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Today’s Quiz

1. What is ground-state electron configuration?2. Define valence electrons and valence shell.3. Explain the exceptions to the octet rule.4. Define Electronegativity.5. What is the difference between a nonpolar covalent bond and a polar

covalent bond.6. What are the 3 molecular shapes?7. What is a hybrid orbital?

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Chapter 1: Covalent Bonding and Shapes of Molecules

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1.1 How do we describe the electronic structure of Atoms?

1.2 What is the Lewis Model of Bonding?

1.3 How do we Predict Bond Angles and the Shapes of Molecules?

1.4 How do we predict if a molecule is polar or non-polar?

1.5 What is resonance?

1.6 What is the Molecular orbital model of covalent bonding?

1.7 What are the Functional Groups? (Group Presentation)

Chapter 1: Covalent Bonding and Shapes of Molecules

Quantum Numbers and Atomic OrbitalsAn atomic orbital is specified by four quantum numbers.

n the principal quantum number - a positive integer, indicates the relative size of the orbital or the distance from the nucleus

l the angular momentum quantum number - an integer from 0 to n-1, related to the shape of the orbital

ml the magnetic moment quantum number - an integer from -l to +l,

orientation of the orbital in the space around the nucleus 2l + 1

Schrodinger Wave Equation

ms spin quantum number

Table 7.2 The Hierarchy of Quantum Numbers for Atomic Orbitals

Name, Symbol(Property) Allowed Values Quantum Numbers

Principal, n(size, energy)

Angular momentum, l

(shape)

Magnetic, ml

(orientation)

Positive integer(1, 2, 3, ...)

0 to n-1

-l,…,0,…,+l

1

0

0

2

0 1

0

3

0 1 2

0

0-1 +1 -1 0 +1

0 +1 +2-1-2

Y = fn(n, l, ml, ms)

spin quantum number ms

ms = +½ or -½

Schrodinger Wave Equation

ms = -½ms = +½

7.6

Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.

1s1

principal quantumnumber n

angular momentumquantum number l

number of electronsin the orbital or subshell

Orbital diagram

H

1s1

7.8

Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s7.7

Ionic Bond

• Is the electrostatic force that holds ions together in an ionic compound

• The metal gives the electrons to the non metal

• Cation is the metal• Anion is the non metal

The Covalent Bond

• A bond in which two electrons are shared by two atoms.

• Only in covalent compounds

• Non metal and non metal

The Covalent Bond

The Covalent Bond

• Pairs of valence electrons that are not involved in covalent bond formation are called: Lone pairs

• We can only draw Lewis structures for compounds that have covalent bonds

Lewis Dot Symbols

Consists of the symbol of an element and one dot for each valence electron in an atom of the element.Note that (except helium) the number of valance electrons each atom has is the same as the group number of the element.

Lewis Structures

• Is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs of dots on individual atoms

• Only valence electrons are shown in a Lewis structure

• Valence electrons: Electrons in the valence (outermost) shell of an atom.

• Valence shell: The outermost electron shell of an atom.

Electronegativity

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

Electronegativity

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I. Lewis StructuresB. Ionic, covalent, and polar bonds

H—HC = 2.1 2.1

DC = 0 equal sharing of electronsCl—Cl = nonpolar covalent bond

C = 3.0 3.0

H—Cl DC = 0.9 unequal sharing of electronsC = 2.1 3.0 = polar covalent bond

Na+Cl– DC = 2.1 transfer of electronsC = 0.9 3.0 = ionic bond

generally: when DC <0.5 non-polar covalent 0.5<DC = 1.9 polar covalent

> 1.9 ionic

d+ d–

nonmetal+

nonmetal

metal+

nonmetal

Polar and Nonpolar Molecules

• A molecule will be polar if:

• 1. It has polar bonds.

• 2. The center of partial positive charge lies at a different place within the molecule than the center of partial negative charge.

The "best" Lewis structure for NO3-

• 1. Determine the total number of valence electrons in a molecule

• Draw a skeleton

• Of the 24 valence electrons in NO3-, 6 were required

to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible

• Are the octets of all the atoms filled? If not then fill the remaining octets by making multiple bonds

• Check that you have the lowest FORMAL CHARGES possible for all the atoms, without violating the octet rule; (valence e-) - (1/2 bonding e-) - (lone electrons).

The Octet Rule

Exceptions to the Octet Rule

• The expanded octet– Atoms in the 2nd period (that

are in families 3A-7A) cannot have more than 8 valance electrons around the central atom

– Atoms of elements in and beyond the 3rd period (that are in families 3A-7A) form some compound in which more than 8 electron surround the central atom

• Odd- electron molecules– Some molecules contain an

odd number of electrons.– We need an even number

of electron for complete pairing the octet rule clearly cannot be satisfied with all the atoms in any of these molecules

• Incomplete octet– The number of electrons

surrounding the central atom in a stable molecule is fewer than eight

Figure 2.10 The modern periodic table.

+1+2 +3 -3 -2 -1

0

NC

The Concept of Resonance

• Resonance means the use of two or more Lewis structures to represent a particular molecule

• Resonance structure, is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure

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E. Resonance structures-two or more equivalent Lewis structures-nuclei remain in fixed positions, but electrons arranged differently

HCO2-

H CO

OH C

O

O= H C

O

O

½ -

½ -

1.5 bond order•neither of these accurately describes the formate ion•actual species is an average of the two (resonance hybrid)

•Resonance hybrid: A molecule or ion that is best described as a composite of a number of contributing structures.

delocalizedelectrons

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I. Lewis StructuresE. Resonance structures

=CO

NH2

CH3 CNH2

OCH3 CH3 C

O

NH2+

-

more stablemajor contributor

less stableminor contributor

x

y

z

90o

90o

2pz

2py2px

Shapes of 2px, 2py, 2pz atomic orbitals and their orientation in space

methane

H

H

HH

109.5o

ethene

C C

H H

HH

120o

ethyne

C CH H

180o

tetrahedralTrigonal planar or pyramidal

linear

Region of electron density around the central atom.

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Sigma bond is covalent bond in which the overlap of atomic orbitals is concentrated along the bond

axis

Hybrid orbital an orbital produced from the combination of two or more atomic orbitals.

Pi bond a covalent bond formed by the overlap of parallel p orbitals.

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Why does hybridization occur?

• In chemistry hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties.

• Hybridized orbitals are very useful in the explanation of the shape of molecular orbitals for molecules. It is an integral part of organic chemistry.

• Sp3 hybrid orbitals – Bond angles of approximately 109.5o

• Four sigma bonds

• 4 groups bonded to carbon

Sp2 hybrid orbitals – Bond angles of approximately 120o

• Three sigma bonds and one pi bond

• 3 groups bonded to carbon

• Sp hybrid orbitals – Bond angles of approximately 180o • Two sigma bonds and two pi bonds

• 2 groups bonded to carbon

tetrahedral

Trigonal planar

linear

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III. Valence Bond ModelA. Hybrid atomic orbitals

1. sp3 hybridization

CH4 facts: tetrahedral,4 equivalent bondsC

H

HH

H

C2s

2ppromoteelectron 2p

2s

hybridize

sp3

hybrida.o.s

sp3 hybrid a.o.s:

C(sp3)tetrahedral (sp3

C + 1sH)

4HC

H

HH

H

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III. Valence Bond ModelB. Hybrid atomic orbitals

1. sp3 hybridization

C N O

CN

HH

H

HH

CC

H

HH

H

HH C

OH

H

HH

(sp3C + sp3

C) (sp3C + sp3

N) (sp3C + sp3

O)

lone pairsin sp3 a.o.s

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III. Valence Bond ModelB. Hybrid atomic orbitals

2. sp2 hybridizationC2H4 facts: all six atoms lie

in same planeC CH

H

H

H

trigonal planar = sp2

2psp2

C Csp2

2p

H1s 1s

H H1s 1s

H

C CH H

H H

2p

(sp2C + 1sH)

(sp2C + sp2

C)

overlapp orbitals C C

H H

H H

bond

all atoms coplanarfor p orbital overlap

= C C

H

H

H

H

double bond =1 bond +1 bond

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III. Valence Bond ModelB. Hybrid atomic orbitals

2. sp2 hybridization

C O

(sp2C + sp2

C) + (sp2C + sp2

O) +

lone pairsin sp2 a.o.s

C CHH

HHC O

H

H

C C

H

H

H

H

C

H

H

O

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III. Valence Bond ModelB. Hybrid atomic orbitals

3. sp hybridization

C2H2 facts: linear = spH C C H

2psp

C Csp

2p

H1s 1s

H

C C HH

2p

(spC + 1sH)

(spC + spC)

C CH H

2 bonds

= C CH H

triple bond =2 bonds +1 bond

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III. Valence Bond ModelB. Hybrid atomic orbitals

3. sp hybridization

N

(spC + spN) + 2

lone pairin sp a.o.

C NH

H C N