Ch. 7 Liquids and Solids

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Properties

generally more dense than gases
stronger attractive forces between individual particles
liquids conform to the shape of a container and fill depending on its mass
solids keep its shape regardless of container

Intermolecular Forces

types of forces attracting molecules to each other

Dipole-Dipole Forces

most compounds unequally share electrons, creating a dipole
- one side partially positive ( $\delta+$ ), other side partially negative ( $\delta-$ )
$\textbf{Force}=\frac{(\delta+)(\delta-)}{r^2}$
need high force to overcome kinetic energy and condense to liquid
- done by increasing pressure (decrease distance) or decreasing temperature (decrease kinetic energy)
- boiling/condensation point indicates required kinetic energy
highly polar molecules have higher boiling points than less polar molecules

London Dispersion Forces

explanation for nonpolar gases
electrons in constant motion create instances of instantaneous dipoles even in nonpolar substances
can attract instantaneous dipole or induce dipole in a nearby atom/moleculex
noble gases have weak LDFs, so have low boiling points
polarizability: how easily an electron cloud can become a dipole
- halogens have larger electron clouds with loosely held electrons so are more polarizable $\rightarrow$ higher boiling points
- in general, more electrons $\rightarrow$ more polarizability
normal alkanes: $n$ $n$ -alkane has formula $\ce{C_nH_{2n+2}}$ $C X_{n} H X_{2 n + 2}$ , and higher $n$ $n$ have higher boiling points
- more hydrogen $\rightarrow$ more instantaneous dipoles $\rightarrow$ more forces and higher boiling point

Hydrogen Bonding

large electronegativity difference between $\ce{H}$ and $\ce{O}$ , $\ce{F}$ , $\ce{N}$ create strong dipole-dipole forces called hydrogen bonds
strongest of the three intermolecular forces

Summary

Force	Description
LDF	weak forces caused by brief unequal electron distribution
Dipole-Dipole	force between partial positive and partial negative ends of molecules
Hydrogen Bond	strong forces between $\ce{H}$ and $\ce{O}$ , $\ce{F}$ , or $\ce{N}$

Physcial Properties of Liquids

surface tension: increase in intermolecular forces at liquid surface
- when cohesive forces (attractions between identical molecules) are stronger than adhesive forces (attractions between different molecules) of a surface, the liquid beads; otherwise, uniformly spreads
- clean glass has low adhesive forces, so dishwater easily beads and evaporates, leaving spots $\rightarrow$ detergent contain surfactants that reduce cohesive forces so dishwater doesn't bead
viscosity: liquid's resistance to flow
- low attractive forces $\rightarrow$ molecules move past each other easily $\rightarrow$ liquid flows easily
- syrup contains sugar molecules, containing $\ce{-OH}$ groups that hydrogen bond to water
- high temperature $=$ weaker IMFs $\rightarrow$ lower viscosity
evaporation: liquid $\rightarrow$ $\to$ gas
- needs sufficient kinetic energy (proportional to Kelvin temperature) to overcome IMFs
- more surface area $=$ faster evaporation (more are close to surface)
- higher temperature $=$ faster evaporation (more have kinetic energy higher than escape energy)
- liquid evaporates at a uniform rate
- at boiling point, molecules do not have to be at surface (bubbles rise to surface)
vapor pressure: pressure of gas formed from liquid in a closed container
- in closed container, liquid evaporates, gas hits walls and liquid, condenses when hitting liquid, eventually evaporation rate $=$ condensation rate $\rightarrow$ dynamic equilibrium
- dependent only on nature of liquid (attractive forces) and temperature (kinetic energy)
natural boiling point: boiling point at $\pu{760 mm Hg}$ $760 m m H g$ ( $\pu{1.00 atm}$ $1.00 a t m$ )
- higher pressure $\rightarrow$ higher boiling point
- at high altitudes (low pressure environments), pressure cookers are used to increase pressure, increase boiling point, cooking food faster
heat of vaporization ( $\pu{J/g}$ $J / g$ ): energy needed to convert $\pu{1g}$ $1 g$ liquid into $\pu{1g}$ $1 g$ gas at normal boiling point
- symbol: $\Delta H_\text{vap}$
- alternatively, if $\pu{1 mol}$ is vaporized, it is called molar heat of vaporization with units $\pu{J/mol}$
- if similar-sized molecules, hydrogen-bonded have largest $\Delta H_\text{vap}$ , $\text{polar} > \text{nonpolar}$ , etc.

Solids

have rigid crystal structure

metallic crystals: rigid structure of metal nuclei and inner electrons
- very mobile valence electrons
  - electric charge and thermal energy conductivity
  - bonding metal atoms together $\rightarrow$ varying boiling points
  - energy needed often called lattice energy
- some soft, some hard/brittle, most malleable, caused by crystal structure
- substitutional alloys replace similarly-sized atoms, resulting in properties between those of the two metals
- interstitial alloys incorporate smaller atoms within existing structure with little volume change, resulting in stronger forces and stronger and harder alloy
ionic crystals: rigid crystalline structures (lattices) made of ionic compounds
- high lattice energy required to separate $\rightarrow$ high melting/boiling points
- extremely brittle, as disruptive force moves like-charged ions together, causing reuplsion
molecular crystals: crystals composed of nonmetals or covalent molecules
- held together by IMFs, weaker than ionic bonds $\rightarrow$ generally soft with low melting points
network covalent crystals: lattice structure where atoms are covalently bonded
- forms one large molecule all with covalent bonds
- can be very hard (diamond, $\ce{SiC}$ )
amorphous substances: noncrystalline materials
- no distinct, sharp melting point; instead gradually softening

Phase Changes

in constant pressure, effect of heat can be explained with a heating/cooling curve
as heat is added at a constant rate:
1. solid temperature increases until melting point
2. temperature stays constant until fully converted to liquid
3. liquid temperature increases until boiling point
4. temperature stays constant until fully converted into gas
5. gas temperature increases
opposite in a cooling curve

heat capacity ( $\pu{J /\degree C}$ ): reciprocal of slope
specific heat ( $\pu{J /g \degree C}$ ): heat capacity divided by mass of sample
molar heat/enthalpy of fusion (melting) ( $\pu{J / mol}$ ): first plateau, heat added divided by number of moles
molar heat/enthalpy of vaporization (boiling) ( $\pu{J / mol}$ ): second plateau, head added divided by number of moles
supercooling (in cooling curve): liquid is cooled below melting point and remaind liquid
- metastable; will rapidly crystallize if disturbed or if a seed crystal is added