Notes on Solids and Liquids

Transcription

1 THE LIQUID STATE Notes on Solids and Liquids Why do liquids tend to bead up when on a solid surface? The effect of uneven pull on surface molecules draws them into the body of the liquid causing droplet to assume shape that has the minimum surface area a sphere. What is surface tension? The resistance of a liquid to an increase in its surface area. Liquids with high intermolecular forces of attraction tend high surface tension. yt cl= &v=45yabrnryxk&x yt ts= &feature=player_detailpage What is capillary action? The spontaneous rising of a liquid in a narrow tube. What two forces are responsible for capillary action? a) cohesive forces intermolecular forces of attraction among molecules of a liquid (these decrease the surface area) b) adhesive forces intermolecular forces of attraction between the liquid molecules and their container (these tend to in surface area) Concave vs convex meniscus in liquids: concave: adhesive forces > cohesive forces; meniscus is lower in the middle convex: cohesive forces > adhesive forces; meniscus is higher in the middle 1

2 What is viscosity? a measure of a liquids resistance to flow the stronger the intermolecular forces of attraction the greater the viscosity the more complex and larger the molecule the greater the viscosity Problem: Which would have a higher surface tension, H 2 O or C 6 H 14? Why? Would the shape of the H 2 O meniscus in a glass tube be the same or different than C 6 H 14? 2

3 SOLIDS crystalline vs amorphous highly regular disordered atomic atomic arrangements arrangements Unit Cell the smallest repeating unit of the lattice. Three types of cubic unit cells are: 1) simple cubic 2) body centered cubic 3) face centered cubic 3

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5 Types of Solids Ionic form electrolytes when dissolved in water Molecular do not form electrolytes when dissolved in water Atomic contain atoms of only one kind of element that are covalently bonded to each other. These are also called network solids The properties of a solid are determined primarily by the nature of the forces that hold the solid together. Structure and Bonding in Metals Properties of metals: high thermal and electrical conductivity, malleability, ductility The proposed model that describes the structure of metals and is able to explain these properties assumes the atoms of metals are uniform hard spheres that are packed to best utilize available space (close packing). Closest Packing Model See text diagrams hexagonal close packing (hcp) has every other layer being spatially equivalent ("abab..."). Examples include Mg, Zn, Cd, Co cubic close packing (ccp) or a face centered cubic structure has every third layer being spatially equivalent ("abcabc...") Examples include Ag, Al, Ni, Pb, Pt 5

6 The coordination number in the above two is 12 (12 equivalent nearest neighbors). An exception is the body centered cubic structure (bcc). In this structure, spheres are not close packed. The coordination number is 8 (Fe and alkali metals). Knowing the net number of spheres (atoms) in a particular unit cell is important for many applications involving solids. 6

7 e.g. ccp = face centered cube = cubic = bcc = body centered cube = 7

8 Of the most reliable nethods for determining Avogadro s number involves combining the density of a crystal, its formula weight, and its interatomic spacing (as determined by X rays via Bragg equation). Let s consider a problem. Problem: Chromium has a body centered cubic structure. Its density is 7.19 g/cm 3, and the length of the edge of a unit cell is pm. Calculate a value for the Avogadro constant. First, let s get the mass of the unit cell frrom its volume and density: Mass = Volume X Density = (2.884 X 10 8 cm) 3 X 7.19 g/cm 3 The unit cell conjsists of two atoms. The mass of two chromium atoms is therefore X g Hence the mass of 2 mol of chromium atoms is X g N A g = 2(molar mass Cr) = 2 X g Hence Avogadro s constant N A = 6.03 X

9 BONDING IN METALS Properties of metals indicate that the bonding in most metals is both strong and nondirectional. electron sea model metallic kernels (which are defined as the nucleus plus all but valence electrons) are arranged in crystal lattice structure and valence electrons are shared by entire crystal. Metals are good conductors because the electrons are delocalized and relatively free to move ("sea of mobile electrons"). Metals are melleable and ductile because deforming the solid does not change the environment immediately surrounding each metal kernel. band model in a metal crystal, electrons are assumed to travel in the metal crystal in molecular orbitals. Molecular orbitals are formed from an atom s valence atomic orbitals consisting a of bonding robitals (high probability of finding electrons) and nonbonding orbitals (low probality of finding electrons). In a metal the number of MO's are closely spaced forming a continuum called bands. No need to concern ourselves too much with this. Alloy = a substance that contains a mixture of elements and has metallic properties. Pure metals and alloys have different physical properties. An interstitial alloy is formed when holes in the close packed metal structure are occupied by small atoms (in high carbon steels the iron holes are occupied by carbon). A substitutional alloy some metal atoms are replaced by others of similar size and electronegativities The atoms must have similar atomic radii and the elements must have similar bonding characteristics e.g. Cu + Zn = brass. 9

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11 NETWORK SOLIDS an atomic solid containing strong directional covalent bonds. example: carbon allotrope forms are diamond and graphite sp 3 tetrahedron sp 2 trigonal planar conductor (pi molecular orbitals de localized electrons) nonconductor (large gap between empty and filled levels) 11

12 Silica empirical formula is SiO 2 CO 2 vs SiO 2 predicted Lewis structure: predicted Lewis structure: CO 2 can exist SiO 2 cannot exist! Why? 12

13 SEMICONDUCTORS Silicon forms a three dimensional network similar to in geometry to diamond. One difference is that silicon is a semiconductor. Thi can be explained by the fact that the energy gap between filled and empty molecular orbitals is smaller in silicon compared to diamond and a few electrons in silicon can cross the gap at room temperature. If the temperaure is increased, more energy is available to excite electrons into the conduction bands and thus increase conductivity. The conductivity of silicon can be increased a small fraction of silicon atoms is replaced by arsenic atoms, which have one more valence electron compared to silicon atoms (called doping ). This produces an n type semiconductor (negative charge carrying), one in which the conductivity of the substance is increased by doping it with atoms having more valence electrons compared to atoms in the host crystal. Doping silicon with an element that has one less valence electron (such as boron) converts silicon into a p type semiconductor (positive charge carrying) material. In this case an electron vacancy or hole is created. As an electron fills this hole, it leaves a new hole, and this process is repeated so the hole advances through the crystal in a direction opposite to movement of the electrons jumping to fill the hole. Junctions between n doped and p doped materials can be used to control electron flow, and therby are the basis of modern electronics. 13

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15 Photovoltaic or solar cells use semiconductors such as silicon. The energy needed to for an electron to cross the gap in silicon is 2.1 X J/atom. Use your knowledge of atomic structure to calculate the wavelength of light needed to accomplish this transition. What part of the electromagnetic spectrum is this light energy? 15

16 The process of converting light (photons) to electricity (voltage) is called the photovoltaic (PV) effect. The PV effect was first discovered by the French physicist Edmund Becquerel in 1839 using copper oxide in an electrolyte. To create the PV effect, radiation from the sun ('sunlight') hits a photovoltaic cell. These cells are made up of two layers of semi conducting material, typically silicon, that have been chemically treated. The industry refers to these layers as P and N. The boundary between P and N acts as a diode allowing electrons to move from N to P, but not from P to N. When photons with sufficient energy hit the cell, they cause electrons to move (from N to P only) causing excess electrons in the N layer and a shortage in the P layer. 16

17 Solids and Liquids 15.notebook MOLECULAR SOLIDS a solid composed of neutral molecules at the lattice points these substances are characterized by strong covalent bonding within the molecules (intramolecular) but relatively weak forces between the molecules (weak intermolecular forces). examples: H2O(s), CO2(s), P4(s), S8(s), I2(s) Recall the stronger the bond, the shorter the distance between atoms. The stronger the intermolecular force, the shorter the distance between molecules. 17

18 IONIC SOLIDS stable, high melting point substances held together by the strong electrostatic forces that exist between opposite charged ions. The structure of most binary ionic solids can be explained by the closest packing of spheres. Anions, which are usually larg than cations with which they combine, are packed in either an hcp or ccp arrangement. Cations fill holes within anions. KEY IDEA: the packing arrangement is done is such a way as to minimize anion anion and cation cation repulsions. The nature o the holes depends on the ratio of the anion to cation size. Trigonal holes are smallest, followed by tetrahedral with octahedral being the largest. 18

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20 Problem: Based on their properties, classify each of the following substances as to the type of solid it forms: a) Fe b) C 2 H 6 c) CaCl 2 d) graphite e) F 2 20

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22 PHASE CHANGES Evaporation: phase change of liquid to gas occuring only at the surface of a liquid below the boiling point temperature; only those molecules with above average KE at the surfave of the liquid may have enough energy to overcome IMFA and escape to the vapor phase. The higher the temperature the greater the percentage of molecules with the minimum KE required to overcome IMFA. 22

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24 Vaporization (boiling): phase change of liquid to gas occuring throughout the liquid at the boiling point of the liquid Sublimation: phase change of solid directly to gas Heat of Vaporization: the heat energy required to vaporize a unit amount of substance at its normal boiling point Heat of Fusion: the heat energy required to melt a unit amount of a substance at its melting point Vapor Pressure: the pressure exerted by the vapor molecules when a liquid in a closed system contains an equilibrium between evaporation and condensation (opposing rates equal) 24

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27 Volatile: a property of a liquid which means the liquid can easily evaporate; in a volatile liquidf the IMFA must be relatively weak The vapor pressure of a liquid is affected by two main factors: a) molecular weight at a given temperature heavy molecules have slower velocities and thus much smaller tendency to escape from the liquid surface lower vapor pressures b) intermolecular forces the stronger the forces the lower the vapor pressure The vapor pressure of a liquid increases with temperature (nonlinear increase). A plot of vapor pressure vs temperature yields "vapor pressure curves": 27

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30 From the vapor pressure curve we can derice the Clausius Claperyon Equation which interrelates vapor pressure, temperature and enthalpy of vaporization of a liquid. V = V gas V liq = V gas = PV = nrt = RT if n = 1 slope = = = integrate (calculus) note: if P1 = 1 atm, then T1 = normal boiling point The beauty of this equation is that you can predict the vapor pressure curve without actually measuring it. 30

31 Problem: The vapor pressure of 1 propanol at 14.7oC is 10.0 torr. The heat of vaporization is 47.2 kj/mol. Calculate the vapor pressure of 1 propanol at 52.8oC. 31

32 HEATING CURVE OF A PURE SUBSTANCE note: the temperature of a substance remains constant during a phase change (heat energy converted to PE) the temperature rises when heat is input while the substance is in one phase (heat energy converted to KE) for water, heat of fusion: 6.0 kj/mol heat of vaporization: 41.2 kj/mol heat capacity of ice: 2.1 J/goC heat capacity of water: 4.2 J/goC heat capacity of steam: 1.8 J/goC heating curve of water: 32

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35 normal melting point: the temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure is 1 atm. normal boiling point: the temperature at which the vapor pressure of the liquid is exactly 1 atm. Problem: How much energy does it take to convert 130. g of ice at 40.0oC to steam at 160oC? 35

36 PHASE DIAGRAMS A phase diagram shows the relationship among the temperature, pressure and phase of a substance. It describes events an conditions in a closed system. The phase diagram for water is illustrated below: 36

37 The diagram is interpreted as follows: BD is the vapor pressure curve (liquid vapor equilibrium) for the liquid phase. At 760 mm Hg pressure and 100oC,point F, is the boiling point (Tb). As you go up the BD line the boiling point temperature increases as the pressure increases. BC is the melting point curve (solid liquid equilibrium). The freezing point occurs at 0oC and 760 mm Hg at point E (Tm). AB is the vapor pressure curve for the solid state (solid vapor equilibrium). At point B the solid, liquid, and vapor phases of water may all exist in equilibrium. This point is referred to as the triple point of water. The triple point of a substance is the only temperature at which all three phases of a pure substance can exist in equilibrium with one another (in a system containing only the pure substance). For pure water this temperature is 0.01oC (at a pressure of 4.5 mm Hg). note: critical temperature: the temperature above which the vapor cannot be liquified no matter what pressure is applied (for water = 374oC) critical pressure: the pressure required to produce liquefaction at the critical temperature (for water = 218 atm). (The critical temperature and critical pressure are known as the critical point. Of special interest in the phase diagram for water is the BC curve. Notice that if the pressure is increased, the temperature must be decreased if equilibrium is to be maintained. Thus, increasing the external pressure actually decreases the melting temperature. This behavior is uncommon among most substances and is partially responsible for the movement of glaciers and ice skating. Note that the BC line for most substances would be shifted away from the y axis showing that an increase in pressure increases the melting point temperature. 37

38 Another way to analyze this is as followsis that along any curve the slope is given as: slope = S/ V liq gas S = +/ V = + therefore slope > 0 (= +) sol gas S = +/ V = + therefore slope > 0 (= +) sol liq S = +/ V = therefore slope < 0 (= ) Problem: Why does the solid/liquid line for water in a phase diagram have a negative slope? Problem: What phase changes does water undergo as the pressure changes while the temperature is held constant at 12oC? 38

39 Phase Diagram for CO 2 Note: The solid/liquid line has a positive slope since solid CO 2 is more dense than liquid CO 2. At 1 atm pressure solid CO 2 sublimes at 78oC (dry ice). 39

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