Chapter 13: Liquids and Solids

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 N2 (why liquid N2 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 (H2O), dry ice (CO2), sucrose (C12H22O6) • 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