2013 Lect2a Physical Properties and Structure Relationship1

7/29/2019 2013 Lect2a Physical Properties and Structure Relationship1

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The structures of organic molecules vary widely.Within this variety there are similarities in structure that result in

similarities in chemical, physical, and/or biological properties.

If one can understand the properties of a compound based on its

structure, then we might be able to predict properties of a compound withsimilar structure.

The power to predict the properties of a substance based on its similarity

to other substances of known structure is a powerful tool for organic

chemists, especially so for the discovery of new pharmaceuticals.

It is useful, therefore, for you to learn to recognize

various categories of structural relationships.

PHYSICAL PROPERTIES RELATIONSHIPS


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Structure Basic 1

of Organic Molecules


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STRUCTURE BASIC 1

Shape of Molecules: Atomic Orbitals, Covalent

bond, sigma bond and pi bond, Hybridization of

carbon orbital : Sp3 (single bond, tetrahedral), Sp2(double bond, planar),sp (triple, linier),

Non covalent Interaction : dipole interaction,

London dispersion, hydrogen bond,Physical properties : Polarity, Solubility, Boiling &

Melting point,,


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:

Atomic Orbitals, Covalent bond, sigma bond and pibond,

Shape of Molecules : Hybridization of carbon orbital

: Sp3 (single bond, tetrahedral), Sp2 (double bond,

planar),sp (triple, linier), 3-D drawing,

Rotationaround single bond : Conformer , Newman

Projection, Isomer ,

Rigid structure of double bond : Stereoisomer

(Cis/trans and E/Z structure designation ), Cahn-

Ingold-Prelog Rules, Enantiomerand R/S structure

designation),


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Objectives

Know how atomic orbitals overlap to form molecular orbitals

Understand orbital hybridization

Using the VSEPR model, predict the geometry of molecules

Understand the formation of molecular orbitals


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The basic unit of matter is an atom. It consists dense centralnucleus surrounded by negatively charged electrons.

The nucleus contains a mix of positively charge proton and electrically

neutral neutrons.

The electrons of an atom are bounded to the nucleus by

electromagnetic force

The electrons determine the chemical properties of an

element .

ATOM AND ITS ELECTRONS


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In quantum mechanics, the orbiting electron around a nucleus could not

fully describe as particles, in the sense of a planet orbiting the sun.

It is treated by the wave-particle duality, so that the orbital exist as

standing waves.

The atomic orbitals are supposed to describe the shape of the

atmosphere(electron) distributed around a relative tiny planet

(nucleus)

ATOMIC ORBITALS


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In quantum mechanics , an atomic orbital is a MATHEMATICAL FUNCTIONthat describes the wave-like behavior the electrons in an atom

The function can be used to calculate the probability of finding any

electron of an atom in any specific region around atoms nucleus.

The simple names s orbital, p orbital, d orbital

and f orbital refer to subshell s, p, d and f

respectively.

ATOMIC ORBITALS


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Graphical representation of the 1s atomic orbital of H


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Waves that are in phase : add together, increase amplitude.

Waves that are out of phase : cancel out

on the same atom is HYBRIDIZATION

between different atoms is BOND FORMATION

ORBITALS are interact each other

like Waves


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SIGMA BONDING

sigma bonds ( bonds) are formed by

head-on overlapping between atomic orbitals

A bond may be formed by s-s,p-p, s-p, or

hybridized orbital overlaps.

a -bond is symmetrical with respect to

rotation about the bond axis


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Shape of Methane

TETRAHEDRAL


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The molecular geometry depends on the electron

pair geometry (i.e. the total number of electronpairs on the central atom).

the shapes of the molecule is explained by Valence

Shell Electron Pair Repulsion (VSEPR) model, that

states that electron pairs arrange themselves a

maximum distance apart.


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GEOMETRY OF METHANEbase on VSEPR MODEL .

in methane, the central carbon atom is bonded to four

other identical atoms

The electron pairs ( electron bond) strive to be as farapart from other electron pairs as possible.

This arrangement places the four identical atoms, the

hydrogens, toward the corners of a regular

tetrahedron with the atom they are bonded to.


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FOR CARBON WITH 6 ELECTRON

What its electron configuration?


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order for filling the orbitals:


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Hund's Rule applied to the filling of the atomic orbitals of carbon

Based on this electronconfiguration,

Carbon has 4 outershell electron

2 in s orbital and 2 in p orbital

Which in different level of energy

So If C bonded with 4 H, the bonds

would have been non identical

Since actually they are identical,

must the for orbital (s 2 and 2 p) arefused or combine to form a new

orbital or hybrid orbital

This fenomena is called hibrization


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SP3 hybridization of C

Electron configuration of carbon (a) before and (b) after hybridization.

Note that the energy level of the hybrid orbitals is between that of the 2s and

that of the 2p orbitals.


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Mixing, or hybridization, of one s orbital with threep orbitals produces four

sp3 orbitals.

Each of the sp3 orbitals has 25% s character and 75%p character.


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Bonding in CH4


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PLANAR SHAPE OF C2H

4

CH

H

H

H

sp2

120 C


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Both carbons in ethylene are sp2hybridized and

bonded to each other

SP2 hybridization CARBON ATOM

in Ethylene (CH2=CH2)


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Pi BONDING

Sideways overlap of parallelp orbitals.


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Orbital Overlap in Ethylene

A double bond consists of A SIGMA BOND and API BOND.


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ORBITAL OVERLAP INI ACETYLENE

linier


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SP hybridization


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The molecular orbitals in acetylene are

formed from overlap of thep orbitals of two

sp hybridized carbon atoms.

SP hybridization


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LINIER SHAPE OF MOLECULE P INI

ACETYLENE

=>


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Orbital Hybridization

BONDS SHAPE HYBRID REMAIN

2 linear sp 2 ps

3 trigonal sp2 1 p

planar

4 tetrahedral sp3 none


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MOLECULAR SHAPES

of oxygen/nitrogen containing compound


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3 Sigma bonds and

one lone pair

:NH3

N

H H

H

Geometry? pyramidal WHY?

See nucle, not electrons

Case 1Amonia


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Case 2

2 sigma bonds and2 lone pair

Water and

CH3OCH3

O

H

H

BENT


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Central Atoms with Single-Bond Pairs and Lone Pairs

Lone pairs of electrons are more repulsive than bonding pair

electrons.

Central Atoms Multiple Bonds


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Central Atoms Multiple Bonds

To predict geometry, count the multiple bond as a single

bonding pair of electrons.


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Case 3

Etanal

One s orbital +

Two pi orbitals

trigonal planar:

120o bond angles.

H3C

O

H

Carbon has 3 sigma

and one pi bond

3 sp2 orbitals


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Central Atoms Surrounded by Only Single-Bond Pairs


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Central Atoms with Single-Bond Pairs and Lone Pairs


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Carbon has sp Hybrid OrbitalsCombine one s 2 with one p 2

sigma orbital

H-C N

Linear 180 o

Note: both C and N are sp hybridized

H-C N

Case 4

HYDROGENCIANIDE


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O

C

O H

H

H

NH

Hsp3

sp

3

sp3

sp2

C

identification of

bonding and geometry

in Glycine


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O

C

O H

H

H

NH

Hsp3

sp

3

sp3

sp2

C

Bonding in

Glycine


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O

C

O H

H

H

NH

Hsp3

sp

3

sp3

sp2

C

Bonding in

Glycine


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O

C

O H

H

H

NH

Hsp3

sp3

sp3

sp2

C

Bonding in

Glycine


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Bonding in

Glycine

O

C

O H

H

H

NH

Hsp3

sp

3

sp3

sp2

C


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QUIZES

Predict the hybridization, geometry, and

bond angle for each atom in the

following molecules: NH2NH2

CH3-CC-CHO

CH3 C

O

CH2

_

Caution! You must start

with a good Lewis

structure


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Tendency of a molecule, or compound, to be attracted or

repelled by electrical charges because of an asymmetrical

arrangement of atoms around the nucleus.

POLARITY


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BOND POLARITY

When two dissimilaratoms bond,

the electrons are shared unequally and result in a partialnegative charge (d-) on one atom and a partial positive

charge (d+) on the other atom.

Bond polarity depends on the

electronegativity of the atoms.

POLAR COVALENT BOND


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ELECTRONEGATIVITY

Electronegativity (c) of an atom is the atoms ability toattract electrons to itself.

c increases

cdecreases


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Measure of POLARITY

Product of the magnitude of the partial charges and the distance

by which they are separated

C O3,52,5

Bond Dipole Moments

(debyes)

= 4.8 x d(electron charge) x d(angstroms

d+ d-

POLAR MOLECULES orientation


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POLAR MOLECULES orientation

under electrical field

HCl highly polar.Between two electric plates with the field off, the molecules lie every which way

(oriented randomly)

With the field on, however, they become oriented with their negative ends facing

the positive plate and their positive ends facing the negative plate.

blue indicates high

electron density and

red indicates low


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RELATIVE STRENGTH OF MOLECULAR

POLARITY

=>

The bigger the partial charges, the more the polarity of the bo


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Lone pairs of electronscontribute to thedipole moment.

=>

Vector sum of the bond

dipole moments.

INFLUENCE MOLECULAR GEOMETRY ON POLARITY

Molecular Polarity


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Molecular Polarity

A molecules polarity is a function of the molecular shape

and the electronegativity of the atoms.


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PHYSICAL PROPERTIES

OF HYDROCARBONS


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PHASES OF MATTER

Matter can exist in four phases (or

states), solid, liquid, gas, and plasma,

A solid is matter in which the

molecules are very close together

and cannot move around

In a liquid, the molecules are

close together and move around

slowlyIn A gas the molecules are widely

separated, move around freely,and move at high speeds.

A plasma is a gas that is composed of free-floating ions (atoms

stripped of some electrons - positively charged) and free electrons


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In a liquid the molecules are packed closely together with many random movements

possible as molecules slip past each other.

As the temperature increases, the kinetic energy increases which causes increasing

molecular motion (vibrations, slipping pass each other).

Eventually, motion becomes so intense that the forces of attraction between themolecules is disrupted to the extent the molecules break free of the liquid, become a gas.


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Water Liquid to Water Gas:

This animation shows how water molecules areable to break the forces of attraction i.e. the

hydrogen bonds to each other and escape as the

gas molecule. This is what is happening inside the

gas bubble as it is rising to the surface to breakand release the water gas molecules.


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The temperature at which the vapor pressure of the liquid

equals the environmental pressure surrounding theliquid.

BOILING POINT

At that T, the Pvapor of the liquid becomes sufficient to

overcome Patm and allow bubbles of vapor to form inside

the bulk of the liquid.

Pure, crystalline solids have a characteristic melting

point, the temperature at which the solid melts tobecome a liquid. The transition between the solid and

the liquid is so sharp for small samples of a pure

substance that melting points can be measured to

0.1o

C. The melting point of solid oxygen, for example, is- o

MELTING POINT


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Principle:

The greater the forces of

attraction the higher the boilingpoint or the greater the polarity

the higher the boiling point.

Molecules which strongly interact or bond with each other

through a variety of intermolecular forces can not move easily

or rapidly and therefore, do not achieve the kinetic energynecessary to escape the liquid state. Therefore, molecules with

strong intermolecular forces will have higher boiling points. This

is a consequence of the increased kinetic energy needed to

break the intermolecular bonds so that individual molecules

may escape the liquid as gases.


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WWU Chemistry

the longer the straight chain, the higher the boiling

point -- van der Waals forces

isomers that are branched have lower boiling points

hydrogen bonding increases boiling points

HYDROCARBON BOILING POINT


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the boiling points of straight chain alkanes are related to

the number of carbon atoms in their molecules.

Increased intermolecular attractions are related to the

greater molecule-molecule contact possible for larger

alkanes.

Branched alkanes have less surface area

contact,

so weaker intermolecular forces.


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BOILING POINTS OF ALKANES

The boiling point vary directly with the number of

carbon atomvary inverselydirectly proportional

inversely proportional

A l f th b th b ili i t i 20 30 C f


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Alkane FormulaBoiling point

[C]

Melting point

[C]

Density

[gcm3] (at

20 C)

Methane CH4 -162 -183 gas

Ethane C2H6 -89 -172 gas

Propane C3H8 -42 -188 gas

Butane C4H10 0 -138 gas

Pentane C5H12 36 -130 0.626(liquid)

Hexane C6H14 69 -95 0.659(liquid)

Heptane C7H16 98 -91 0.684(liquid)

Octane C8H18 126 -57 0.703(liquid)

Nonane C9H20 151 -54 0.718(liquid)

Decane C10H22 174 -30 0.730(liquid)

Undecane C11H24 196 -26 0.740(liquid)Dodecane C12H26 216 -10 0.749(liquid)

Icosane C20H42 343 37 solid

Triacontane C30H62 450 66 solid

Tetracontane C40H82 525 82 solid

Pentacontane C50H102 575 91 solid

As a rule of thumb, the boiling point rises 2030 C for

each carbon added to the chain

MELTING POINTS OF ALKANES


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MELTING POINTS OF ALKANES

Branched alkanes pack more efficiently into

a crystalline structure, so have higher m.p.

=>


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What is the state (solid, liquid, gas) of alkane

compound at 20C


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2 methylpropane and n butane which boil at higher


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(2-methylpropane) and n-butane (butane), which boil at 12 and 0 C

2-methylpropane and n-butane which boil at higher

temperatur and why

2,2-dimethylbutane and 2,3-dimethylbutane which

boil at higher temperature?

2,2-dimethylbutane and 2,3-dimethylbutane which boil at 50 and 58 C,

two molecules 2,3-dimethylbutane can "lock" into each other better than the cross-shaped 2,2-dimethylbutane, hence the greater van der Waals forces.

due to the greater surface area in contact, thus the greater van der W aals forces


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Match each boiling point to the appropriate C7H16

isomer: 98.4 C, 92.0 C, 79.2 C.


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m.p., b.p., and solubility; esp. for solids and

liquids are influenced by the strength of

attractions between molecules or

INTERMOLECULAR INTERACTION


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INTERMOLECULAR FORCES

Classification depends on structure.

Dipole-dipole interactions

London dispersionsHydrogen bonding

=>


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DIPOLE-DIPOLE FORCES

(permanent dipole)

Between polar molecules

Positive end of one molecule aligns with

negative end of another molecule.

Lower energy than repulsions, so net

force is attractive.

Dipole Dipole INTERACTION


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Dipole-Dipole INTERACTION

=>

Larger dipoles cause higher boiling points and higher heats of

vaporization.


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LONDON DISPERSION FORCE

instantaneously induced dipoles

Between nonpolar molecules

Temporary dipole-dipole interactions

Larger atoms are more polarizable..


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Dispersions

=>

L d Di i


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London Dispersions

Branching lowers b.p. because of

decreased surface contact between

molecules.

=>

CH3 CH2 CH2 CH2 CH3

n-pentane, b.p. = 36C

CH3 CH

CH3

CH2 CH3

isopentane, b.p. = 28C

C

CH3

CH3

CH3

H3C

neopentane, b.p. = 10C

B h d Alk


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Branched Alkanes

Lower b.p. with increased branching Higher m.p. with increased branching

H

CH3CH

CH3

CH2 CH2 CH3

bp 60Cmp -154C

CH3CH

CH3

CHCH3

CH3bp 58Cmp -135C

=>

bp 50Cmp -98C

CH3 C

C 3

CH3

CH2 CH3


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What can you infer about intermolecular attractions indecane compare to those in pentane?

Intermolecular forces also help explain other liquid

properties such as viscosity and freezing points.

Based on their intermolecular attractions, try to rank

pentane, octane, and decane in order of increasing

viscosity.

Assign "1" to the least viscous ("thinnest") of the three.


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CH3CH2OCH2CH3 and CH3CH2CH2CH2OH

Are they isomer? What isomer type?

Which one has higher boiling point /density

BP= 34.5C BP= 117.2C

d= 0.7138 g/mL d= 0.8098 g/mL

insoluble in water soluble in water


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Hydrogen Bonding

Strong dipole-dipole attraction

Organic molecule must have N-H or O-H.

The hydrogen from one molecule isstrongly attracted to a lone pair of

electrons on the other molecule.

O-H more polar than N-H, so strongerhydrogen bonding =>

H B d


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H Bonds

=>


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More normal behavior is seen in dimethyl ether

(CH3)2O which has no hydrogen bonds possible.

If waterbehaved as a

normal polar molecule itwould have boiled at

about - 100 C (shown in

red). Instead, water boils at

+100 C, which is veryabnormal.

The major reason is the strongattractions afforded by the

hydrogen bonds.

It takes a lot more kinetic energy in

increasing T to break the H-bonds

to free the water molecules as the

gas.

hid b di i th B ili P i t


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hidrogen bonding increase the Boiling Points

CH3

CH2

OH

ethanol, b.p. = 78C

CH3 O CH3

dimethyl ether, b.p. = -25C

rimethylamine, b.p. 3.5C

N CH3H3C

CH3

propylamine, b.p. 49C

CH3CH2CH2 N

H

H

ethylmethylamine, b.p. 37C

N CH3CH3CH2

H

=>

CH3 CH2 OH

ethanol, b.p. = 78C ethyl amine, b.p. 17

CH3 CH2 NH2

polar dipole forces exerts a very strong effect to keep

the molecules in a liquid state until a fairly high

temperature is reached.


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Soluble and insoluble. Alcohol is soluble in water; whenadded to water, it forms a clear solution.

Oil is insoluble in water; when added to water, the two

liquids form separate layers.

SOLUBLE AND INSOLUBLE

SO


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SOLUBILITY

Like dissolves like

Polar solutes dissolve in polar solvents.

Nonpolar solutes dissolve in nonpolarsolvents.

Molecules with similar intermolecular

forces will mix freely.=>


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SOLUBILITY is an ability of a substance to dissolve.

In the process of dissolving, the substance which isbeing dissolved is called a solute and

the substance in which the solute is dissolved is called a

solvent.

A mixture of solute and solvent is called a solution.

SOLUBILITY

is understood as a maximum amount of solute that dissolves

in a solvent at so called equilibrium.

An equilibrium is a state where reactants and products reach

a balance - no more solute can be dissolved in the solvent in

the set conditions (T,P). Such a solution is called a

saturated solution.


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CH3CH2OCH2CH3 and CH3CH2CH2CH2OH

Are they isomer? What isomer type?

Which one dissolves in water? Both?

BP= 34.5C BP= 117.2C

d= 0.7138 g/mL d= 0.8098 g/mL

insoluble in water soluble in water


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Why some things dissolve and some do not?

When dissolving, new bonds (Intermolecular forces)

between solvent and solute are created. During thisprocess energy is given off. The amount of this energy is

sufficient to brake bonds between molecules of solvent

and between molecules of solute.

If the energy resulted from the intermolecular forces is not

enough to disrupt molecular bond of solvent and

molecular bond of solute, then the solute will not dissolve.

PROCESS OF DISSOLVING


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Ionic Solute withPolar Solvent

Energy released is enough

=>


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Ionic Solute with Nonpolar Solvent

=>

Energy released is not enough

Nonpolar Solute with


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Nonpolar Solute with

Nonpolar Solvent

=>Energy released is enough

Nonpolar Solute


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Nonpolar Solute

with PolarSolvent

=>

Energy released is not enough

II Structure and


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II. Structure and

Physical Properties: RX Bonding in alkyl halides:Bond lengths increase as size of halogen

increases:

C-I > C-Br > C-Cl > C-F

polar covalent bonds

sp3 hybridized

tetrahedral

II St t d Ph i l P ti RX


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II. Structure and Physical Properties: RX

Polarity:

dipole-dipole interactions

dipole-induced dipole interactios

Boiling points are highest for iodine containing

compounds, lowest for fluorine compounds of similar

molecular weight: R-I > R-Br> R-Cl > R-F Polarizability: ease with which electron distribution is

affected by nearby electric field

accounts for induced dipole interactions

Insoluble in water Higher density than water


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II. Structure and Physical Properties: R-OH

Polar

sp3 hybridized

bent

hydrogen-bonding

relatively high boiling points

RO

H

R O

H

O

R

H O

R

H

H

OR


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Boiling Points:

Alcohols have high boiling pointsBoiling Point Intermol.

Attractions

CH3CH2CH3 - 42

o

C London

CH3OCH3 - 24o C Dipole-

dipole

CH3CH

2OH 78o C Hydrogen

bonding


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Solubility:

alcohols having from 1 to 4 carbons:

miscible with water

alcohols having from 4 to 6 carbons:slightly soluble

alcohols having greater than 6 carbons

are insoluble

pol ols (diols triols etc ) are er

2013 Lect2a Physical Properties and Structure Relationship1

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