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Chapter 14:

Chapter 14:. Liquids and Solids. Students will learn about…. Water phase Solid, liquid, gas Energy to change phase q calculations Intermolecular forces Dipole-dipole, hydrogen bonding, London dispersion forces Vapor pressure. Phase. Synonymous to state

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Chapter 14:

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  1. Chapter 14: Liquids and Solids

  2. Students will learn about….. • Water phase • Solid, liquid, gas • Energy to change phase • q calculations • Intermolecular forces • Dipole-dipole, hydrogen bonding, London dispersion forces • Vapor pressure

  3. Phase • Synonymous to state • 4 states of matter: solid, liquid, gas, and plasma • States of matter determined by the degree of motions of particles • Gas: great deal of motion = weak attraction between particles • Liquid: considerable amount of motion = somewhat strong attraction between particles • Solid: vibration of particles = strong attraction between particles

  4. Properties of Gas vs. Solid • Gases • Low density • Highly compressible • Fill container • Solids • High density • Slightly compressible • Rigid (keeps its shape)

  5. Changing Phase (Heating/Cooling Curve) Notice that T is NOT changing Notice that T is NOT changing Going from left to right = energy required = endothermic Going from right to left = energy released = exothermic

  6. Melting point? Freezing point? Boiling point? Condensation point?

  7. Normal • “Normal” boiling point and “Normal” freezing point • Normal = 1 atm pressure • The boiling point varies greatly with the atmospheric pressure Ex) The BP of water at 1 atm = 100 °C The BP of water at 0.92 atm = 98.7 °C

  8. Energy Requirements for Changing States • When states changing, use: • q = m·s·ΔT • q = amount of heat energy • m = mass • s = specific heat (capacity) • ΔT = Tf‒ Ti • When state not changing, use: • solid to liquid or vice-versa • q = ΔHf·m • ΔHf = heat of fusion per gram or one mole (the amount of heat needed to melt/freeze 1 g (or 1 mol) of solid/liquid) • liquid to gas or vice-versa • q = ΔHv·m • ΔHv = heat of vaporization per gram or one mole (the amount of heat needed to boil/condense 1 g (or 1 mol) of liquid/gas)

  9. Examples 1. (Pg 451) Calculate the energy required to melt 8.5 g of ice at 0 °C. The molar heat of fusion for ice is 6.02 kJ/mol.

  10. (Pg 452) Calculate the energy (in kJ) required to heat 25 g of liquid water from 25 °C to 100. °C and change it to steam at 100. °C. The specific heat capacity of liquid water is 4.18 J/g ·°C, and the molar heat of vaporization of water is 40.6 kJ/mol.

  11. Forces between molecules *Bonds are intramolecular forces Intermolecular Forces

  12. Types of Intermolecular Forces 1) Dipole-Dipole attraction • Among polar molecules • 1% strength of covalent or ionic bond

  13. 2) Hydrogen Bonding • An extreme case of dipole-dipole due to a greater difference in electronegativity and relatively small size of atom • Only happens when a molecule has H ‒ N, H ‒ O or H ‒ F • Unusually high BP of H2O – two hydrogen bonds in each molecule

  14. Example Draw two Lewis structures for the formula C2H6O and compare the boiling points of the two molecules.

  15. 3) London dispersion force = instantaneous dipole • Between all molecules, including non-polar molecules • More electrons  greater London dispersion force  higher BP and FP

  16. Example Consider the following compounds: NH3 CH4 H2 How many of the compounds above exhibit London dispersion forces?

  17. Vapor Pressure • Amount of liquid first decreases then becomes constant. • Condensation – gas converting to a liquid. • Rate of evaporation = Rate of equilibrium (called equilibrium) • The gas pressure at the equilibrium is the vapor pressure • More volatile liquid (=lower intermolecular forces) = higher vapor pressure • Higher the temperature = higher vapor pressure

  18. Examples 1. Which will show the larger vapor pressure? H2O (l) vs. CH3OH (l) CH3OH (l) vs. CH3CH2CH2CH2OH(l)

  19. 2. Which of the following would be expected to have the highest vapor pressure at room temperature? a) CH3CH2CH2OH b) CH3CH2CH2NH2 c) CH3CH2CH2CH3 d) CH3CH2CH3

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