron poor and slightly positive, while another region becomes slightly negative. • This creates a temporary dipole, and two molecules with temporary dipoles are attracted to each other. Chapter 13

7. Dispersion Forces, Continued • Dispersion forces are the weakest intermolecular force. • Dispersion forces are present in all molecules. • The strength of the dispersion forces in a molecule is related to the number of electrons in the molecule: • The more electrons in a molecule, the stronger the dispersion forces. Chapter 13

8. Dipole Forces Polar molecules have a permanent dipole. The oppositely charged ends of polar molecules are attracted to each other; this is the dipole force. The strength of a dipole force is typically 10% of a covalent bond’s strength. Dipole forces are stronger than dispersion forces. Chapter 13

9. Hydrogen Bonds Hydrogen bonds are a special type of dipole attraction. Hydrogen bonds are present when a molecule has an N—H, O—H, or F—H bond. Hydrogen bonds are especially important in water and living organisms. Chapter 13

10. Physical Properties of Liquids • There are four physical properties of liquids that we can relate to the intermolecular attractions present in molecules: • Vapor pressure • Boiling point • Viscosity • Surface tension Chapter 13

11. Vapor Pressure • At the surface of a liquid, some molecules gain enough energy to escape the intermolecular attractions of neighboring molecules and enter the vapor state. This is evaporation. • The reverse process is called condensation. • When the rates of evaporation and condensation are equal, the pressure exerted by the gas molecules above a liquid is called the vapor pressure. Chapter 13

12. Vapor Pressure, Continued • The stronger the intermolecular forces between the molecules in the liquid, the less molecules that escape into the gas phase. • As the attractive force between molecules increases, vapor pressure decreases. Chapter 13

13. Vapor Pressure Comparison Let’s compare water and ether. Water has strong intermolecular attractions, and ether has weak intermolecular attractions. At 0 C, neither has a significant vapor pressure. At 35 C, ether has a significant vapor pressure and water does not. Chapter 13

14. Vapor Pressure Versus Temperature As the temperature increases, the vapor pressure of a liquid increases. Again, the stronger the intermolecular attractions, the lower the vapor pressure at a given temperature. Chapter 13

15. Boiling Point • The normal boiling point of a substance is the temperature at which the vapor pressure is equal to the standard atmospheric pressure. • As we saw in the previous graph, the stronger the intermolecular attractions, the higher the boiling point of the liquid. • A liquid with a high boiling point has a low vapor pressure. Chapter 13

16. Viscosity • The viscosity of a liquid is a liquid’s resistance to flow. • Viscosity is the result of an attraction between molecules. • The stronger the intermolecular forces, the higher the viscosity. Chapter 13

17. Surface Tension • The attraction between molecules at the surface of a liquid it called surface tension. • For an object to sink in a liquid, it must first break through the surface. • The stronger the intermolecular attractions, the stronger the surface tension of a liquid. Chapter 13

18. Properties of Solids • Unlike gases, solids do not respond dramatically to temperature and pressure changes. • We can study the solid state and make five general observations. • Solids have a fixed shape and volume. • Unlike liquids, solids are rigid. • Solids are either crystalline or noncrystalline. • Crystalline solids contain particles in a regular, repeating pattern. Chapter 13

19. Properties of Solids, Continued • Solids do not compress or expand to any degree. • Assuming there is no change in physical state, temperature and pressure have a negligible effect on the volume of a solid. • Solids have a slightly higher density than their corresponding liquid. • One important exception is water; ice is less dense than liquid water. • Solids do not mix by diffusion. • The particles are not free to diffuse in a solid heterogeneous mixture. Chapter 13

20. Crystalline Solids There are three types of crystalline solids, examples of which are shown below: Ionic solids, such as NaCl Molecular solids, such as S8 Metallic solids, such as Cu Crystalline network solids, such as diamonds Chapter 13

21. Ionic Solids A crystalline ionic solid is composed of positive and negative ions arranged in a regular, repeating pattern. In table salt, NaCl, sodium ions and chloride ions are arranged in a regular three-dimensional structure referred to as a crystal lattice. • Other ionic compounds will have different crystal lattices. Chapter 13

22. Molecular Solids A crystalline molecular solid has molecules arranged in a particular conformation. In sulfur, S8, the molecules are arranged in a regular three-dimensional structure. • Other examples of crystalline molecular solids are table sugar, C12H22O11, and water, H2O. Chapter 13

23. Metallic Solids A crystalline metallic solid is composed of metal atoms arranged in a definite pattern. A metallic crystal is made up of positive metal ions surrounded by valance electrons. Metals are good conductors of electricity because electrons are free to move about the crystal. • This is referred to as the “electron sea” model. Chapter 13

24. Diamond • Diamond is a special type of crystalline solid that has covalent bonds between large numbers of atoms. • This type of crystalline solid is referred to as a network solid. • Diamond is very hard and has a very high melting point. Chapter 13

25. General Properties of Solids Chapter 13

26. Changes in Physical State • Heat is necessary to raise the temperature and change the physical state of a substance. • Specific heat is the amount of heat required to raise 1.00 g of a substance 1 C. • Water is the reference and its specific heat is 1.00 cal/(g xC). • The specific heats of ice and steam are about half that of liquid water. Chapter 13

27. Solid–Liquid Phase Changes • As a solid melts, the temperature is constant until all of the solid is changed to liquid. • The amount of heat required to melt 1.00 g of substance is called the heat of fusion(Hfusion). For water, the heat of fusion is 80.0 cal/g. • When a liquid changes to a solid, the heat change is the heat of solidification (Hsolid). • The value of Hfusion is the same as the value of Hsolid. Chapter 13

28. Liquid–Gas Phase Changes • As a liquid vaporizes, the temperature is constant until all of the liquid is changed to gas. • The amount of heat required to vaporize 1.00 g of substance is called the heat of vaporization(Hvapor). For water, the value is 540 cal/g. • When a gas changes to a liquid, the heat change is the heat of condensation (Hcond). • The value of Hvapor is the same as the value of Hcond. Chapter 13

29. Solid–Gas Phase Changes • Some substances convert directly between the solid and gas phases. • The process of a solid changing directly to a gas is called sublimation. • The process of a gas changing directly to a solid is called deposition. • Carbon dioxide (CO2) and iodine (I2) are two common substances that undergo sublimation–deposition phase changes. Chapter 13

30. Temperature–Energy Graphs We can graph the amount of energy required to change the temperature and physical state of a substance. • The heating curve for water is shown here. • As energy is added, the temperature increases and changes the physical state. Chapter 13

31. Energy from Heating Curves • We can use the heating curve and heat values for water to calculate how much energy is required to change the temperature of a sample of water. • The amount of energy required to raise the temperature of a substance is calculated using the following formula: heat = (specific heat) x(change in temperature) x(mass of sample) • The amount of energy required to change the state of a substance is calculated as follows: heat = (Hxxx) x(mass) Chapter 13

32. Energy Calculation • How much energy is required to raise 25.5 g of ice at –5.0 C to steam at 100.0 C? • Looking at the heating curve for water, there are four regions: • Heating of solid ice from –5.0 C to 0.0 C. • Melting of ice at 0.0 C. • Heating of liquid water from 0.0 C to 100.0 C. • Vaporization of water at 100.0 C. Chapter 13

33. Energy Calculation, Continued • The total energy is the sum of the energy in Steps 1 through 4. Calculate the energy for each step. • (25.5g)x [0.0–(–5.0)]Cx(0.50 cal/g x C)= 64 cal 2. (25.5 g) x(80.0 cal/g)=2040cal 3. (25.5g)x[100.0–0.0)]Cx(1.00cal/g x C)=2550cal 4. (25.5 g)x(540 cal/g)=13,800cal • The total energy is: • 64 cal + 2040 cal + 2550 cal + 13800 cal = 18,500 cal. Chapter 13

34. Structure of Water Let’s start with the electron dot formula for water. • Water has a bent molecular shape and the bond angle is 104.5. • Water is a polar molecule that exhibits strong hydrogen bonding. Chapter 13

35. Properties of Ice • Water is one of the few substances that is less dense as a solid than as a liquid. • As water freezes, the hydrogen bonds organize the water molecules into a three-dimensional structure where the molecules are farther apart then in the liquid. • Liquid water has a density of 1.00 g/mL, while ice has a density of 0.917 g/mL. Chapter 13

36. Water Purification • In many areas, water has lots of dissolved minerals leading to high concentrations of ions. • This is referred to as hard water. • It is often not suitable for use in agriculture or drinking. • The water is purified in a water softener by exchanging the cations and anions for H+ ions and OH– ions. Chapter 13

37. Physical Properties of Water Water has unusual melting and boiling points, especially compared to the other hydrogen compounds of Group VIA/16. • This is due to hydrogen bonding that is present in water, but not present in H2S, H2Se, or H2Te. Chapter 13

38. Chemical Properties of Water • Water can undergo an electrolysis reaction to produce hydrogen and oxygen: 2 H2O(l) → 2 H2(g) + O2(g) • Water reacts with active metals to produce hydrogen and a metal hydroxide: 2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g) • Water reacts with metal oxides to produce a base: CaO(s) + H2O(l) → Ca(OH)2(aq) • Water reacts with nonmetal oxides to produce an acid: CO2(g) + H2O(l) → H2CO3(aq) Chapter 13

39. Reactions that Produce Water • Water is obtained as a product in several types of chemical reactions. • Combustion reactions: • 2 C2H2 (g) + 5 O2 (g) → 4 CO2 (g) + 2 H2O (g) • C2H5OH (g) + 3 O2 (g) → 2 CO2 (g) + 3 H2O (g) • Neutralization reactions: • H3PO4 (aq) + 3 LiOH (aq) → Li3PO4 (aq) + 3 H2O (l) • Dehydration reactions: • Water is driven off from a hydrate by heating. Chapter 13

40. Hydrates • A hydrate is a crystalline ionic compound that contains water: CuSO4  5 H2O • The dot indicates that water molecules are bonded directly to each unit of hydrate. • Heating a hydrate drives off the water and produces an anhydrouscompound (without water). CuSO4  5 H2O(s) →CuSO4(s) + 5 H2O(l) heat Chapter 13

41. Critical Thinking: Bottled Water • Bottled water is a very large industry. • Water from our home faucet is tap water. • Tap water that has been processed by distillation or deionization is purified water. • Spring water is obtained from natural underground springs. • In most cases, bottled water has lower purity standards than tap water. Chapter 13

42. Chapter Summary • There are three types of intermolecular bonds: • Dispersion forces • Dipole forces • Hydrogen bonds • Dispersion forces are the weakest, and hydrogen bonds are the strongest. • These intermolecular attractions affect the physical properties of substances. Chapter 13

43. Chapter Summary, Continued • There are five properties of liquids that are affected by intermolecular bonds: • Vapor pressure decreases as intermolecular force increases. • Boiling point increases as intermolecular force increases. • Viscosity increases as intermolecular force increases. • Surface tension increases as intermolecular force increases. Chapter 13

44. Chapter Summary, Continued • There are three types of crystalline solids: • Ionic solids • Molecular solids • Metallic solids • Energy is required to raise the temperature of a substance. • Water displays many unique properties due to the presence of strong hydrogen bonds. Chapter 13