Ch 9 Liquids & Solids (IMF) Masterson & Hurley
Intra- and Intermolecular AP Questions: 2005 Q. 7, 2005 (Form B) Q. 8, 2006 Q. 6, 2007 Q. 2 (d) and (c), Periodic Trends AP Questions: 2001 Q. 8, 2002 Q. 6, 2002 Q. 6, 2003 (Form B) Q. 7, 2006 (Form B) Q. 7 and 8, 2007 (Form B) Q. 6
Bond energy is the energy required to break the bond, and can be utilized to learn the strength of a bonding interaction by measuring the bond energy.
Ch 9.1 Liquid-Vapor Equilibrium
9.2 Phase Diagrams
8.2 Electronegativity
Electronegativity: the ability of an atom in a molecule to attract shared electrons to itself Trend: the smaller the atom the greater its electronegativity. Electronegativity increases across a period with increasing atomic number and decreases going down a group. Moving left to right increases number of protons, causing a greater nuclear-electron pull.
The 4 most Electronegative: F, O, N, Cl
Pure covalent (nonpolar covalent) Two identical atoms share electrons equally Ionic and covalent bonds are the extreme bond types, between these extremes are intermediate cases Polar Covalent bond electrons are shared unequally and atoms end up with fractional charges: + and - So how do we determine polarity?
Ionic compounds generally have electronegativity differences greater than 1.6
Example: Place the following atoms in order of increasing electronegativity: Cs, Rb, S, Al, Sr, O
Example: Order the following bonds according to polarity (use 2 sig. Figs.): H-H, O-H, Cl-H, S-H, and F-H
8.3 Bond Polarity and Dipole Moments
A Dipolar molecule (polar covalent) is one that has a slight (+) and slight (-) charge This is called a dipole moment Dipole moment is indicted with an arrow, the arrow points to the most electronegative element
Molecules with polar bonds but not having a dipole moment This occurs when the individual bond polarities are arranged in such a way that they cancel each other out.
Molecules With Polar Bonds but no Dipole Moment: Linear, trigonal planar, tetrahedral symmetry of charge distribution (Note page 356 Table 8.2) CO 2 linear SO 3 trigonal planar CCl 4 tetrahedral
9.3 Intermolecular Forces
This chapter will focus on the condensed states of matter: liquids and solids and the forces that cause them to form a liquid or a solid
Intermolecular forces are weaker forces than covalent or ionic. These forces occur between molecules, NOT WITHIN molecules!
Intermolecular Forces (attractions occurring between molecules), may involve covalent or ionic bonding or they may involve weaker interactions Intramolecular Forces Intermolecular Forces The properties of the condensed states of matter (liquids and solids) involve these forces. For example, when water changes from solid to liquid to gas, the molecules remain intact. The changes are due to forces among the molecules rather than in those within the molecules
Remember temperature is a measure of random motions of the particles in a substance. When water changes from a liquid to a gas, the molecules remain intact the change in state is due to forces among the molecules rather than within them. If energy is added the motions of the molecules increase (greater movement and disorder) The energy needed to vaporize 1 mole of liquid water (40.7 kj) The energy needed to break the O-H bond in 1 mole of water is (934 kj)
London dispersion forces are forces that exist among the noble gases and nonpolar molecules A temporary dipole is formed due to an induced dipole in a neighboring atom--it is weak and short lived Important for large atoms--noble gases having low freezing points (the larger the atom, the greater the number of electrons Important for nonpolar molecules in attraction for each other, such as: H 2, CH 4, CCl 4, and CO 2
London dispersion forces exist between all molecules and account for the b.p. of the noble gases if there were no attractions between molecules, then they would never liquefy. It is believed that the electron cloud can shift temporarily to one side of the molecule becoming more negative than the other, causing a temporary dipole, inducing a similar dipole on a neighboring atom
The greater the charge on the nucleus (the atomic number), the larger the number of e- in the molecule, the greater the induced dipole Cl 2 (g), Br 2 (l) I 2 (s) for example, the halogens all experience London dispersion forces, but the force becomes stronger toward the bottom of the group: F 2 (g) not shown, Cl 2 (g), Br 2 (l), & I 2 (s)
Dipole Forces
Dipole-dipole attraction is a force that acts between polar molecules. In a condensed, liquid state, where many molecules are in close proximity, the dipoles find a compromise between attraction and repulsion Dipole-dipole forces are only 1% as strong as covalent or ionic bonds
Hydrogen Bonding
Molecules containing hydrogen have strong dipole-dipole forces called hydrogen bonding Hydrogen bonding has a very important effect on physical properties
Dipole-Dipole Attractions polar molecules attract each other, lining up so that their positive and negative poles are close to each other polar molecules generally have higher boiling points than nonpolar molecules of similar molar mass because they have dipole-dipole attractions in addition to London forces. An example is H 2 O
(a) Oxygen atoms are arranged in layers of distorted hexagonal rings. Hydrogen atoms lie between pairs of O atoms, closer to one (covalent bond) than to the other (hydrogen bond.) (b) At the macroscopic level, this structural pattern is revealed in the hexagonal shape of snowflakes. Hydrogen bonding is an unusually strong dipoledipole force among molecules in which H is bound to a highly electronegative atom such as: N, O, or F. This is the strongest type of intermolecular force (NH 3 and HF also exhibit this )
Nonpolar (Group 4A) show an increase in b.p. with molar mass For other groups, the lightest has higher b.p. Larger hydrogen bonding occurs between smaller molecules with the most polar X-H bonds. This is due to: large electronegativity difference between X-H bonds and the small size of the first element which allows close interactions strengthening the intermolecular forces
AP likes this diagram Smaller molecules (smaller amu) and larger electronegativity values lead to a more polar molecule These types of molecules will remain together in the liquid state even at high temperatures--hence very high boiling points
9.4 Network Covalent, Ionic, & Metallic Solids
Network Covalent Solids
Ionic Solids
An ionic compound results when a metal reacts with a nonmetal and forms an ionic bond The metal loses e- relatively easily, (low ionization energy), to the nonmetal which has a high affinity for e-. Electrostatic attraction of oppositely charged ions forms an ionic bond. Any compound that conducts an electric current when melted is an ionic compound Ions paired have lower energy (greater stability) than separated ions
Coulomb s Law The energy interaction between a pair of ions can be calculated by using Coulomb s Law: E = (2.31 x 10-19 J nm)(q 1 Q 2 ) E = energy in Joules ( r ) r = the distance b/w the ion centers in nm Q 1 and Q 2 = numerical ion charges
For example: solid NaCl has a distance between the centers of the Na+ and Clions is 2.76 Å (0.276 nm), what is the ionic energy per pair of ions? E = (2.31 x 10-19 J nm)(q 1 Q 2 ) ( r ) = (2.31 x 10-19 J nm)(+1)(-1) = -8.37 x 10-19 J 0.276 nm *The (-) sign indicates an attractive force, the ion pair has lower energy than the separated ions. * Coulomb s Law can be used to calculate the repulsive energy when two likecharged ions are brought together, the value will have a (+) sign.
Metals
9.5 Crystal Sturctures
Old Material
8.6 Partial Ionic Character of Covalent Bond
Calculating Percent Ionic Character Percent ionic character = (measured dipole moment of X-Y) x 100% (calculated dipole moment of X + Y - ) Ionic cmpds. generally have greater than 50% ionic character Ionic cmpds. generally have electronegativity differences greater than 1.6 Percent ionic character is difficult to calculate for compounds containing polyatomic ions
Forces At Work Attractive Forces: Repulsive Forces: proton-electron electron-electron proton-proton Energy is given off (bond energy) when two atoms achieve greater stability together than apart Let s watch the formation of Na + Cl 2 --> NaCl
Bonding Between Identical Atoms A bond will form if the energy of the aggregate is lower than that of the separated atoms. The distance where the energy is minimal is called the bond length. The system will act to minimize the sum of the positive (repulsive) energy terms and the negative (attractive) energy term. Let s watch the movie
8.4 Ions: Electron Configurations and Sizes
When two nonmetals react (covalent) they share e- to complete the valence e- configurations of both atoms. Both nonmetals attain noble gas e- configurations.
When nonmetal and rep.-group meal react (ionic) the nonmetal achieves the e- configuration of the next noble gas and valence orbitals from the metal are emptied to achieve the e- configuration of the previous noble gas.
Predicting Formulas of Ionic Compounds Placement of elements on the periodic table suggests how many e- are lost or gained to achieve a noble-gas configuration Group I loses 1 e-, Group II loses 2 e-, Group VI gains 2 e-, Group VII gains 1 e-, etc Formulas for compounds are balanced so that the total positive ionic charge is equal to the total negative ionic charge Al 2 +3 O 3-2 Total positive = +6 Total negative = -6
Sizes of Ions Anions larger than the parent atom Cations smaller than the parent atom Ion sizes increases within a family Isoelectronic ions: are ions containing the same number of e- For example: O 2-,F-, Na+, Mg 2+, Al 3+ each has 10 e- but the number of protons increases from 8 to 13 as we go from O 2- to Al 3+ causing a greater attraction as the positive charge on the nucleus increases, the ions become smaller
Sizes of Ions
Motor Oil (SAE 40 and SAE 10)
10.8 Vapor Pressure and Changes of State
Molecules of a liquid can escape the liquid s surface and form a gas in a process called vaporization or evaporation. Vaporization is endothermic because energy is required to overcome intermolecular forces in the liquid.
The energy required to vaporize 1 mole of a liquid at a pressure of 1 atm is called the heat of vaporization or the enthalpy of vaporization H vap Volatile liquids have a higher vapor pressure The stronger the IMF s the lower the vapor pressure The weaker the IMF s the higher the vapor pressure Vapor pressure increases with increasing temperature
Page 487 Exercise 10.6 Vapor pressure of water: 298 K and 23.8 torr and H vap = 43,900 J/mol Calculate vapor pressure of water at: 323 K ln (P vap, T1 ) = H vap (1-1) (P vap, T2 ) R T 2 T 1 ln (23.8 torr) = 43,900 J/mol (1-1) P 8.3145 J/K mol (323K 298K) P = 93.7 Torr As temp increases press increases
10.9 Phase Diagrams
Phase Diagram of Water
At 1 atm = Temp below 0 C Ice changes to liquid as energy is added (normal melting point) (no vapor pressure) As heating continues to 100 C boiling occurs, vapor pressure is 1 atm
Below 0.006 atm = sublimation occurs (solid to gas, no liquid conditions)
At 0.006 atm and 0.01 C = triple point, all three states of water are present
At 217.7 atm and 374.4 C = critical point, liquid changes to vapor, beyond the critical point an intermediate fluid occurs (not liquid, not gas)
Note at 1 atm what happens to dry ice (CO 2 ) Sublimation occurs
Read 10.3-10.7 On Your Own