1. Chapter 13: Liquids and Solids Chemistry 1020: Interpretive chemistry Andy Aspaas, Instructor

2. Phases of water Water is a liquid between 0 °C and 100 °C at normal atmospheric pressure –As liquid water is heated, its temperature rises It begins to boil at 100 °C, and its temperature will not raise any further until all the water has converted to steam When the steam is heated, its temperature will rise beyond 100 °C –As liquid water is cooled, its temperature falls until it begins to freeze at 0 °C, and will not fall further until water is completely frozen

3. Liquid and solid water If held at 0 °C, a mixture of liquid water and ice will coexist indefinitely Unlike most other substances, water expands as it freezes –Liquid water at 0 °C: d = 1.00 g/mL –Ice at 0 °C: d = 0.917 g/mL Relatively high specific heat of water: 4.184 J/g·°C –A relatively large amount of energy is required to change water’s temperature

4. State changes State change: change between solid, liquid, or gas –A physical change: no chemical (ionic or covalent) bonds are broken in the process Intermolecular forces: forces that attract water molecules to each other –Occur when a molecule has a dipole moment (a partial positive side and a partial negative side) Intermolecular forces must be broken when ice melts or water boils, so energy is required for both these processes

5. State changes Molar heat of fusion: energy required to melt 1 mole of a solid substance –In a solid, molecules are locked together and can only vibrate –When energy is added, the vibrations increase until molecules break apart and move freely to form a liquid (still many intermolecular forces though) Molar heat of vaporization: energy required to change 1 mol of a liquid to its vapor –Energy added to a liquid will break nearly all intermolecular forces so the molecules spread out and form a gas

6. Intermolecular forces Dipole-dipole interaction: when polar molecules attract each other Hydrogen bonding: attraction between an electropositive hydrogen and an electronegative element of another molecule –A very strong intermolecular force, accounts for the relatively high boiling point of water London dispersion forces: instantaneous dipoles caused by random dispersion of electrons –The only intermolecular force in nonpolar molecules like N 2 (why liquid N 2 can only exist at very low temperatures)

7. Types of solids Crystalline solids: regular arrangement of particles –Ionic solids: like NaCl –Molecular solids: like sugar (sucrose) or ice –Atomic solids: contain only one element (all metals, diamond, silicon)

8. Ionic solids Held together by very strong ionic bonds (full positive and full negative charges attracted to each other) –Very high melting points (NaCl is over 800 °C) –Ions are packed as efficiently as possible - small ions fit in the holes left by packing large ions

9. Molecular solids Molecule is the fundamental particle of a molecular solid –Ice (H 2 O), dry ice (CO 2 ), sucrose (C 12 H 22 O 6 ) Compared to other forms of solids, molecular solids usually have low melting points –Dipole-dipole interactions and London dispersion forces are nowhere near as strong as ionic or covalent bonds

10. Atomic solids Noble gases (group 8) are only solid at very low temperatures –Full valence shell = only London dispersion forces Diamond, crystalline solid carbon: one of the strongest solids known –all covalent bonds –1 diamond = 1 large molecule!

11. Bonding in metals Metals change shape easily but usually have very high melting points –Metal atoms are arranged in a regular crystal-like arrangement –But the valence electrons flow together around the atoms to form a “sea” of electrons –Why metals can conduct electricity