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Chapter 16 liquids and solids

Chapter 16 liquids and solids - Chapter 16 Liquids and...

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Chapter 16: Liquids and Solids Gases have low densities, high compressibilities, and completely fill a container. Solids have much greater densities, are compressivle only to a very slight extent, and are rigis—a solid maintains its shape irrespective of its container. These properties indicate that the components of a soid are close together and exert large attractive forces on each other. Enthalpy change for the melting of ice at 0 o C (the heat of fusion) with that for vaporizing liquid water at 100 o C (the heat of vaporization): H 2 O(s) → H 2 O(l) ∆H o fus = 6.02 kJ/mol H 2 O(l) → H 2 O(g) ∆H o vap = 40.7 kJ/mol Above suggests that there are extensive attractive forces among the molecules in liquid water, similar to but not as strong as those in the solid state. Densities of the three states of water: State Density (g/cm 3 ) Solid (0 o C, 1 atm) 0.9168 Liquid (25 o C, 1 atm) 0.9971 Gas (400 o C, 1 atm) 3.26 X 10 -4 The liquid and solid states show many similarities and are strikingly different from the gaseous state. 16.1 Intermolecular Forces Intramolecular bonding : (within the molecule) atoms can form stable units called molecules by sharing electrons. Condensed states of matter: liquids and solids Intermolecular forces: weak interactions that occur between, rather than within molecules When a substance like water changes from solid to liquid to gas, the molecules remain intact. Changes of state are caused by changes in the forces among the molecules rather than those within the molecules. If energy is added, the motions of the molecules increase, and they eventually achieve the greater movement and disorder characteristic of liquid water; the ice has melted. As more energy is added, the gaseous state is eventually reached, where the individual molecules are far apart and thus interacting relatively little. Molecules with dipole moments have a center with positive charge and a center with negative charge. They can attract each other electrostatically by lining up so that the positive and negative ends are close to each other: called dipole-dipole attraction . Dipole forces are about 1% as strong as covalent or ionic bonds, and rapidly become weaker as the distance between the dipoles increases. Strong dipole-dipole forces involve hydrogen bound to a highly electronegative atom, such as nitrogen, oxygen, or fluorine. Two reasons: the great polarity of the bond and the close approach of the
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dipoles, allowed by the very small size of the hydrogen atom. They are called hydrogen bonding. Boiling points of covalent hydrides in group 4A show steady increase with molar mass (going down a group). In 5A, 6A, and 7A, the lightest member has an unexpectedly high boiling point. Reason is because of large hydrogen-bonding interactions that exist among smallest molecules with most polar X—H bonds.
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