Thursday, Nov. 29 th : “A” Day Friday, Nov. 30 th : “B” Day Agenda  Section 11.1 Quiz  Start Section 11.2: “Intermolecular Forces”  Comparing Ionic/Covalent Bonds, Dipole-Dipole Forces, Hydrogen Bonds  Homework:  Section 11.2 review, pg. 392: #1-5  Writing Assignment: “What Would Life Be Like?”  Concept Review: “Intermolecular Forces”

Section 11.1 Quiz “States and State Changes”

 You can use your notes and your book to complete the quiz on your own… Use the Force!

Intramolecular Vs. Intermolecular Forces

 INTRAmolecular forces are the forces within a molecule.  Examples: ionic and covalent bonds  INTERmolecular forces are the forces of attraction between molecules.

 Examples: Dipole-dipole forces and London dispersion forces

Sec. 11.2

“Intermolecular Forces”  It takes energy to overcome the forces holding particles together.

 It takes energy to cause a substance to go from the liquid to the gaseous state.

 The melting point and boiling point of a substance are good measures of the strength of the forces that hold the particles together.

Comparing Ionic and Covalent Compounds  Ionic: metal bonded to non-metal (L side/R side)  Covalent: two non-metals bonded together (R side)

Ionic Vs. Covalent Compounds  Most covalent compounds melt at lower temperatures than ionic compounds because the forces holding them together are weaker.

 Ionic substances with small ions tend to be solids that have high melting points. (NaCl)  Covalent substances tend to be gases and liquids or solids that have low melting points. (Cl 2 )

The strength of the ionic forces depends on the size of the atoms and the amount of charge.

 Large ions can’t get close together, so the distance between them is greater and the forces are weaker.

(Homework hint: polyatomic ions are large)

 Example: KCl melts at a lower temperature (770˚C) than NaCl (801˚C) because K + is larger than Na + .

The strength of the ionic forces depends on the size of the atoms and the amount of charge.

 Ions with larger charges have a greater ionic force between them than ions with smaller charges.

 Example: MgF 2 melts at a higher temperature (1,263˚C) than NaF (993˚C) because Ca 2+ has a larger charge than Na + .

Intermolecular Forces Attract Molecules to Each Other 

Intermolecular force:

the forces of attraction between molecules of covalent compounds.

 Intermolecular forces include dipole-dipole forces and London dispersion forces.

 Intermolecular forces are:  Short-range  Decrease rapidly as molecules get farther apart  Don’t have much impact on gases because gas molecules are too far apart

Dipole-Dipole Forces  In dipole-dipole forces, the positive end of one molecule attracts the negative end of a neighboring molecule.

 Bonds are polar because of electronegativity differences between bonded atoms. 

Dipole-dipole forces:

interactions between polar molecules.

Dipole-Dipole Forces Affect Melting and Boiling Points  When molecules are very polar, the dipole dipole forces are very significant.

 The more polar the molecules are, the higher the boiling points of the substances.

How Do You Know if a Substance is Polar?

 Is it an ionic substance or a covalent compound?

 Ionic substances always have positive and negative ions.  In covalent compounds, check the differences in electronegativity between the atoms. (pgs. 194-195) The greater the electronegativity difference, the more polar.

 Symmetrical compounds tend to be non-polar because the atoms share electrons equally.

(Example: CCl 4 )

Hydrogen Bonds 

Hydrogen bond:

the dipole-dipole force occurring when a hydrogen atom that is bonded to a highly electronegative atom of one molecule is attracted to two unshared electrons of another molecule.

 In general, compounds with hydrogen bonding have higher boiling points than comparable compounds.

 Example: H 2 S and H 2 O  H 2 S: Boiling point = -60.7˚C (gas at room temp)  H 2 O: Boiling point = 100.0˚C (liquid at room temp)

Hydrogen Bonds Form with Electronegative Atoms  Strong hydrogen bonds can form with a hydrogen atom that is covalently bonded to very electronegative atoms in the upper-right part of the periodic table: nitrogen, oxygen, and fluorine.

Why Are Hydrogen Bonds Such Strong Dipole-Dipole Forces?

1. The large electronegativity difference (high polarity gives molecules more ionic character) 2. Hydrogen’s small size  When hydrogen’s one electron is pulled away by a highly electronegative atom, there are no more electrons under it and the single proton of hydrogen’s nucleus is partially exposed. This proton is strongly attracted to the unbonded pair of electrons of other molecules.

Hydrogen Bonding

 Water is a liquid at room temperature because of hydrogen bonding.

Hydrogen Bonding Explains Water’s Unique Properties  Each water molecule forms multiple hydrogen bonds, so the intermolecular forces in water are strong.

 When water forms ice, the ice crystals have large amounts of open space. Thus, ice has a low density.

 Water is unusual in that its liquid form is denser than its solid form.

 The energy of hydrogen bonds is lower than normal chemical bonds, but can be stronger than other intermolecular forces.

Ice and Water

Homework

 Section 11.2 review, pg. 392: #1-5  Writing Assignment: “What Would Life Be Like?”  Concept Review: “Intermolecular Forces”