Chemistry Chapter 10 & 11

Writing Lewis Structures

Writing Lewis Structures

• Draw skeletal structures of compound showing what atoms are bonded to each other. Put least electronegative element in the center

• Count total number of valence e- . Add 1 for each negative charge. Subtract 1 for each positive charge

• Complete an octet for all atoms except hydrogen

• If short electrons, add bonds

• If there are extra electrons, add them to the central atom

Electronegativity

• The ability of an atom to attract toward itself the electrons in a chemical bond

• F is highest

Electrons are shared unequally in what type of bond?

Polar covalent

Type of bond is determined by what?

Difference in electronegativity

Electronegativity < 0.4

Pure (nonpolar) covalent

Electronegativity 0.4-2.0

Polar covalent

Electronegativity > 2.0

Ionic

Polar molecules must

• Have polar bonds

• Have a shape (geometry) that allows it to be polar

• Dipoles that are equal and opposite cancel each other out

H2O

• Dipole moment

• Polar molecule

SO2

• Dipole moment

• Polar molecule

CO2

• No dipole moment

• Nonpolar molecule

CH4

• No dipole moment

• Nonpolar molecule

Non-polar

As long as dipoles are equal (same atoms) and opposite

Linear

Non-polar

Trigonal Planar

Non-polar

Tetrahedral

Non-polar

Trigonal Bipyramidal

Non-polar

Octahedral

Non-polar

Polar

Shapes that are not symmetrical are polar

Bent

Polar

Trigonal Pyramidal

Polar

Distorted Tetrahedral (see-saw)

Polar

Intermolecular Forces

Attractive forces between molecules

Intramolecular Forces

Hold atoms together in a molecule

Intermolecular vs Intramolecular

• 41 kJ to vaporize 1 mole of water (inter)

• 930 kJ to break all O-H bonds in 1 mole of water (intra)

“Measure” of intermolecular force

• Boiling point

• Melting point

• ΔHvap

• ΔHfus

• ΔHsub

Kinetic Energy vs Interactive Forces

• KE lowest, IF highest

• KE intermediate, IF intermediate

• KE highest, IF lowest

The boiling point increases as the __ increases

Atomic number

The stronger the particles’s attractions for each other, the _

Greater the amount of energy needed to separate them

The greater the energy required to separate these particles, the

Higher the boiling points

Dispersion (London) Forces (LDF)

Intermolecular force between nonpolar molecules or atoms caused by the presence of temporary dipoles in the molecules

Temporary (induced) Dipole

Separation of charge produced in an atom or molecule by a molecule by a momentary uneven distribution of electrons

Polarizability

Relative ease with which the electron cloud in a molecule, ion, or atom can be distorted, inducing a temporary dipole

All molecules and atoms have…

Dispersion forces

Dispersion

Momentary shift in e- density

Factors affecting strength of dispersion

• Size of atoms/molecules

• Shape of molecules

Dispersion - Size of atoms/molecules

• Larger atoms/molecules are more polarizable than smaller atoms/molecules

• Dispersion increases with polarizability

Dispersion - Shape of molecules

• Increased surface area = increased interactions between molecules

• Linear molecules have higher dispersion than branched molecules of similar molecular weight

Effect of Shape on Dispersion

The more spread out a molecule, the stronger the opportunity for dispersion forces and the higher the boiling point

Dipole-Dipole Interactions

Attractive force between polar molecules

Hydrogen Bond

• Strongest dipole-dipole interaction

• Occurs between H atom bonded to a small, highly electronegative element (F, O, N) and an atom of O or N or F in another molecule

Ion-Dipole Interaction

Attractive force between an ion and a molecule that has a permanent dipole

Sphere of Hydration

• Cluster of water molecules surrounding an ion in aqueous solution

• Sphere of solvation if solvent is other than H2O

Dipole-Induced Dipole

• Proximity of polar molecule causes dipole-induced dipole

Relative Strengths of Intermolecular Forces

• Weakest - Dispersion Forces

• Dipole Induced Dipole Forces

• Dipole-Dipole Forces

• Hydrogen Bonding

• Strongest - Ion Dipole Forces

Solubility

Depends on relative strength of solute-solvent interactions compared to solute-solute or solvent-solvent

Like dissolves like

• Ionic/polar solutes will be soluble in polar solvents

• Nonpolar solutes will be soluble in nonpolar solvents

Combination of Forces

More than one intermolecular force may need to be considered when examining larger molecules dissolved in a liquid solvent

Solubility decreases are relative energy of…

Hydrogen bonding decreases and dispersion increases

Hydrophobic (water-fearing)

Interaction that repels water, diminished water solubility

Hydrophilic (water-loving)

Interaction that attracts water, promotes water solubility

Equilibrium Vapor Pressure

• The vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation

• H2O (L) → H2O (g)

Dynamic Equilibrium

Rate of Condensation = Rate of Evaporation

Molar Heat of Vaporization (ΔHvap)

The energy required to vaporize 1 mole of liquid at its boiling point

Clausius-Clapeyron Equation

• ln P = [ (-ΔHvap) / (RT) ] + C

• P = equilibrium vapor pressure

• T = temperature (K)

• R = gas constant (8.314 J/K x mol)

Clausius-Clapeyron Equation at two temperatures

ln (p1 / p2) = (ΔHvap / R) x [ (1 /T2) - (1 / T1) ]

Boiling Point

The temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure

Normal Boiling Point

The temperature at which a liquid boils when the external pressure is 1 atm

Factors Affecting Physical States

• Intermolecular Forces: strength of attractive forces compared to kinetic energy of atoms/molecules

• Temperature: affects kinetic energy of atoms/molecules

• Pressure: affects distance between atoms/molecules

Solid-Liquid Equilibrium

• H2O (s) → H2O (L)

• The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium

Molar Heat of Fusion (ΔHfus)

The energy required to melt 1 mole of a solid substance at its freezing point

Solid-Gas Equilibrium

H2O (s) → H2O (g)

Molar Heat of Sublimation (ΔHsub)

The energy required to sublime 1 mole of a solid

Hess’s Law

ΔHsub = ΔHfus + ΔHvap

Cohesion

The attraction of like molecules due to intermolecular forces

Adhesion

The attraction of different molecules due to intermolecular forces

Cohesive Forces are responsible for…

Some properties of liquids such as surface tension and viscosity

Phase Diagram

A graphical representation of how the stabilities of the physical states of a substance depend on temperature and pressure

Equilibrium Lines

Represent phase changes between solid/liquid, liquid/gas, and solid/gas under specific conditions (T,P)

Supercritical Region

4th state of matter with properties intermediate between gas and liquid

Critical Temperature (Tc)

The temperature above which the gas cannot be made to liquify, no matter how great the applied pressure

Critical Pressure (Pc)

The minimum pressure that must be applied to bring about liquifacation at the critical temperature

Phase Diagram - Triple Point

Temperature/pressure where all 3 phases of a substance coexist

Phase Diagram - Critical Point

Specific temperature/pressure at which the liquid and gas phases have the same density

Phase Diagram - Supercritical Fluid

A substance above its critical temperature and pressure

The equation to calculate amount of heat (q) entering/leaving a substance

• q = c x m x ΔT

• q = c x m x (Tf - Ti)

T final > T initial

If a substance gains thermal energy, then the value of q is positive

T final < T initial

If a substance loses thermal energy, then the value of q is negative

Molar Heat of Fusion (ΔHfus)

The energy required to melt 1 mole of a solid substance at its freezing point

Enthalpy in Change of State equations

• q = nΔHfus

• q = nΔHvap

Solutions are __ made up of two or more substances

Homogenous mixtures

Solute

Substance being dissolved

Solvent

• Substance that is doing the dissolving

• Typically, present in largest amount

Solid Solution

Metal alloys (two or more metals)

Gas Solution

Example: air - made up of multiple gases

Three types of intermolecular forces influence whether substances will form a solution

• Solute-Solute

• Solvent-Solvent

• Solute-Solvent

Step 1 Solution Formation

• Solute-Solute IMFs must be overcome

• Energy is consumed

Step 2 Solution Formation

• Solvent-Solvent IMFs must be overcome

• Energy is consumed

Step 3 Solution Formation

• Solvation-Solute-Solvent attractive forces are established

• Energy is released

Nonpolar solvents dissolve

Nonpolar solutes

Polar solvents dissolve

Polar solutes and many ionic solutes

Electrolytes

Substances that dissolve in water and undergo a physical or chemical change to produce ions

Nonelectrolytes

Substances that do not produce ions when dissolved in water

Strong Electrolytes

Essentially 100% of the dissolved substance generates ions

Weak Electrolytes

Only a relatively small fraction of the dissolved substance generates ions

Ion-Dipole Attraction

The electrostatic attraction between an ion and a molecule with a dipole (polar molecule)

Dissociation

Water molecules surround and solvate the separated ions

Competing forces influence the extent of solubility…

Of an ionic compound

If Ion-Dipole attraction is the dominant force, then

The compound it “pulled” into solution and has very high water solubility

The force of attraction between oppositely charged ions

If this is the dominant force then the compound tends to remain un-dissolved and in the solid state, and have very low water solubility