Phase changes

• Phase diagrams show the states of matter (solid, liquid, gas) under different temperatures and pressures.
• Forces that hold matter together include: Ionic bonds, Covalent bonds, Metallic bonds, Hydrogen bonds, Dipole-dipole attractions, and London dispersion forces
• Chemical bonds are of primary importance, but for molecular substances, there are three new subcategories of weaker forces that hold molecules together
• In condensed phases (liquids and solids), atoms/ions/molecules are in direct contact with one another and interact
• In the case of single-celled organisms, substances can easily enter the cell due to a short distance, while in multicellular organisms, the distance is larger due to a higher surface area to volume ratio
• Phase changes occur as heat is added: at low temperatures, atoms and molecules stick together (solid), at high temperatures, molecules and atoms separate from one another (gas)
• Gases completely fill whatever space they occupy and exert pressure
• Network compounds require chemical bonds to be broken for melting or boiling, while molecular compounds only need intermolecular forces to be broken
• Hydrogen bonds are the strongest of the intermolecular forces, resulting from strong dipole-dipole attractions
• Hydrogen bonds occur in molecules like H-F, H-O, and H-N
• Dipole-dipole attractions are weaker than hydrogen bonds and occur between molecules with permanent dipoles
• Dipole-Dipole Attractions:

  • Occur in all polar compounds not exhibiting "hydrogen bonds"
  • Similar to hydrogen bonds but weaker
  • Involves the attraction between the positive pole of one polar molecule and the negative pole of another
• London Dispersion Forces:

  • Weakest among intermolecular forces
  • Occur in nonpolar compounds
  • Formed from temporary, transient dipoles induced by asymmetric electron distributions
• IMF Decision Summary:

  • Questions to ask when evaluating a compound:
  • Is it a network compound?
  • Does it have O-H, N-H, or F-H bonds?
  • Is the molecule polar?
  • Is the molecule nonpolar?
• Examples of IMF's for different compounds:

  • CH3NH2: Hydrogen bonds
  • CH3CH3: London dispersion forces
  • NaCH3COO: Ionic
  • CH3CH2F: Dipole-dipole attractions
• Vapor Pressure:

  • Liquid molecules/atoms jump off the surface into the gas phase
  • At equilibrium, rate of leaving = rate of returning
  • Gas pressure measured in the headspace at equilibrium is the liquid's vapor pressure
• Vapor Pressure Curves:

  • Vapor pressure increases with temperature
  • Boiling begins when vapor pressure equals air pressure
  • Normal boiling point is when vapor pressure is 760 torr
• Different substances have different vapor pressures at the same temperature
• Ethyl alcohol has a normal boiling point of 78.5 °C
• Dimethyl ether has a boiling point of -24 °C
• Methylene chloride in bubble lights boils at 39.75 °C
• Boiling Point Determiners:

  • Type and extent of interactions: bonds beat intermolecular forces
  • H-bonds > dipole-dipole > London forces
  • If interactions are the same, then higher molecular weight has a higher boiling point
  • If interactions and molecular weight are similar, geometries with more contact have a higher boiling point
• Nature and Extent of Interaction:

  • H2O is completely H-bonded with a boiling point of 100 °C
  • CH3CH2OH is H-bonded at one end with a boiling point of 78.4 °C
  • CH3OCH3 has dipole-dipole interactions with a boiling point of -24 °C
• Molecular Weight:

  • CH4 (methane) has London dispersion forces with a boiling point of -164 °C
  • CH3CH3 (ethane) has London dispersion forces with a boiling point of -89 °C
  • CH3CH2CH3 (propane) has London dispersion forces with a boiling point of -42 °C
• Molecular Shape:

  • n-pentane is linear with a boiling point of 36.2 °C
  • Dimethyl propane is "spherical" with a boiling point of 9.5 °C
• Predicting the Types of Intermolecular Forces:

  • Bonding forces are stronger than nonbonding (intermolecular) forces
  • Hydrogen bonding occurs only when there is N, O, or F covalently bonded to H
  • Dispersion forces are decisive when the difference is molar mass or molecular shape
• Sample Problems:

  • CH3NH2 has a higher boiling point than CH3F
  • CH3OH has a higher boiling point than CH3CH2OH
  • Hexane has a higher boiling point than 2,2-dimethylbutane
• Solids:

  • Change to solid called freezing, solidification, or fusion
  • As liquids are cooled and/or compressed, atoms/molecules get locked into place
  • "Locks" are bonds or intermolecular forces
  • Solids slowly locked in an orderly 3-D array are crystalline
  • Solids quickly cooled with no orderliness are amorphous
• Amorphous solids:

  • Molecules or atoms just stick together as temperature drops; no organized orientation
  • Crystalline solids have a sharp melting point; amorphous materials soften over a large temperature range
  • Examples: glass, plastics, tar, hard candy, opal, obsidian
• Network Solids vs. Molecular Solids:

  • Network solids are held together by bonds (ionic, covalent, metallic) and have high melting points because bonds must break (e.g., salt, quartz, iron)
  • Molecular solids have covalent molecules held together with intermolecular forces and have lower melting points because only intermolecular forces must break (e.g., ice, sugar, wax)
• Behavior of Solids:

  • Ionic network: high electrical conductivity as liquids
  • Covalent network: high electrical conductivity, except graphite
  • Metallic network: medium electrical conductivity as liquids and solids
  • Molecular: low electrical conductivity, none
• Allotropes:

  • Some elements come in more than one solid form called allotropes
  • Five of carbon's allotropes: graphite, diamond, buckyball, nanotubes, soot
• Carbon Allotropes:

  • Diamond: 3-D network
  • Graphite: sheets
  • C60: buckyball molecule
  • Glassy carbon: 3-D amorphous network
• Snowflakes:

  • Hydrogen bonds in H2O molecules become more important in its solid form
  • The crystal lattice is such that each molecule is strictly oriented into a 3-dimensional hexagonal array
• Snowflake Development:

  • When humidity is very low, end plates grow (plate stacking)
  • As humidity increases, edge plates grow
  • When humidity is high, points grow
  • Snowflakes travel through a range of conditions, leading to various forms of growth and artistic shapes
• Freezing and Boiling:

  • The final topic in Chapter deals with the energy changes that accompany phase changes
  • Temperature changes as heat is added to a sample of H2O that starts out as ice
• Phase Changes:

  • Solid to liquid: melting (fusion)
  • Liquid to solid: freezing
  • Liquid to gas: vaporization
  • Gas to liquid: condensation
  • Endothermic and exothermic processes
  • Deposition and sublimation
• Heating Curve for H2O:

  • Shows the temperature changes and heat added during phase changes
  • Specific heat of ice, heat of fusion, and heat of vaporization are key points on the curve
• In single-celled organisms, substances can easily enter the cell due to a short distance, while in multicellular organisms, the distance is larger because of a higher surface area to volume ratio
• Multicellular organisms require specialised exchange surfaces for efficient gas exchange of carbon dioxide and oxygen due to their higher surface area to volume ratio