cm1501-week 1a & 1b-struc bonding

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CM1501 Week 1A & 1B – Structure & Bonding

For lab matters, visit IVLE module

CM1501_LAB/CM1501

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Organic Chemistry, Second EditionJanice Gorzynski Smith

University of Hawai’i

Textbook

Monday Lecture: 45 mins

DR Zeng Huaqiang: [email protected] Ong Yue Ying: [email protected]

Teaching Staff

Thursday Lecture: 90 mins(no break)

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Chapter 1: Structure and BondingChapter 2: Acids and BasesChapter 3: Introduction to Organic Molecules & Functional GroupsChapter 4: AlkanesChapter 5: StereochemistryChapter 6: Understanding Organic ReactionsChapter 7: Alkyl Halides and Nucleophilic SubstitutionChapter 8: Alkyl Halides and Elimination SubstitutionChapter 9: Alcohols, Ethers and EpoxidesChapter 10: AlkenesChapter 11: AlkynesChapter 16: Conjugation, Resonance and DienesChapter 17: Benzene and Aromatic CompoundsChapter 22: Carboxylic acids and Their DerivativesChapter 19: Carboxylic Acids and the Acidity of the O-H BondChapter 21: Aldehydes and KetonesChapter 25: Amines and whole class tutorial

CM1501 Course ScheduleWeek 1Week 2

Week 3

Week 4

Weeks 5-6

Week 7Week 8

Week 12

Week 13

Weeks 9-11

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Mid-Term Exam 30 %

Final Exam 50 %

Practical Section 20 %

Total 100 %

CM1501 Assessments

Final Exam: Monday, 03-12-2012 (1:00-3:00 pm)

CLOSED BOOK EXAMINATION

Mid Exam: 04-10-2012, 4:10-5:50pm at MPSH1, section A+B

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What is Organic Chemistry?

• Organic chemistry is the study of carbon compounds.

• Why is it so special?- 90% of more than 30 million chemical compounds contain

carbon.

- Examination of carbon in periodic chart answers some of

these questions.

- Carbon is group 4A element, it can share 4 valence

electrons and form 4 covalent bonds.

Volume of atom depends on the size of the electron cloud.

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Silicon As An Alternative Element for Life?

Smaller in size and lighter in mass.

Can form double or triple bonds.

Can form alkane hydrocarbons or long carbon chains

CO2 is gas & soluble in water

Larger in size and heavier in mass.

Difficult to form double or triple bonds.

Silanes, the analogs of alkanehydrocarbons, are highly reactive and may decompose easily

SiO2 is non-soluble solid.

SiliconCarbon126

2814

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The nucleus (very dense; very small: 10-15 m; 1 fm) contains positively charged protons and uncharged neutrons.

The electron cloud (10-10m; ~ 2Ǻ) is composed of negatively charged electrons.

The Periodic Table

Neutral atomprotons = electrons

Cationsprotons > electrons

Anionsprotons < electrons

Atomic number Isotope

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A comparison of two isotopes of the element carbon

Atomic weight: the weighted average of the mass of all isotopes reported in atomic mass units (amu)

Carbon as the example: 12.011 = 12 x 98.9% + 13 x 1.1%

The Periodic Table

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The Periodic Table

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181832

8

32http://www.dayah.com/periodic/

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A

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The Periodic Table

• Elements in the same row are similar in size.• Elements in the same column have similar electronic and

chemical properties.

From left to right From top to bottom

Atomic radius

Electronegativity

Ionization energy

Atomic radius

Electronegativity

Ionization energy

Periodic Trends

+-

+-

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The Periodic Table

• Four types: s, p, d, and f

• s and p orbitals most important in organic and biological chemistry.

• There are ONE s orbital, THREE p orbitals, FIVE d orbitals and SEVEN f orbitals.

• Each orbital can hold two electrons.

Atomic Orbital:

A region of space that ishigh in electron density

1s (1)

2s (1)

3s (1)

3d (5)

2p (3)

3p (3)

4f

Energy gradient

Electron fillingpriority

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• Orbitals are grouped in shells of increasing size and energy• Different shells contain different numbers and kinds of orbitals• Each orbital can be occupied by two electrons with opposite spins

The Periodic Table

Orbitals and Shells

1s

2s

3s

3d

2p

3p

Pauli exclusion principle

Hund's rule

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The Periodic Table

Orbitals and Shells

• Orbitals are grouped in shells of increasing size and energy

• Different shells contain different numbers and kinds of orbitals

• Each orbital can be occupied by two electrons with opposite spins

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The Periodic Table

1s2

2s22p6

4s23d104p6

5s24d105p6

2

8

818

18

3s23p6

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2

2+8

2+8+8

2+8+8+18

2+8+8+18+18

2+8+8+18+18+326s24f145d106p6

Noble Gas Configuration

8A

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The Periodic Table

Columns

1s2 2s1

1s2 2s22p6 3s1

1A

1s1

1s2 2s22p6 3s23p6 4s1

1s2 2s22p6 3s23p6 4s23d104p6 5s1

1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s1

Vary the inner shell while keeping outer shell constant

2 8 8 18 18

2 8 8 18

2 8 8

2 8

2

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The Periodic Table

Rows

2 1s2 2s22p6

2 1s2

2s1 2s2 2s22p2 2s22p4 2s22p61s1 1s2

2s22p1 2s22p3 2s22p5

First row

Second rowVary the outer shell while

keeping inner shell constant

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• s orbitals: spherical, nucleus at center• p orbitals: dumbbell-shaped, nucleus at middle• d orbitals: cloverleaf-shaped, nucleus at center

The Periodic Table

1s orbital 2p orbital 3d orbitalNo electron density at the planar nodes

Shapes of Atomic Orbitals

http://winter.group.shef.ac.uk/orbitron/AOs/6s/index.html

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The Periodic Table

Shapes of Atomic Orbitals

X

Y

Z

http://winter.group.shef.ac.uk/orbitron/AOs/6s/index.html

• s orbitals: spherical, nucleus at center• p orbitals: dumbbell-shaped, nucleus at middle• d orbitals: cloverleaf-shaped, nucleus at center

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• Through bonding,

1) two atoms are joined in a stable arrangement.

2) atoms attain a complete outer shell of valence

electrons, a stable noble gas configuration.

• Ionic bonds: transfer of electrons from one element

to another.

• Covalent bonds: sharing of electrons between two

nuclei.

Bonding

Review2

2+8

2+8+8

2+8+8+18

2+8+8+18+18

2+8+8+18+18+32

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Bonding Forces

Ionic Bond

1s22s22p63s23p5

1s22s22p63s1

Change in electronic configuration

• Generally occurs when elements (e.g., Na) on the far left side of the periodic table combine with elements (e.g., Cl) on the far right side, ignoring noble gases.

2+8

2+8+8

2+8+8 (2+8+7)

2+8 (2+8+1)

Na

Cl

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• Hydrogen forms one covalent bond. Each has a filled valence shell of two electrons

Bonding in Molecular Hydrogen (H2)

Bonding Forces

Covalent bond

Lewis structure representation

• First row elements such as hydrogen can accommodate two electrons around it, making it like the noble gas helium.

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Bonding Structure

Noble gas configuration & Octet Rule

Examples: 1) Carbon has four valence electrons (2s22p2), forming four

bonds.

2) Nitrogen has five valence electrons (2s22p3) but formsonly three bonds.

3) Oxygen has six valence electrons (2s22p4) but forms twobonds.

• Second row elements can have no more than eight electrons around them. Atoms such as C, N, O, F with four or more valence electrons form enough bonds to give an octet.

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The Periodic Table

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Elements in Groups 2A and 3A in the Second Row

Elements in the Third Row

Exceptions to the Octet Rule

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In a Lewis structure, a solid line indicates a two-electron covalent bond.

Bonding StructureLewis structures

Lewis structures are electron dot representations for molecules.Three general rules for drawing Lewis structures:

1. Draw only the valence electrons.2. Give each hydrogen two electrons3. Give second-row elements (C, N, O, & F) an octet of electrons, if possible.

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Lewis structures

Bonding Structure

Lewis structures are electron dot representations for molecules.Three general rules for drawing Lewis structures:

1. Draw only the valence electrons.2. Give each hydrogen two electrons3. Give second-row elements (C, N, O, & F) an octet of electrons, if possible.

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• Formal charge is the charge assigned to individual atoms in a Lewis structure.

• Formal charge is used to determine how the number of electrons around a particular atom compares to its number of valence electrons.

• The number of electrons “owned” by an atom is determined by its number of bonds and lone pairs.

• An atom “owns” all of its unshared electrons and half of its shared electrons.

Formal Charge

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The number of electrons “owned” by different atoms is indicated in the following examples:

Example 1: neutral

Example 2: neutral

Example 3: anion

Formal Charge

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Formal Charge

Formal charge observed with common bonding patters

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More than one arrangement of atoms is possible for a given molecular formula.

The above two are valid Lewis structures and both exist. These two compounds are called isomers.

Ethanol and dimethyl ether are constitutional isomers.

Isomers

Structural (constituitional) isomers1. Chain isomerism2. Positional isomerism3. Functional group isomerism

Stereoisomer (spatial isomer)1. Cis-trans isomer/geometric isomer2. Chiral carbon/optical isomer

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Some molecules cannot be adequately represented by a single Lewis structure.

The hybrid is a weighted average of the contributing resonance structures.

Resonance structures are two Lewis structures having the same placement of atoms but a different arrangement of electrons

Neither resonance structure is an accurate representation for the true structure, which is a composite of both resonance forms and is called a resonance hybrid.

Electron delocalization adds stability.

Resonance Structures

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1. Resonance structures are not real. An individual resonance structure does not accurately represent the structure of a molecule or ion. Only the hybrid does.

2. Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another.

3. Resonance structures are not isomers. Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons.

Basic Principles


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A “better” resonance structure is one that has 1) more bonds, 2) fewer charges, and 3) negative charges preferentially localized on more electronnegative elements.


Resonance Hybrids

Major resonance structure

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Drawing Resonance Structures

Rule [1]: Two resonance structures differ in the position of multiple bonds and nonbonded electrons. The placement of atoms and other single bonds always stays the same.

Rule [2]: Two resonance structures must have the same number of unpaired electrons.


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Rule [3]: Resonance structures must be valid Lewis structures: hydrogen must have two electrons and no second-row element can have more than eight electrons.



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• Shows the movement of an electron pair: the tail of the arrow always begins at the electron pair, either in a bond or lone pair. The head points to where the electron pair “moves.”

Example 1:

Example 2:

Resonance StructuresCurved Arrow Notation

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A lone pair is located on an atom directly bonded to a double bond



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CM1501 Week 1A & 1B – Structure & Bonding In the above examples,


An atom bearing a (+) charge is bonded either to a double bond or an atom with a lone pair.


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Two variables define a molecule’s structure: bond length and bond angle.

• Bond length decreases across a row of the periodic table as the size of the atom decreases.

• Bond length increases down a column of the periodic table as the size of an atom increases.

Molecular Dimensionality

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Bond length

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Bond angle determines the shape around any atom bonded to two other atoms.

• The number of groups surrounding a particular atom determines its geometry. A group is either an atom or a lone pair of electrons.

• The most stable arrangement keeps these groups as far away from each other as possible.


Bond angle

Predicting Geometry Based on Counting Groups Around the Central Atom

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Two groups around an atom—


Bond angle

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Three groups around an atom—


Bond angle

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Four groups around an atom—


Bond angle

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Drawing Three Dimensional Structures

• A solid line is used for a bond in the plane.

• A wedge is used for a bond in front of the plane.

• A dashed line is used for a bond behind the plane.


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The molecule can be turned/rotated in many different ways, generating many equivalent representations. All of the followingare acceptable drawings for CH4.

Molecular DimensionalityDrawing Three Dimensional Structures

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Note that wedges and dashes are used for groups that are really aligned one behind another. It does not matter in the following two drawings whether the wedge or dash is skewed to the left or right, because the two H atoms are really aligned.


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In water (H2O), two of the four groups attached to the central O atom are lone pairs. The two H atoms and two lone pairs around O point to the corners of a tetrahedron. The H-O-H bond angle of 105° is close to the theoretical tetrahedral bond angle of 109.5°. Water has a bent shape, because the two groups around oxygen are lone pairs of electrons.


A Nonbonded Pair of Electrons is Counted as a “Group”

Drawing Three Dimensional Structures

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In ammonia (NH3), one of the four groups attached to the central N atom is a lone pair. The three H atoms and the lone pair point to the corners of a tetrahedron. The H-N-H bond angle of 107° is close to the theoretical tetrahedral bond angle of 109.5°. This shape is referred to as a trigonal pyramid.


A Nonbonded Pair of Electrons is Counted as a “Group”

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Methane (CH4)

Ammonia (NH3)

Water (H2O)

In both NH3 and H2O, the bond angle is smaller than the theoretical tetrahedral bond angle because of repulsion of the lone pairs of electrons. The bonded atoms are compressed into a smaller space with a smaller bond angle.


109o

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• Drawing every bond in organic molecule can become tedious.• Several shorthand methods have been developed to write

structures.• Condensed structures typically don’t have C-H or C-C single

bonds shown.

Molecular DimensionalityDrawing Organic Molecules—Condensed Structures

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Drawing Organic Molecules—Condensed Structures

• All atoms are drawn in with two-electron bond lines omitted.• Atoms are usually drawn next to the atoms to which they are

bonded.• Parentheses are used around identical groups bonded to the same

atom.• Lone pairs are omitted.


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Condensed Structures

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Skeletal structures

• Assume there is a carbon atom at the junction of any two lines or at the end of any line.

• Assume there are enough hydrogens around each carbon to make it tetravalent.

• Draw in all heteroatoms and hydrogens directly bonded to them.


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Condensed StructuresSkeletal structures

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• A charge on a carbon atom takes the place of one hydrogen atom.

• Negatively charged carbon atoms have one lone pair and positively charged carbon atoms have none.

Condensed StructuresSkeletal structures

Cannot be bonded to 3 H atoms, otherwise octet will be lost

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When the 1s orbital of one H atom overlaps with the 1s orbital of another H atom, a sigma () bond that concentrates electron density between the two nuclei is formed.

This bond is cylindrically symmetrical because the electrons forming the bond are distributed symmetrically about an imaginary line connecting the two nuclei.

Hydrogen

Orbitals and Bonding

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Carbon has four valence electrons; two in the 2s orbital and one each in 2p orbitals.

Carbon forms four identical bonds in methane

Methane


Methane (CH4)

Ground state

Excitedstate

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The mixing of a spherical 2s orbital and three dumbbell shaped 2p orbitals together produces four hybrid orbitals, each having one large lobe and one small lobe.

The four hybrid orbitals are oriented towards the corners of a tetrahedron, and form four equivalent bonds.

sp3 Hybrid Orbitals


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Bonding Using sp3 Hybrid Orbitals

Each bond in CH4 is formed by overlap of an sp3 hybrid orbital of carbon with a 1s orbital of hydrogen. These four bonds point to the corners of a tetrahedron.

Orbitals and Bondingsp3 Hybrid Orbitals

4+

H (1s)C (sp3)

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• One 2s orbital and two 2p orbitals form three sp2 hybrid orbitals.

• One 2s orbital and one 2p orbital form two sp hybrid orbitals.

sp2 & sp Hybrid Orbitals


sp3

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Determining Hybridization Patterns

To determine the hybridization of an atom in a molecule, we count the number of groups around the atom. The number of groups (atoms and nonbonded electron pairs) corresponds to the number of atomic orbitals that must be hybridized to form the hybrid orbitals.


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Hybrid orbitals of NH3 and H2O


2s22p5

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The bonds between B and F are the result of sp2 hybrid orbital overlapThis leaves the B atom and each of the F atoms with a p orbital not used in the bond formation. The filled F orbital and unfilled B orbital give rise to localized interactions

Orbitals and BondingHybridization in BF3

B: sp2 + one empty 2p orbital

F: sp2 + one filled 2p orbital

Dotted lines refers to extra electron density

Total must have 1s and 3p orbitals altogether.

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sp2

filled 2p empty 2p

B F

F

F

F B+3

+

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B F

F

F


The presence of the resonance contributors in the bottom line account for the partial double bond character in BF3

Resonance Structures for BF3



CM1501 Lecture 1 – Structure & Bonding

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Orbitals and BondingResonance Structures for BF3

CM1501 Lecture 1 – Structure & Bonding

Although F is more electronegative than B, there is still a positive charge on F because fluorine has to share some of its electron density with boron so that the whole molecule will be stable.

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Each carbon is tetrahedral. Each carbon is sp3 hybridized

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Each carbon is trigonal and planar. Each carbon is sp2 hybridized


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Each carbon is trigonal and planar. Each carbon is sp2 hybridized


Unlike the C—C bond in ethane, rotation about the C—C double bond in ethylene is restricted as it involves the breaking of bond .

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Each carbon is diagonal and linear.

Each carbon is sphybridized

Two σ bonds and two bonds


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Each carbon is diagonal and linear.

Each carbon is sphybridized

Two σ bonds and two bonds

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• As the number of electrons between two nuclei increases, bonds become shorter and stronger.


Bond Length and Bond Strength

Stronger bond does not mean better stablility

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• The length and strength of C—H bonds vary depending on the hybridization of the carbon atom.



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Although sp3, sp2 and sp hybrid orbitals are similar in shape, they are different in size

s is smaller than p. Degree of overlapping with s is larger compared to p, hence larger percent of s-character, the stronger the bond.

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Electronegativity is a measure of an atom’s attraction for electrons in a bond.

Electronegativity values for some common elements

Electronegativity and Bond Polarity

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Periodic Electronegativity Values


Calculated values, not determined by experiments

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Electronegativity values are used as a guideline to indicate whether the electrons in a bond are equally shared or unequally shared between two atoms. When electrons are equally shared, the bond is nonpolar. When differences in electronegativity result in unequal sharing of electrons, the bond is polar, and is said to have a “separation of charge” or a “dipole”.


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Bonding between atoms of different electronegativity values results in unequal sharing of electrons.

Example: In the C-O bond, the electrons are pulled away from C (2.5) toward O (3.4), the element of higher electronegativity. The bond is polar, or polar covalent. The bond is said to have dipole; that is, separation of charge.

The direction of polarity in a bond is indicated by an arrow with the head of the arrow pointing towards the more electronegative element. The tail of the arrow, with a perpendicular line drawn through it, is drawn at the less electronegative element.

+ means the indicated atom is electron deficient.

- means the indicated atom is electron rich.


E = 2.5 E = 3.4

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Use the following two-step procedure to determine if a molecule has a net dipole:

1. Use electronegativity differences to identify all of the polar bonds and the directions of the bond dipoles.

2. Determine the geometry around individual atoms by counting groups, and decide if individual dipoles cancel or reinforce each other in space.

Electrostatic potential plot of CH3Cl

Polarity of Molecules

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A polar molecule has either one polar bond, or two or more bond dipoles that reinforce each other. An example is water:

A nonpolar molecule has either no polar bonds, or two or more bond dipoles that cancel. An example is carbon dioxide:

Polarity of Molecules

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The Periodic Table

1s2

2s22p6

3d104s24p6

4d105s25p6

Rows

4f145d106s26p6

2

8

8

18

18

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

(1)(2+1)

(2+8+1)

(2+8+8+1)

Column

3s23p6

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1s1

2s1

3s1

4s1

2s1 2s2 2s22p2 2s22p4 2s22p6

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