chapter 01 structure determines...

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©2012 Gregory R Cook

Chapter 01Structure Determines

Properties

CHEM 341: Spring 2012

Prof. Greg Cook

cook.chem.ndsu.nodak.edu/chem341

Thursday, January 12, 12

http://cook.chem.ndsu.nodak.edu/chem341




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Structure Determines Properties

• Mostly a review of general chemistry

• Atomic and Molecular Structure

• Bonding

• Polarity and Properties

• Acid/Base concepts

2




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Structure

• The sequence of connections that defines a molecule, including the spatial orientation of these connections

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Atoms, Electrons and Orbitals

Section 1.1




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Subatomic Particles

• Nucleus is made up of Protons and Neutrons

• mass of Proton = 1.6726 x 10-27 kg

• mass of Neutron = 1.6760 x 10-27 kg

• Surrounding the Nucleus are electrons

• mass of Electron = 9.1096 x 10-31 kg << proton

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Atomic Structure

6

+-




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Atomic Structure

• Electrons surround nucleus in orbitals

• Atomic Number (Z) = # protons in nucleus

• Mass Number (A) = # protons + # neutrons

• Atomic Weight = average mass of a large number of atoms

7

+

-

XAZ H11 C12612.01071.0079




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Wave Function

• Electrons have properties of both Particles and Waves

• Quantum Mechanics help us understand the structure and behavior of electrons

• Schrödinger Wave Equation

• Describes the energy of anelectron in an atom

• Wave functions ψ (psi)

8

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Heisenberg uncertainty principle

• We can’t tell exactly where an electron is

• but we can tell where it will most likely be

• Probability of finding an electron at a particular spot relative to the nucleus is given by ψ2 (psi)2

9

+




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Orbitals

• Wave functions are also called orbitals

• each orbital characterized by 3 quantum numbers

10

• n : principle quantum number

• l : angular momentum quantum number

• ml : magnetic quantum number




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Quantum Numbers

• principle quantum number n

• an integer

• determines major part of orbital energy - the shell

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s orbitals

• s orbitals are spherical in shape

• energy increases with n

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Quantum Numbers

• Angular momemtum l determines the shape of the orbital

• for a given value of n : l = 0, 1, 2, , , n - 1

• l = 0 : s

• l = 1: p

• l = 2 : d

• l = 3 : f

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p Orbitals

• p orbitals are shaped like dumbells with a node in between the lobes (n = 2 and higher)

• Three orbitals with the same energy

14




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4th Quantum Number - Spin

• Each electron also has a spin quantum number ms

• +½ and -½

• Pauli Exclusion Principle - two electrons may occupy the same orbital only when they have opposite or “paired” spins.

• No orbital can contain more than 2 electrons

• H, He, Li

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Electron Configuration

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Periodic Table - Periods

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Second Period Electron Configurations

18

Z

Li 3

1s 2s 2p

C 6

N 7

O 8

Ne 10




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Second Period Electron Configurations

19

Z

Li 3

1s 2s 2p

C 6

N 7

O 8

Ne 10




cook.chem.ndsu.nodak.edu/chem341


Ionic BondsSection 1.2




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Ions

21

1s22s1 1s22s22p5




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Ionization

22

1s22s1 1s22s22p5

Na Cl• ••• •• •

•




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Ionization

23

1s22s0 1s22s22p6

Na Cl• ••• •• ••+ -




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Ionic Bond

24

Na Cl• ••• •• ••+ -

• An ionic bond is a force of attraction between oppositely charged species (ions).

• Ionic bonds are common in inorganic compounds but are more rare in organic compounds.




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Ionization

25

1s22s1 1s22s22p5

1s22s22p2




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Covalent Bonds, Lewis Structures and the Octet

RuleSection 1.3




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Covalent Bonds

• In 1916 G. N. Lewis proposed that atoms combine in order to achieve a more stable electron configuration.

• Maximum stability results when an atom is isoelectronic with a noble gas.

• An electron pair that is shared between two atoms constitutes a covalent bond.

27

1s22s0 1s22s22p6

Na Cl• ••• •• ••+ -

He Ne




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Covalent Bonding in H2

• Two hydrogen atoms each have one electron

• Instead of ionizing, they come together to share both electrons between them

• Sharing the electron pair allows both hydrogen atoms to have a filled orbital

28

•H • H

•H • H




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Covalent Bonding in F2

• Two fluorine atoms each have 7 valence electrons

• They can share them in a covalent bond

• Each fluorine atom has the same electron configuration as Ne (8 electrons)

29

•F • F

F •• F••••• •

• •• •

• •

••

• •• • ••

• •• •




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Octet Rule

• When forming compounds, atomswill gain, lose or share electrons togive a stable electron configurationcharacterized by 8 valence electrons

• The octet rule is most useful in cases involving covalent bonds to C, N, O and F

30

F •• F••••• •

• •• •

• •




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Example with CF4

• Carbon has 4 valence electrons, F has 7

• The Lewis structure for CF4

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•C • F•

•• ••

• •• •

•C •F••

• ••

• •• •

••F••• •

•

••F••

•• •

••F•• •

• •




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Example with CF4

• It is more common to represent a covalent bond (shared pair of electrons) with a line

32

•C •F•

•• ••

• •• •

••F••

• ••

••F••

•• •

••F•

• •• • F C F

F

F=




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Example with CF4

• We often don’t write the lone pairs (Kekulé structures)

• You should know how many lone pairs an atom should have!

33

F C F

F

F




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Double and Triple BondsSection 1.4




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Inorganic Examples

35

O C OO C O

H C NH C N




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Organic Examples

36

C CH

H

H

H

C C HH

Ethylene (Ethene)

Acetylene (Ethyne)




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Polar Covalent Bonds, Electronegativity, and

DipolesSection 1.5




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Electronegativity

• Electrons are not always shared equally between atoms in a covalent bond

• Electronegativity is a measure of the ability of an element to attract electrons toward itself when bonded to another element

• An electronegative element attracts electrons gathering negative charge

• An electropositive element releases electrons gathering a positive charge

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Pauling Electronegativity Scale

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Electronegativity

• Increases from left to right and bottom to top (decreases going down a group)

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EN Generality

• The greater the difference in EN between two bonded atoms, the more polar the bond

• Nonpolar bonds connect atoms with the same EN

41

H H Cl Cl O O




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Polar Covalent Bonds

• Polar Covalent Bonds connect atoms that have different EN

• partial negative charges on atoms with higher EN and partial positive charges on atoms with lower EN

• Bond dipoles pointing from ∂+ toward ∂-

42

H F H O H O C Oδ+ δ- δ- δ- δ-δ+ δ+ δ+




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Electrostatic Potential Maps

• Electrostatic potential maps show the charge distribution within a molecule

43

H Fδ+ δ-




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Polar vs Ionic Bonds

44

A A A B A Bδ+ δ-

Covalent Ionicpolar covalent

δEN >2 IONIC BONDSδEN <2 COVALENT BONDS




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Formal ChargeSection 1.6




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Formal Charge

• The Formal Charge is the calculated charge for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs

• When atoms have more or less number of bonds than their valency requires, they must have a charge that is not zero

• A formal charge is a way of keeping track of electrons and who owns them in a molecule

FC = (# valence electrons) - (# bonds) - (# nonbonded electrons)

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Example HNO3

47

O N

O

OH




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More Examples

48

N

H

H

H

H

B

F

F

F

F




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Structural Formula of Organic Molecules

Section 1.7




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Constitution

• The order in which the atoms of a molecule are connected is called its constitution or connectivity

• The constitution of a molecule must be determined in order to write a Lewis (or Kekulé) structure

• Isomers are different compounds with the same molecular formula

• Constitutional (structural) isomers differ in the order that atoms are connected

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Constitutional Isomers

• Take, for example, a compound with the formula C2H6O

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51

C C O

H

H

H

H

H

H




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C C O

H

H

H

H

H

H C O C

H

H

H

H

H

H




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• How many isomers of C3H7Cl are there?

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• How many isomers of CH3NO2 are there?

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