Chapter 11: Liquids and Solids and Intermolecular Forces. States of Matter: Liquids and Solids. The Phases of Matter

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1 Chapter 11: Liquids and Solids and Intermolecular Forces Principles of Chemistry: A Molecular Approach,1 st Ed. Nivaldo Tro Dr. Azra Ghumman Memorial University of Newfoundland States of Matter: Liquids and Solids 11.2 A molecular comparison of Gases, Liquids, and Solids 11.3 Intermolecular Forces: The forces that hold condensed phase together 11.5 Vapourization and Vapour Pressure (Excluding the Clausius-Clapeyron Equation) 11.6 Sublimation and Fusion 11.7 Heating Curve for Water Crystalline Solids: The Fundamental Types 2 The Phases of Matter Three states of matter-gas, Liquid and Solid 3

2 A Molecular comparison of Gases, Liquids, and Solids Gases are compressible fluids. Molecules are widely separated and are in constant random motion throughout mostly empty space. Assume the shape of container Negligible forces of interaction Negligible molecular size compared with total volume (K.T) Liquids- incompressible fluids. Molecules are more closely spaced than gases Assume the shape of container 4 Solids Solids- Molecules or ions are in close contact and only oscillate or vibrate about fixed positions definite shape, incompressible and rigid. Crystalline: particles are arranged in an orderly geometric pattern e.g. salt and diamonds Amorphous: particles have no regular geometric pattern over a long range e.g. plastic and glass 5 Changes between Phases Phase transition: Change of a substance from one physical state to another Caused by change in temperature and or pressure melting vapourization freezing condensation (LPG) Increase in pressure favour the dense state i.e. Condensation of gases 6

3 11.3 Intermolecular Forces Intermolecular Forces(IMF)- The interactive forces between the molecules hold the condense phase together Normally weak forces than chemical bonds Determines Physical properties of liquids and solids Three types of forces are known to exist between neutral molecules Dispersion forces (Fritz W.London) van-der Waals Dipole-dipole forces forces Hydrogen bonding 7 Intermolecular Forces Intermolecular attractions are due to attractive forces between opposite charges (q 1 and q 2 ) separated by a distance r Coulomb s Law larger the charge, stronger the attraction longer the distance, weaker attraction H-bonding especially strong 8 Intermolecular Forces London dispersion forces- Result from temporary dipole in the molecules due to unequal electron distribution. Present in all substances Dipole Dipole attractions- Permanent polarity in the molecules due to their structure leads to especially strong dipole dipole attraction Hydrogen bonding- Occurs in substances containing hydrogen atoms bonded to certain very electronegative atoms (O, N, F). Stronger than van der Waals forces (Dispersion and dipole-dipole) 9

4 Instantaneous Dipole in He 10 Dispersion Forces The magnitude of the dispersion forces depends on several factors. 1. Polarizability: the tendency for charge separation to occur in a molecule. The larger a molecule, more easily it can be distorted to give an instantaneous dipole Polarizability increases with molar mass resulting in strong dispersion forces 2. Shape of the molecule- Elongated and branched molecules are easily polarizeable compared to compact structures more surface-to-surface contact = larger induced dipole stronger attraction 11 Sample Problem n-pentane and neopentane (C 5 H 12 ) has same molar mass g mol -1. But boiling point of n-pentane is higher than neopentane. Why? 12

5 Effect of Molecular Size on Size of Dispersion Force Effect of size on the boiling points 13 Dipole Dipole Attractions Dipole-dipole force- is an attractive intermolecular force resulting from the tendency of polar molecules to align themselves positive end to the negative end Miscibility Like dissolve like The permanent dipole adds to the attractive forces between the molecules. Note- All molecules have dispersion forces Strong IMF s increase the bp and mp of the compounds H Cl H Cl 14 Practice problem Choose the substance in each pair with the highest boiling point. b) 15

6 Hydrogen Bonding Hydrogen bonding is a force that exists between a hydrogen atom covalently bonded to a very electronegative atom, X (F, O and N) and a lone pair of electrons on a very electronegative atom,y in the vicinity. X-H----:Yone of the following three structures must be present Only N, O, and F are electronegative enough to leave the hydrogen nucleus exposed Hydrogen bond is not a chemical bond Type of strong dipole-dipole force strongest of three intermolecular forces 16 Boiling point versus molecular weight for Hydrides If London dispersion forces are dominant, The boiling points of hydrides should increase with increasing molar mass but H 2 O, NH 3 and HF do not follow this trend. Why? 17 Hydrogen Bonding in water Properties of water and IMF The e - s in the O-H bond are drawn to the O atom, leaving the dense positive charge of the hydrogen nucleus exposed. The strong attraction of this exposed nucleus for the lone pair on an adjacent molecule accounts for the strong hydrogen bond

7 H- Bonding in Water and properties On freezing the molecules are arranged in open hexagonal pattern that gives ice lower density at 0 C (0.917g.cm -3 ) than water(1.000 g. cm -3 ). Moderate temperatures near lakes Water expands on freezing (equatic life under ice) H-bonding in Biological molecules A similar mechanism explains the attractions in HF and NH 3 19 Identifying intermolecular forces 1. What kind of intermolecular forces are expected in the following substances? Arrange them in order of increasing boiling points. CH 4, CHCl 3, C 2 H 6 and CH 2 CH 2 OH. 2. Choose the substance in each pair that is a liquid at room temperature (the other is a gas). 20 Ion Dipole Attraction Ion Dipole forces-occurs when an ionic compound is mixed with a polar compound ions from an ionic compound are attracted to the dipole of polar molecules. E.g. Solubility of ionic compounds in water. 21

8 22 Phase transition revisted Phase Transition:A change of a substance from one state to another is called a change of state or phase transition Heat for phase transition q= n H H sub = H fus + H sub = H fus + H vap H vap Vapourization and Vapour Pressure The average kinetic energy is proportional to the temperature. Vapourization- Some molecules with more kinetic energy than the average, overcome the attractive forces and escape the liquid to gas phase as vapour Rate of vapourization depends; Temperature Surface area Strength of IMF Condensation: Opposite of vapourization 24

9 Distribution of Thermal Energy The higher the temperature, the greater the average energy of the collection of the molecules 25 Effect of IMF on Vapourization Volatile substances -Liquids and solids that evaporate easily at normal temperatures Relatively high vapour pressure Weak IMF e.g., gasoline, fingernail polish remover Nonvolatile - Liquids that that do not vapourize easily Low vapour pressure Strong IMF e.g. H 2 O compared to acetone or motor oil Vaporization is an endothermic process. Why? Condensation is an exothermic process. Why? 26 Heat of Vaporization Heat of vaporization, H vap -The amount of heat energy required to vaporize one mole of the liquid sometimes called the enthalpy of vaporization always endothermic, H vap is +ve somewhat temperature dependent e.g. for water H 2 O(l) H 2 O(g) H vap = 44.7 kj mol -1 (25 C) = 40.7 kj mol -1 (100 C) H cond. = - H vap = kj mol-1 (100 C) Heats of vapourization of several liquids (Table 11.7) 27

10 Heat of Vapourization Using heat of vapourization in calculation (recall stoichiometry of H section 6.5) Calculate the mass of water that can be vaporized with 155 kj of heat at 100 C. Solution- 28 Dynamic Equilibrium Dynamic Equilibrium-The condition at which two opposite processes, evaporation and condensation occur at same rate rate of vapourization = rate of condensation 29 Vapour Pressure Vapor pressure-the pressure exerted by the vapor when it is in dynamic equilibrium with its liquid The weaker the attractive forces between the molecules, the higher the vapor pressure. The higher the vapor pressure, the more volatile is the liquid 30

11 Vapor Liquid Dynamic Equilibrium Le Chatelier s Principle-When a system in dynamic equilibrium is disturbed the system responds so as to minimize the disturbance and return to a state of equilibrium 31 Boiling Point Boiling Point- The temperature at which the vapour pressure of a liquid equals the external pressure ( by the atmosphere) Boiling point varies with external pressure e.g. bp of H 2 O =100 C at 1.00 atm = 95 C at 0.83 atm Normal boiling point -The temperature at which vapour pressure of a liquid equals 1.00 atm (760 mmhg = 101.3kPa) Normal bp of H 2 O at 1.00 atm = 100 C The normal bp is related to vapor pressure and is lowest for liquids with the weakest intermolecular forces. When intermolecular forces are weak, little energy is required to overcome them resulting in low boiling points 32 Vapour Pressure of Several liquids at different temperature Which liquid has the strongest intermolecular attractions? Vapour pressure increases with increasing temperature 33

12 Heating Curve of a Liquid As you heat a liquid, its temperature increases linearly until it reaches the boiling point. q = mass C s T During boiling the temperature remains constant 34 Critical Point: Another phase of matter Supercritical fluids- have properties of both gas and liquid states Critical point-at which gas and liquid phases mingles; the meniscus between gas and liquid disappears Critical Temperature(T c )-The temperature at which this transition occurs Critical Pressure(P c )- The pressure at which this transition occurs Sublimation and fusion Sublimation- The phase transition from solid to gas Deposition- The phase transition from gas to solid The opposite of sublimation. The solid and vapor phases exist in dynamic equilibrium in a closed container. 36

13 Melting or Fusion Melting point (fusion)- The temperature at which a crystalline solid changes to a liquid (endothermic) E.g. H 2 O(s) H 2 O(l) H fus = 6.02kJ /mol Freezing- Freezing Point-The temperature at which pure liquid changes to a crystalline solid or freezes (Exothermic) opposite to melting H cryst = - H fus H 2 O(l) H 2 O(s) H cryst = kJ /mol can be used to identify the substance Heat of fusion, H fus The amount of heat energy required to melt one mole of the solid H fus is always +ve somewhat temperature dependent H sub = H fus + H vap (Heat of fusion for several substances See Table11.8) 37 Heating curve for one mole of Water 38 Crystalline solids: Fundamental Types There are four types of solids. Molecular solids (Van der Waals forces) composite particles are molecules. E.g. solid water (ice), and solid carbon dioxide (dry ice) Ionic solids(ionic bond)- composite particles are ions Atomic solids- composite particles are atoms Metallic (Metallic bond) e - - sea model Nonbonding atomic solids are held together by dispersion forces. Network covalent atomic solids are held together by covalent bonds 39

14 Classification of Solids A molecular solid- composed of molecules held together by intermolecular forces dispersion forces, dipole dipole attractions, and H-bonds Many solids are of this type e.g. solid neon, ice, solid sulfur (S 8 ) and solid carbon dioxide (dry ice) Weak IMF leads to low melting points. generally <300 C An ionic solid consists of cations and anions held together by electrical attraction of opposite charges (ionic bond) e.g. NaCl Strength of Ionic bonds depends on The magnitude of ionic charge Size of the ions the m.p. of NaCl 801 C compared to MgO is 2800 C Lattice energy(energy needed to separate a crystal into isolated ions in gaseous form) 40 Atomic Solids Nonbonding Atomic Solids- A group that only consists of noble gases in solid form e.g. Xe held together by weak dispersion forces very low melting points which increases uniformly with molar mass e.g. mp of Ar is -189 C and of Xe is -112 C tend to arrange atoms in closest-packed structure maximizes attractive forces and minimizes energy A metallic solid consists of positive cores of atoms held together by a surrounding sea of electrons (metallic bonding). positively charged atomic cores are surrounded by delocalized electrons e.g. Fe, Cu, Au etc 41 Network Covalent Solids A covalent network solid consists of atoms held together in large networks or chains by covalent bonds. Examples include different forms of carbon, as diamond (3-D) or graphite (2-D sheets) asbestos (long chains), and silicon carbide. very high melting points generally >1000 C The dimensionality of the network affects other physical properties. 42

15 The crystalline Structure of Diamond and Graphite Covalent bonds extend throughout a crystalline solid 3-D Network Two-dimensional structure 43 The Structure and Properties of Diamond Diamond(bonding) -Each C is sp 3 hybridized and covalently bonded to four other C atoms tetrahedrally (very stable 3 D structure) each crystal is one giant molecule held together by covalent bonds Propertiesvery high mp ~3800 C very rigid (directionality of the covalent bonds) Very hard (used as abrasives) electrical insulator & nonreactive 44 The Structure and properties of Graphite Graphite(bonding)- Each C is sp 2 hybridized and bonded to three other C atoms in trigonal planar geometry (Two- Dimensional Network) has three bonds and one bond. form flat sheets fused together each sheet is a giant molecule The sheets are then stacked and held together by dispersion forces. Properties- slippery feel, used as lubricants high mp electrical conductor due to delocalized pi electrons chemically nonreactive 45

16 Silicates Most common network covalent atomic solids ~90% of Earth s crust Basic compound silica (SiO 2 )- extended arrays of Si O Al substituted for Si aluminosilicates Quartz - most common crystalline form SiO 4 three-dimensional array of Si covalently bonded to four O in tetrahedral geometry impurities add color mp ~1600 C. very hard Glass is SiO 2 the amorphous form (Fig 11.49b) 46 Operational Skills Calculating the heat required for a phase change of a given mass of substance. Identifying intermolecular forces and describing physical properties in terms of IMF s Identifying types of solids. Determining the relative melting points based on types of solids. 47

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