Intermolecular Forces, Liquids, and Solids

Transcription

1 Slide 1 / 92 Intermolecular Forces, Liquids, and Solids Intermolecular Forces Slide 2 / 92 Intermolecular forces are the piece we need to add to the puzzle to explain the world around us. We first explained atoms, and how to build up the periodic table from quantum numbers. Then we explained how atoms combine to form molecules: the most common way we find most atoms in nature. Now, we're going to use intermolecular forces to combine molecules to create the common states of matter. Without intermolecular forces, we wouldn't have tables, lakes, wall...or even our bodies. Intermolecular forces shape our world. States of Matter Slide 3 / 92 While there are many states of matter, the three common states that dominate our world are gases, liquids and solids. These are the states of matter we'll be studying. We won't be discussing more exotic states such as plasma, nuclear matter, etc.

2 States of Matter Slide 4 / 92 The fundamental differences between states of matter is: the distance between particles the particles' freedom to move cool or increase pressure heat or decrease pressure cool heat Gas Liquid rystalline solid Particles are far apart, total freedom, much of empty space, total disorder disorder, freedom, free to move relative to each other, close together orderd arrangement, particles are in fixed positions, close together haracteristics of the States of Matter Slide 5 / 92 Gas Liquid Solid ssumes the shape of its container Expands to the volume of its container Is compressible Flows easily Diffusion within a gas is rapid ssumes the shape of the part of a container it occupies Does not expand to the volume of its container Is virtually incompressible Flows easily Diffusion within a liquid is slow Retains its own shape, regardless of container Does not expand to the volume of its container Is virtually incompressible Does not flow Diffusion within a solid is very very slow ondensed Phases Slide 6 / 92 In the solid and liquid states particles are closer together, we refer to them as condensed phases. cool or increase pressure heat or decrease pressure cool heat Gas Liquid rystalline solid Particles are far apart, total freedom, much of empty space, total disorder disorder, freedom, free to move relative to each other, close together orderd arrangement, particles are in fixed positions, close together

3 1 Which of the following is not a type of solid? Slide 7 / 92 ionic molecular covalent-network D supercritical E metallic 2 Which of the below is a characteristic of a gas? Slide 8 / 92 D E It fills only a portion of its container. Its molecules are in relatively rigid positions. It takes on the shape of its container. It is not compressible. Diffusion is very slow within it. 3 Which of the below is a characteristic of a liquid? Slide 9 / 92 D E It fills only a portion of its container. Its molecules are in relatively rigid positions. It takes on the shape of its container. It is compressible. Diffusion is very rapid within it.

4 4 Which of the below is a characteristic of a solid? Slide 10 / 92 D E It fills all of its container. Its molecules are in relatively rigid positions. It takes on the shape of its container. It is compressible. Diffusion is very rapid within it. 5 Together, liquids and solids, constitute phases of matter. Slide 11 / 92 the compressible the fluid the condensed D all of the above E the disordered States of Matter Slide 12 / 92 The state of a substance at a particular temperature and pressure depends on two opposing properties: Intermolecular Forces: the strength of the attractions between the particles, which pulls them together the kinetic energy of the particles, which pulls them apart Without intermolecular forces, all molecules would be ideal gases...there would be no liquids or solids.

5 Intermolecular Forces & oiling Points oiling represents a transition from a liquid to a gas. Slide 13 / 92 To make that transition, molecules in the liquid must break free of the intermolecular forces that bind them. The kinetic energy of the molecules is proportional to the temperature: as temperature rises, so does kinetic energy. The temperature where the molecules' energy overcomes intermolecular forces is called the boiling point. The boiling point is a measure of the strength of the intermolecular forces: the higher the boiling point - the stronger the intermolecular forces. Intermolecular Forces Slide 14 / 92 H l H l ovalent bond (strong) Intermolecular attraction ( week) The attractions between molecules, intermolecular forces, are not nearly as strong as the intramolecular attractions that hold compounds together. They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. Intermolecular Forces Slide 15 / 92 There are three types of Intermolecular Forces: they are sometimes called van der Waals Forces Dipole-dipole interactions Hydrogen bonding London dispersion forces

6 - + Dipole-Dipole Interactions Slide 16 / 92 Molecules that have permanent dipoles are attracted to each other. The interaction between any two like charges is repulsive (black) + - The positive end of one is attracted to the negative end of the other and vice-versa These forces are only important when the molecules are close to each other The interaction between any two opposite charges is attractive ( red) + Dipole-Dipole Interactions Slide 17 / 92 The polarity of a molecule is measured by its dipole moment, m. The more polar the molecule, the greater its dipole moment. The more polar the molecule, the higher its boiling point. That's because the attraction between the dipoles holds the molecules together, not letting them boil away. Substance Molecular Dipole oiling Weight (amu) Moment u(d) Point(K) cetonitrile, H3N cetaldehyde, H3HO Methyl chloride, H3l Dimethyl ether, H3OH Propane, H3H2H Which of the below molecules will have the highest boiling point? Slide 18 / 92 H 3H 2H 3 H 3OH 3 Substance Molecular Dipole Wt. Moment H 3l D H 3HO H3H2H E H 3N H3OH H3l H3HO H3N

7 7 Which of the below molecules will have the lowest boiling point? Slide 19 / 92 H 3H 2H 3 H 3OH 3 H 3l D H 3HO E H 3N Substance Molecular Wt. H3H2H H3OH H3l Dipole Moment H3HO H3N London Dispersion Forces Slide 20 / 92 London Dispersion Forces occur between all molecules. They result from the fact that electrons are in constant motion and sometimes are the same side of the molecule. When they are on one side, the molecule is polarized: one side is negative and the other is positive; the molecule acts like a dipole. That creates an electric field that oppositely polarizes nearby molecules...leading to an attraction. Let's see how that works using Helium as an example. London Dispersion Forces Slide 21 / 92 While the electrons in helium repel each other, they occasionally wind up on the same side of the atom. t that instant, the helium atom is polar, with an excess of electrons on one side and a shortage on the other. nother helium atom nearby, has a dipole induced in it, as the electrons on the left side of the first atom repel the electrons in the second. London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. δ - e- 2+ e- Helium atom δ + electrostatic attractio e- e e- e- Helium atom Helium 1 atom δ - δ + δ - δ +

8 London Dispersion Forces Slide 22 / 92 These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability. The larger the molecule, the more polarizable it is...and the stronger the London Dispersion Force. That means, the higher the molecular weight of a molecule, the more London Dispersion Force it experiences. London Dispersion Forces Slide 23 / 92 The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms and molecules have larger electron clouds which are easier to polarize. Molecular Halogen Weight ( amu) oiling Noble Point (K) gas Molecula rweight (amu) F He l Ne r r oiling point (K) I Kr Xe Which of the molecules below will have the highest boiling point? Slide 24 / 92 F 2 l 2 r 2 D I 2

9 9 What intermolecular force is responsible for ice being less dense than liquid water? Slide 25 / 92 London Dispersion Forces Dipole-Dipole Forces Ion-Dipole Forces D Hydrogen onding E Ionic onding 10 Intermolecular force(s) responsible for the fact that H 4 has the lowest boiling point in the set H 4, SiH 4, GeH 4, SnH 4 is/are. Slide 26 / 92 Hydrogen onding Dipole-Dipole Interactions London Dispersion Forces D Mainly H2 onding with some dipole-dipole interactions E mainly London Dispersion Forces with dipole-dipole interactions 11 Which of the below molecules will have the lowest boiling point? Slide 27 / 92 F 2 l 2 r 2 D I 2

10 12 Which of the below molecules will have the highest boiling point? Slide 28 / 92 D E He Ne r Kr Xe 13 Which of the below molecules will have the lowest boiling point? Slide 29 / 92 D E He Ne r Kr Xe Which Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces Slide 30 / 92 If two polar molecules are of comparable size, dipole-dipole interactions are the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties. If molecules are nonpolar, dispersion forces will dominate, since all molecules experience dispersion forces.

11 Hydrogen onding In these graphs, boiling points increase with the mass of the molecules...as we'd expect: larger dispersion forces due to larger masses. Slide 31 / 92 The nonpolar series (H 4 to SnH 4 ) follow the expected trend. The polar series follows the trend from H 2Te through H 2S, ut water defies the trend. It is the lightest in its series, but has the highest boliing point. Hydrogen onding Slide 32 / 92 H The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds. These powerful bonds occur only when H is bonded directly to N, O or F. H H O H O H H H H N H O H H H N H O H Hydrogen onding Slide 33 / 92 Hydrogen bonding arises in part from the high electronegativity and small radius of nitrogen, oxygen, and fluorine. When hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

12 Hydrogen onding Ice is the only solid that floats in its liquid form. Slide 34 / 92 If it didn't, life on Earth would be very different. For instance, lakes would freeze from the bottom and fish couldn't survive winters. Hydrogen bonding creates the space in ice that explains its low density. Water-Ice transition simulation.htm 14 Which of the following molecules has hydrogen bonding as one of its IM forces? Slide 35 / 92 D E HF Hl Hr HI ll of the above. Ion-Dipole Interactions Slide 36 / 92 There is a fourth molecular force that will be important as we explore solutions later this year: Ion-dipole interactions are not considered a van der Waals force. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents. + _ nion- dipole attractions ation -dipole attractions

13 Summarizing Intermolecular Forces Slide 37 / 92 Interacting molecules or ions re polar molecules involved? NO re ions involved? YES re polar molecules and ions both present? NO YES re hydrogen atoms bonded to N,O,or F atoms? YES NO NO YES Dispersion forces only (induced dipoles) r, I2 Dipole-dipole forces H2S, H3l Hydrogen bonding H2O, NH3 Ion- dipole forces Nal in H2O Ionic bonding Nal, KI van der Waals forces 15 Which of the following molecules has London dispersion as its only IM force? Slide 38 / 92 PH 3 H 2S Hl D SiH 4 E None of the above. 16 How many of these substances will have dipole-dipole interactions? (How many are polar molecules?) H 2 O O 2 H 4 NH 3 Slide 39 / D 3 E 4

14 17 Which of the following molecules will have the highest boiling point? Slide 40 / 92 H 2 O O 2 H 4 D NH 3 18 Which liquid will have the lowest freezing point? Slide 41 / 92 D E pure H2O aq M glucose aq M sucrose aq M FeI3 aq M KF 19 Of the following diatomic molecules, which has the highest boiling point? Slide 42 / 92 N 2 r 2 H 2 D l 2 E O 2

15 20 Of the following diatomic molecules, which has the lowest boiling point? Slide 43 / 92 N 2 r 2 H 2 D l 2 E O 2 21 Which one of the following derivatives of methane (H 4 ) has the lowest boiling point? Slide 44 / 92 r 4 F 4 l 4 D I 4 22 Which one of the following derivatives of methane (H 4 ) has the highest boiling point? Slide 45 / 92 r 4 F 4 l 4 D I 4

16 23 For an aqueous solution of a nonvolatile compound, the vapor pressure will be, the boiling point will be, and the freezing point will be than pure water. Slide 46 / 92 D E lower, lower, lower lower, higher, lower lower, higher, higher higher, higher, lower higher, lower, higher Intermolecular Forces ffect Many Physical Properties Slide 47 / 92 The strength of the attractions between particles can greatly affect the properties of a substance or solution. Viscosity Resistance of a liquid to flow is called viscosity. Slide 48 / 92 It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature. Substance Formula Viscosity ( kg/m-s) Hexane H3H2H2H2H2H x 10-4 Heptane H3H2H2H2H2H2H x 10-4 Octane H3H2H2H2H2H2H2H x 10-4 Nonane H3H2H2H2H2H2H2H2H x 10-4 Decane H3H2H2H2H2H2H2H2H2H x

17 Surface Tension Slide 49 / 92 Surface tension results from the net inward force experienced by the molecules on the surface of a liquid. Phase hanges Slide 50 / 92 GS Vaporization ondensation Energy of system LIQUID Sublimation Deposition Melting Freezing SOLID Energy hanges ssociated with hanges of State Slide 51 / 92 hemical and physical changes are usually accompanied by changes in energy. When energy is put into the system, the process is called endothermic. When energy is released by the system, the process is called exothermic.

18 24 Which of the following is not a phase change? Slide 52 / 92 D vaporization effusion melting sublimation 25 Of the following, is an exothermic process? Slide 53 / 92 D E melting subliming freezing boiling all are exothermic 26 The first molecules to evaporate from a liquid are. Slide 54 / 92 Those with the lowest KE Those farthest from the surface of the liquid Those with the highest KE

19 27 The direct change of a substance from a solid to a gas is called. Slide 55 / 92 D boiling evaporation sublimation condensation oiling vs. Evaporation oiling and evaporation are two ways in which a liquid can vaporize into a gas. However, there are important distinctions between these processes. Slide 56 / 92 oiling occurs at a specific temperature, the boiling point (.P.) occurs throughout the entire liquid Evaporation occurs below the boiling point occurs only at the surface of a liquid achieved when atmospheric pressure equals vapor pressure (Patm = Pvap) Vapor Pressure Slide 57 / 92 t any temperature some molecules in a liquid have enough energy to escape. s the temperature rises, the fraction of molecules that have enough energy to escape increases.

20 Vapor Pressure Slide 58 / 92 s more molecules escape the liquid, the pressure they exert increases. The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate. Vapor Pressure Slide 59 / 92 The boiling point of a liquid is the temperature at which it s vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr. (K 760 mm Hg = 1 atm) 28 When an aqueous salt solution is compared to water, the salt solution will have Slide 60 / 92 D E a higher boiling point, a lower freezing point, and a lower vapor pressure a higher boiling point, a higher freezing point, and a lower vapor pressure. a higher boiling point, a higher freezing point, and a higher vapor pressure a lower boiling point, a lower freezing point, and a lower vapor pressure a lower boiling point, a higher freezing point, and a higher vapor pressure

21 29 What is the normal boiling point of ethanol? Slide 61 / D 78.3 E What is the boiling point ( in 0 ) of diethyl ether at 200 torr? Slide 62 / D What is the boiling point of water at 300 torr? Slide 63 / D 75 E 90

22 Vapor Pressure liquid will boil when its vapor pressure equals atmospheric pressure. pressure cooker works by increasing the "atmospheric" pressure inside it, so water will not boil at 100 o ; instead, it may be heated up to 120 o before turning to steam. Raising the cooking temperature cuts cooking time drastically. Slide 64 / 92 Vapor Pressure Slide 65 / 92 Vapor Pressure Slide 66 / 92 Recall that boiling occurs when P vap = P atm. Since atmospheric pressure is so low at high altitudes, (e.g. top of Mount Everest) water will boil at a much lower temperature than here at ergen Tech. P atm = 33 kpa on Mt. Everest P atm = kpa at sea level

23 32 It will take longer to hard-boil an egg Slide 67 / 92 on top of Mt. Everest here at ergen Tech cooking times are equal 33 volatile liquid is one that. Slide 68 / 92 D E is highly flammable is highly viscous is highly hydrogen-bonded is highly cohesive readily evaporates 34 If a liquid is sealed in a container and kept at constant temperature, how does its vapor pressure change over time? Slide 69 / 92 it rises at first, then remains constant it rises as first, then falls it remains constant

24 35 ased on this figure, the boiling point of diethyl ether under an external pressure of 1.32 atm is. Slide 70 / D 40 E 0 Phase Diagrams Slide 71 / 92 Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases. Volatility Slide 72 / 92 The more volatile a liquid: The more quickly it evaporates The higher its vapor pressure at a given temperature The weaker its Intermolecular Forces

25 Phase Diagrams Slide 73 / 92 The circled line is the liquid-vapor interface. It starts at the triple point (T), the point at which all three states are in equilibrium. Phase Diagrams Slide 74 / 92 It ends at the critical point (); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other. Phase Diagrams Slide 75 / 92 The circled line in the diagram below is the interface between liquid and solid. The melting point at each pressure can be found along this line.

26 Phase Diagrams elow the triple point the substance cannot exist in the liquid state. Slide 76 / 92 long the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line. Phase Diagram of Water Slide 77 / 92 Note the high critical temperature and critical pressure. These are due to the strong van der Waals forces between water molecules. The slope of the solid-liquid line is negative. This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid. Phase Diagram of arbon Dioxide arbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; O 2 sublimes at normal pressures. Slide 78 / 92 The low critical temperature and critical pressure for O 2 make supercritical O 2 a good solvent for extracting nonpolar substances (like caffeine).

27 Phase Diagram of H 2 O and O 2 Slide 79 / 92 The slope of the solid-liquid line is negative. If you increase the pressure at a given temperature, near the freezing point, the ice will melt. The Water slope of is the solid-liquid only known line is positive, as it is for most of substance with this behavior. the substances. if you increase the pressure at a given temperature near -55 0, the liquid will freeze. Note that, since the Triple point is at a high pressure, you will see O 2 only subliming under normal atmospheric condition. 36 The phase diagram of a substance is shown to the right. The region that corresponds to the solid phase is. Slide 80 / 92 w x y D x E x and y 37 In this phase diagram, the substance is a at 25 and 1.0 atm Slide 81 / 92 solid liquid gas D supercritical fluid E crystal

28 38 On a phase diagram, the melting point is the same as the Slide 82 / 92 triple point critical point freezing point D boiling point E vapor-pressure curve 39 On the phase diagram shown to the right, segment corresponds to the conditions of temperature and pressure under which the solid and the gas of the substance are in equilibrium Slide 83 / 92 D D D E Solids Slide 84 / 92 We can think of solids as falling into two groups rystalline, in which particles are in highly ordered arrangement. morphous, in which there is no particular order in the arrangement of particles.

29 Solids Slide 85 / 92 Some examples of amorphous solids are: rubber, glass, paraffin wax and cotton candy. rystalline solids include ionic compounds, metals and another group called covalent-network solids. rystalline solids are categorized by bonding type as shown on the next slide. Types of onding in rystalline Solids Slide 86 / In a saturated solution of salt water,. Slide 87 / 92 D E the rate of crystallization > the rate of solution the rate of solution > the rate of crystallization seed crystal addition may cause massive crystallization rate of solution = rate of crystallization addition of more water causes massive crystallization

30 ovalent-network Solids Slide 88 / 92 Diamonds are an example of a covalent-network solid, in which carbon atoms are covalently bonded to four other carbon atoms. They tend to be hard and have high melting points. ovalent-network Solids Slide 89 / 92 Graphite is another example of a covalent-network solid. Each carbon atom is covalently bonded to 3 others in layers of interconnected hexagonal rings. The layers are held together by weak dispersion forces. The layers slide easily across one another, so graphite is used as a lubricant as well as the "lead" in pencils. Metallic Solids Slide 90 / 92 Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. In metals, valence electrons are delocalized throughout the solid. This means that the "sea" of electrons moves freely around all the nuclei.

31 Properties of Metallic Solids Slide 91 / 92 The delocalized nature of the electrons in metals accounts for many physical properties. For example, metals are generally: good conductors of heat and electricity malleable ductile, (i.e. may be drawn into wires) Slide 92 / 92