CHAPTER 9: LIQUIDS AND SOLIDS

Transcription

1 CHAPTER 9: LIQUIDS AND SOLIDS Section 9.1 Liquid/Vapor Equilibrium Vaporization process in which a liquid vapor open container - evaporation continues until all liquid evaporates closed container 1) Liquid evaporate. 2) Vapor particles collect and Condense. 3) Eventually, Rate of Evaporation Rate of Condensation DYNAMIC EQUILIBRIUM

2 Vapor Pressure At equilibrium, # molecules/volume is constant. Pressure of gas over liquid is constant. As long as both liquid and vapor are present, the pressure exerted bythe vapor is independent of the volume of the container. Vapor Pressure is dependent on: a) Characteristics of liquid b) Temperature Important As long as BOTH liquid and solid are present, the vapor pressure will be constant.

3 If volume More liquid will evaporate Equilibrium Re-establish Vapor Pressure vs. Temperature In general, Vapor pressure For example, H 2 O Vapor Pressure of H 2 O Temp 24 mmhg 25 C 92 mmhg 50 C 760 mmhg 100 C as Temp. Higher the Temperature More Molecules Vaporize

4 A plot of pressure vs. temperature does not produce a straight line. This is not a direct relationship! Graphs of curves are often difficult to interpret. The solution to this problem: Graph manipulated variables of pressure and temperature that will represent a straight line.

5 Instead of P vs. T, graph ln P vs. 1/T Recall that the general equation of a straight line is y = mx + b (m = slope & b = y-intercept) Here, y = 1n P x = 1/T m = - H vap /R Therefore, ln P = - H vap 1 R T + b

6 If 2 different temps are evaluated: at T 2 : ln P 2 = - H vap R 1 T 2 + b at T 1 : ln P 1 = - H vap R 1 T 1 + b Clausius- Clapeyron Equation ln P 2 ln P 1 ln P 2 P 1 H R vap 1 T 2 1 T 1 where R = J/K mole

7 Boiling Point Vapor Pressure of Liquid EQUALS Pressure Above the Surface of the Liquid Normal Boiling Point = the temperature a liquid boiling at 1 atm of pressure above the liquid. The boiling point of any liquid can be lowered by reducing pressure above liquid. Varies with altitude. Critical Temperature the temperature above which the liquid state of a pure substance cannot exists regardless of the pressure.

8 Pressure (atm) Critical Pressure the pressure that be applied to cause the condensation of a pure liquid at the critical temperature. Section 9.2 Phase Diagrams Phase Diagram a graphical way to summarize the conditions under which the different states of a substance are stable. Water s Phase Diagram D Critical Point C Normal Freezing Point Liquid Normal Boiling Pt Solid.0060 B A Triple Point Vapor *Not to Scale Temperature ( o C)

9 Melting-Point Curve - Observe the solid/liquid states at different pressures. - Along the curve both phases are in equilibrium. - Special Note: When conditions indicate that a substance is in the liquid or solid state, the vapor of that substance is also present (in equilibrium). - The question one should ask is How much vapor is present? - Answer: It depends!!! Sublimation Transformation of a solid directly into a vapor. Melting Point The opposite process freezing.

10 Triple Point the point on a phase diagram representing the temperature and pressure at which three phases of a substance coexist in equilibrium. For example, water s triple point is 0.1 C, atm and all phases coexist.

11 Section 9.3 Molecular Substances: Intermolecular Interactions 1. Nonconductors (when pure) examples: I 2, C 3 H 8, C 2 H 5 OH Most water solutions are also nonconductors Some polar molecules form ions when they react with H 2 O conduct electricity For example: HF(g) H +1 (aq) + F -1 (aq) 2. Generally, molecular compounds are insoluble in water. 3. They have low melting & boiling points. Many are gases (N 2, O 2, ) Some are liquids with melting points <25 o C (like H 2 O, mp = 0 o C).

12 Some are solids with melting points <300 o C (like I 2, mp = 114 o C). The boiling point and melting point of molecular substances is directly related to the strength of their INTERMOLECULAR ATTRACTIVE FORCES among molecules. Intermolecular Forces 1. (London) Dispersion Forces Found in all molecular substances. involves a temporary or induced dipole. Consider the H 2 molecule Nonpolar bond equal sharing of electrons

13 For an instant, the electrons within the molecule can concentrate closer to one atom in the molecule. Produces a + / molecule. - (dipole) within the This temporary dipole induces a similar dipole within another molecule. These temporary dipoles result in the two molecules attracting each other. This attraction is the Dispersion Force! The attraction is dependant on: 1. The # of electrons in the molecules involved. 2. The ease of the electrons in the molecules to be dispersed within the individual molecules.

14 Larger atoms/molecules Produce Greater Temporary Dipoles. In general, As Molar Mass, The dispersion forces, The bp & mp of nonpolar molecules.

15 2) Dipole Forces These interactions occur in polar molecules. The + / - (dipole) of one polar molecule lines up with + / - (dipole) of another polar molecule (opposites attract). The greater the dipole moments (the measure of the polarity of a molecule) of the molecules, the stronger the attractive force. This interaction (attraction) really only works when the molecules are close together. When the molecules are in the gas phase, the dipole forces of attraction are negligible (as is the case for dispersion forces).

16 3) Hydrogen Bonding This attraction occurs in polar molecules, HOWEVER, only in molecules that have X H bonds where X = N, O, or F. This attractive force is an unusually strong dipole force of attraction. Why is the hydrogen bonding such a powerful attractive force? 2 Reasons: 1. There is a large difference in the electronegativity of the X and H H(2.2) F(4.0) H(2.2) O(3.5) H(2.2) N(3.0) The H atom almost behave as a naked proton. 2. The H atom is very small. The lone pair of electrons on F, O, and N can get really close to H.

17 Let s look at the pattern of boiling points. bp( o C) bp( o C) bp( o C) NH 3-33 H 2 O 100 HF 19 PH 3-88 H 2 S -60 HCl -85 AsH 3-63 H 2 Se -42 HBr -67 SbH 3-18 H 2 Te -2 HI -35 Note the effect of hydrogen bonding in the first row of boiling points. IMPORTANT Although these intermolecular force are very important, they are very weak compared to a covalent bond.

18 Section 9.4: Network Covalent, Ionic, and Metallic Solids Most Molecular substances are gases or liquids at room temperature. Most NON-molecular substances (network covalent, ionic, and metallic) are solids at room temperature. 1. Network Covalent Solids These solids are made of atoms joined by a continuous network of covalent bonds. In general, these solids are: a. High melting (over 1000 o C) - In order for this type of solid to melt, bonds need to be broken - Remember: When molecular solids to melt, only interactions need to be broken!

19 b. This type of solid is typically insoluble in all common solvents. - Why? Bonds need to be broken! NOT EASY TO DO!!! c. These solids are poor conductors of electricity. - Why? No mobile electrons are available

20 Example: Carbon 2 types of solids exist: Graphite & Diamond Both have very high melting points >3500 o C.

21 Ionic Solids These solids are held together by very strong electrostatic attractive forces (ionic bonds). 1. They are composed of cations/anions. 2. They are non-volatile (do not become gases very easily). 3. They are high melting ( o C). 4. They do not conduct electricity (they only do when they form aqueous solutions or they are molten). 5. Many (but not all) are soluble in water.

22 Metallic Solids A structural unit of electrons and metal cations. Positive metal ions anchored in position with electrons moving around from one metal ion to another. Metals are highly conductive - Very mobile electrons metals have very low electronegativities. Metals have a high thermal conductivity - Very mobile electrons vibrate Metals are ductile and mobile. Metals have high luster. - Electrons within a metal can absorb and emit light energy very easily. Metals are insoluble in common solvents.