Unit 2: Nomenclature, Intermolecular Forces and Properties of Solutions

Unit 2: Nomenclature, Intermolecular Forces and Properties of Solutions

Nomenclature Intermolecular Forces Phase Changes and Phase

Diagrams Saturated Solutions Solubility Concentration Units Colligative Properties

Nomenclature You are responsible for reviewing the

conventions for naming and writing formulas for ionic compounds, binary molecular compounds, and acids that were presented in Chem I.

This material will not be covered in lecture but will represent about 20% of the Unit 2 Exam.

You will be responsible for ionic compounds that contain any of the ions given in the Ion Chart handout.

Nomenclature Resources for reviewing nomenclature:

Chapter 2 of your text

Nomenclature slides on KMB’s website under Unit 2 Powerpoint lectures.

Ion chart handout

General Chemistry Tutorial for naming and writing formulas for ionic compounds on KMB’s website.

Intermolecular Forces

The fundamental difference between states of matter is the distance between particles.


The solid and liquid states are referred to as condensed phases because the particles are closer together.

Intermolecular forces The physical state of a substance at a

particular temperature and pressure depends on two antagonistic entities:

The kinetic energy of the particles

The strength of the intermolecular forces between the particles the attractive forces between particles in a

solid, liquid, or gas

Converting from one physical state to another requires the molecules to gain enough kinetic energy to overcome the intermolecular forces


Intermolecular forces are not nearly as strong as the intramolecular attractions (ionic or covalent bonds) that hold compounds together.

Intermolecular forces, however, have a significant impact on the physical properties of compounds: boiling point melting point vapor pressure solubility


There are four types of intermolecular forces: Dipole-dipole interactions (forces)

polar molecules London dispersion forces

all molecules Hydrogen bonding

molecules with H-F, O-H, or N-H bonds

Ion-dipole forces An ion and a polar molecule


Polar Molecules contain polar covalent bonds which

are asymmetrically distributed within the moleculecontain a “positive” end and a “negative”end

Examples:HClH2OCH3OH

-+

OH H

-

+ +

OHC

H

H H

-

++


To determine if a molecule is polar, identify all polar covalent bonds:

0 polar covalent bonds = nonpolar molecule

1 polar covalent bond = polar molecule

>2 polar covalent bonds = polar or nonpolar

Must use the electron domain geometry to determine polarity


If a molecule has two or more polar covalent bonds, it may be either a polar or nonpolar molecule: Draw the molecule in 3 dimensions

using its electron domain geometry

Draw the bond dipole moments for the polar covalent bonds

Nonpolar if they are symmetrical or offset each other

Polar if they are asymmetrically arranged


Example: Identify each of the following molecules as a polar or nonpolar molecule.


Example: Which of the following molecules are polar: CHCl3, CO2, Br2, HF, CH3OH, CH3OCH3?


Polar molecules have large dipole moments A measure of the separation

between the positive and negative charges in polar molecules.

+ -

H – F

+ -

Intermolecular Forces Dipole-dipole interactions (forces) are

found between polar molecules. attractive intermolecular forces

resulting from the attraction of the positive and negative ends of the dipole moments of polar molecules

The most stable arrangement of polar molecules is the one in which the positive end of one molecule is oriented toward the negative end of the other molecule.


For example, molecules of CH3Cl are held together by dipole-dipole interactions:


When a liquid vaporizes, the intermolecular forces must be overcome. As polarity increases, the strength of

the dipole-dipole interactions increase.Large Hvap

High BP


London dispersion forces are present in ALL molecules. Temporary dipole moments lasting

for fractions of a second are induced in one molecule by other nearby moleculeselectrons in molecules are displaced from symmetrical arrangement


London dispersion forces require close surface contact between two (or more) molecules. The strength of the London

dispersion forces is roughly proportional to surface area.As surface area increases, LDF increase and BP increases


For molecules with similar molecular formulas: Long, skinny (unbranched)

molecules have greater surface areahigher BP

Shorter, branched molecules are more spherical and have less surface arealower BP

C C

H

H

H

H

H

C

H

H

C

H

H

C

H

H

H

CH3C

CH3

CH3

CH3


The strength of dispersion forces tends to increase with increasing molecular weight. Larger atoms have larger electron

clouds, which are easier to polarize.

In general, as MW increases, BP increases.


Although BP tends to increase with increasing MW, there are some exceptions.

Why is the BP of water so much higher than expected?


Compounds containing H-F, O-H, and/or N-H bonds exhibit hydrogen bonding: a strong dipole-dipole interaction

between a hydrogen atom that is covalently bonded to either O, N, or F and a lone pair of electrons on a different O, N, or F atom


Due to the large difference in electronegativity, O-H, N-H, and F-H bonds are highly polar H has a partial positive charge in

such bonds

The H atom is strongly attracted to the nonbonding electrons on other N, O or F atoms.

O - H N - H- +-+

F - H +-


Example: Which of the following compounds can form hydrogen bonds with another molecule of the same substance?

C C

H

H

H

H

H

N

H

C

H

H

C

H

H

H

C C

H

H

H

H

H

C

H

C

H

H

C

H

F

H

H

C C

H

H

H

H

H

O C

H

H

C

H

H

H

C C

H

H

H

H

H

C

H

C

H

H

C

H

OH

H

H

C C

H

H

H

H

H

N

CH3

C

H

H

H

Intermolecular Forces Impact of hydrogen bonding

on BP: Hydrogen bonding leads to

higher BPH2O forms H-bondsH2S, etc cannot form H-bonds

As the number of hydrogens capable of forming hydrogen bonds increases, the boiling point increases.

Intermolecular Forces Impact of hydrogen bonding on BP:

Due to greater differences in electronegativity, OH forms stronger hydrogen bonds than NH

Compounds with OH have higher BP’s than similar compounds with NH

CH3CH2CH2CH2NH2

BP = 77-78oC

CH3CH2CH2CH2OH

BP = 117-118oC


Example: Consider each pair of compounds separately. Which compound in each pair has the higher BP? Explain why.

C C

H

H

H

H

H

O C

H

H

C

H

H

HC

H

H

C

H

C

H

H

C

H

OH

H

H

H vs.

CH3CCH3

O

vs. CH3CCH3

CH3

H

HOCH2CH2OH vs. CH3CH2OH

CH3CH2CH2CH2CH3 vs. CH4

Intermolecular Forces and Solubility

Intermolecular forces also determine the solubility of organic compounds. “Like dissolves like”

polar compounds dissolve in polar solvents

nonpolar compounds dissolve in nonpolar solvents

Ionic compounds generally dissolve readily in water because water hydrates or solvates the individual ions


Ionic compounds tend to dissolve in polar solvents like water because of ion-dipole forces attractive force between an ion and

the partially “charged” end of a polar molecule.


An aqueous solution of an ionic compound such as NaCl contains solvated cations and anions:

- +

Solvation of anion

Solvation of cation

Examples of ion-dipole forces


Polar compounds dissolve in polar solvents due to: dipole-dipole interactions H-bonding

Hydrogen bonding between methyl amine and water. H

H

N

OH

HCH3


Nonpolar compounds do not dissolve appreciably in water because they cannot break the hydrogen-bonding network that exists.


Example: Will each of the following vitamins be water soluble or fat soluble?

O

OHO

HO

CH2OHHO

O

OH

CH3H3C

H3C

H3C

CH3 CH3 CH3

Unit 2: Nomenclature, Intermolecular Forces and Properties of Solutions

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