Chemistry Chapter 10 & 11

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Writing Lewis Structures

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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

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Electronegativity

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

  • F is highest

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Electrons are shared unequally in what type of bond?

Polar covalent

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Type of bond is determined by what?

Difference in electronegativity

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Electronegativity < 0.4

Pure (nonpolar) covalent

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Electronegativity 0.4-2.0

Polar covalent

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Electronegativity > 2.0

Ionic

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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

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<p>H2O</p>

H2O

  • Dipole moment

  • Polar molecule

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<p>SO2</p>

SO2

  • Dipole moment

  • Polar molecule

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<p>CO2</p>

CO2

  • No dipole moment

  • Nonpolar molecule

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<p>CH4</p>

CH4

  • No dipole moment

  • Nonpolar molecule

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Non-polar

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

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<p>Linear</p>

Linear

Non-polar

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<p>Trigonal Planar</p>

Trigonal Planar

Non-polar

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<p>Tetrahedral</p>

Tetrahedral

Non-polar

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<p>Trigonal Bipyramidal</p>

Trigonal Bipyramidal

Non-polar

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<p>Octahedral</p>

Octahedral

Non-polar

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Polar

Shapes that are not symmetrical are polar

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<p>Bent</p>

Bent

Polar

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<p>Trigonal Pyramidal</p>

Trigonal Pyramidal

Polar

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<p>Distorted Tetrahedral (see-saw)</p>

Distorted Tetrahedral (see-saw)

Polar

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Intermolecular Forces

Attractive forces between molecules

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Intramolecular Forces

Hold atoms together in a molecule

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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)

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“Measure” of intermolecular force

  • Boiling point

  • Melting point

  • ΔHvap

  • ΔHfus

  • ΔHsub

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Kinetic Energy vs Interactive Forces

  • KE lowest, IF highest

  • KE intermediate, IF intermediate

  • KE highest, IF lowest

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The boiling point increases as the __ increases

Atomic number

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

Greater the amount of energy needed to separate them

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The greater the energy required to separate these particles, the

Higher the boiling points

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Dispersion (London) Forces (LDF)

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

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Temporary (induced) Dipole

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

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Polarizability

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

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All molecules and atoms have…

Dispersion forces

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Dispersion

Momentary shift in e- density

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Factors affecting strength of dispersion

  • Size of atoms/molecules

  • Shape of molecules

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Dispersion - Size of atoms/molecules

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

  • Dispersion increases with polarizability

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Dispersion - Shape of molecules

  • Increased surface area = increased interactions between molecules

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

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Effect of Shape on Dispersion

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

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Dipole-Dipole Interactions

Attractive force between polar molecules

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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

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Ion-Dipole Interaction

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

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Sphere of Hydration

  • Cluster of water molecules surrounding an ion in aqueous solution

  • Sphere of solvation if solvent is other than H2O

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Dipole-Induced Dipole

  • Proximity of polar molecule causes dipole-induced dipole

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Relative Strengths of Intermolecular Forces

  • Weakest - Dispersion Forces

  • Dipole Induced Dipole Forces

  • Dipole-Dipole Forces

  • Hydrogen Bonding

  • Strongest - Ion Dipole Forces

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Solubility

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

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Like dissolves like

  • Ionic/polar solutes will be soluble in polar solvents

  • Nonpolar solutes will be soluble in nonpolar solvents

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Combination of Forces

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

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Solubility decreases are relative energy of…

Hydrogen bonding decreases and dispersion increases

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Hydrophobic (water-fearing)

Interaction that repels water, diminished water solubility

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Hydrophilic (water-loving)

Interaction that attracts water, promotes water solubility

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Equilibrium Vapor Pressure

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

  • H2O (L) → H2O (g)

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Dynamic Equilibrium

Rate of Condensation = Rate of Evaporation

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Molar Heat of Vaporization (ΔHvap)

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

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Clausius-Clapeyron Equation

  • ln P = [ (-ΔHvap) / (RT) ] + C

  • P = equilibrium vapor pressure

  • T = temperature (K)

  • R = gas constant (8.314 J/K x mol)

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Clausius-Clapeyron Equation at two temperatures

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

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Boiling Point

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

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Normal Boiling Point

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

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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

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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

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Molar Heat of Fusion (ΔHfus)

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

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Solid-Gas Equilibrium

H2O (s) → H2O (g)

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Molar Heat of Sublimation (ΔHsub)

The energy required to sublime 1 mole of a solid

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Hess’s Law

ΔHsub = ΔHfus + ΔHvap

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Cohesion

The attraction of like molecules due to intermolecular forces

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Adhesion

The attraction of different molecules due to intermolecular forces

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Cohesive Forces are responsible for…

Some properties of liquids such as surface tension and viscosity

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Phase Diagram

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

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Equilibrium Lines

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

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Supercritical Region

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

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Critical Temperature (Tc)

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

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Critical Pressure (Pc)

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

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Phase Diagram - Triple Point

Temperature/pressure where all 3 phases of a substance coexist

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Phase Diagram - Critical Point

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

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Phase Diagram - Supercritical Fluid

A substance above its critical temperature and pressure

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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)

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T final > T initial

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

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T final < T initial

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

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Molar Heat of Fusion (ΔHfus)

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

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Enthalpy in Change of State equations

  • q = nΔHfus

  • q = nΔHvap

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Solutions are __ made up of two or more substances

Homogenous mixtures

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Solute

Substance being dissolved

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Solvent

  • Substance that is doing the dissolving

  • Typically, present in largest amount

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Solid Solution

Metal alloys (two or more metals)

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Gas Solution

Example: air - made up of multiple gases

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

  • Solute-Solute

  • Solvent-Solvent

  • Solute-Solvent

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Step 1 Solution Formation

  • Solute-Solute IMFs must be overcome

  • Energy is consumed

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Step 2 Solution Formation

  • Solvent-Solvent IMFs must be overcome

  • Energy is consumed

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Step 3 Solution Formation

  • Solvation-Solute-Solvent attractive forces are established

  • Energy is released

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Nonpolar solvents dissolve

Nonpolar solutes

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Polar solvents dissolve

Polar solutes and many ionic solutes

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Electrolytes

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

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Nonelectrolytes

Substances that do not produce ions when dissolved in water

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Strong Electrolytes

Essentially 100% of the dissolved substance generates ions

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Weak Electrolytes

Only a relatively small fraction of the dissolved substance generates ions

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Ion-Dipole Attraction

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

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Dissociation

Water molecules surround and solvate the separated ions

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Competing forces influence the extent of solubility…

Of an ionic compound

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If Ion-Dipole attraction is the dominant force, then

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

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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

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