CHAPTER 2 CHEMICAL BONDSrhbestsh.x-y.net/.../Chemistry/Lecture_Note/Ch.02.pdf · 2010. 4. 12. · Ionic bond : Na+Cl– Covalent bond : H 2 Metallic bond : Fig. 2.1 Sodium Chloride

2009년도 제1학기 화 학 1 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 2

CHAPTER 2 CHEMICAL BONDS

Ionic bond : Na+Cl– Covalent bond : H2Metallic bond :

Fig. 2.1 Sodium Chloride (NaCl) Fig. 1.53 Metallic bond

IONIC BONDS Ionic solids Crystalline solids

2.1 The Ions That Elements Form Ions of s-block metallic elements

Li : [He]2s1 Li+ : [He] duplet (2전자계)

Na : [Ne]3s1 Na+ : [Ne]

cf. Outer electron configuration of Ne : 2s22p6 Octet (8전자계) Ions of p-block metallic elements of periods 2 and 3

Al : [Ne]3s13p1 Al3+ : [Ne]

Ions of p-block metallic elements of period 4 and later periods

Ga : [Ar]3d104s24p1 Ga3+ : [Ar]3d10

Ions of d-block metallic elements

Fe : [Ar]3d64s2 Fe3+ : [Ar] 3d5

Ions of nonmetal elements : Gain electrons to acquire the next noble gas configuration

N3– : [Ne]

Fig. 2.2 Electron configurations of cations Fig. 2.3 Electron configurations of anions of main-group metals become those of the of nonmetal atoms become those of the preceding noble gas atoms. next noble gas.


Ex. 2.1 Writing the electron configuration of cations

In : [Kr]4d105s25p1 ln+ : [Kr]4d105s2 ln3+ : [Kr]4d10

2.2 Lewis Symbols

Gilbert Newton Lewis (美,1875-1946)

Lewis Symbols Dots for valence electrons N : 2 1 12 2 2 2x ys p p p

1z

Cl K MgH He N O : : : :⋅ ⋅ ⋅ ⋅ ⋅ ⋅ii iiiii ii

i

2Cl Ca Cl Cl Ca Cl : + : :: : : :− −

+⎡ ⎤ ⎡ ⎤⎯⎯→ ⎢ ⎥ ⎢ ⎥⎣ ⎦ ⎣ ⎦⋅ + ⋅

ii ii ii ii

ii ii ii ii

2.3 The Energies of Ionic Bond Formation Why NaCl(s) is more stable than Na(g) + Cl(g)?

(1) Ionization energy of Na atoms

+ 1Na(g) Na (g) + e (g) energy required = 494 kJ mol− −⎯⎯→ ⋅ (2) Electron affinity of Cl atoms

1Cl(g) + e (g) Cl (g) energy released = 349 kJ mol− − −⎯⎯→ ⋅ (3) Net change in energy : 494 – 349 kJ·mol–1 = +145 kJ·mol–1

(4) Experimentally, + 1Na (g) + Cl (g) NaCl(s) energy released = 787 kJ mol− −⎯⎯→ ⋅

Due to attraction between oppositely charged ions

(5) Net change in energy for the overall process

Na(g) + Cl(g) NaCl(s)⎯⎯→ is 145 – 787 kJ·mol–1 = – 642 kJ·mol–1 Huge decrease in energy !


Fig. 2.4 NaCl(s) is more stable in energy than Na(g) + Cl(g).

2.4 Interactions Between Ions ◈ Lattice energy of a solid

Difference in energy between the ions packed together in a solid and

the ions widely separated as a gas

Fig. 2.5 The sequence of images illustrates why ionic compounds are brittle.

▷ Ionic solids are brittle due to strong repulsions between like charges of ions.


Coulomb potential energy between two ions: 2

1 2 1 2,12

0 12 0 12

( ) ( )4 4p

z e z e z z eEr rπε π

×= =

ε

Calculation of the potential energy of a one dimensional ionic solid

For the interaction arising from ions to the RHS, total potential energy of the central ion is

2 2 2 2 2 2 2 2

0

14 2 3 4P

z e z e z e z eEd d d dπε

⎛ ⎞= × − + − + − ⋅ ⋅⎜ ⎟

⎝ ⎠⋅

2 2 2 2

0 0

1 1 11 ln 24 2 3 4 4z e z e

d dπε πε⎛ ⎞= − − + − + ⋅⋅ ⋅ = − ×⎜ ⎟⎝ ⎠

Fig. 2.6 1-d ionic lattice

Total potential energy per mole of ions 2 2 2 2

A A

0 0

2 ln 2 1.3864 4Pz N e z N eE

d dπε π= − × = − ×

ε

For 3-d lattices of ions 2

1 2 A

0

| |4P

z z N eE Adπε

= − ×

A : Madelung constant

Repulsive potential between ions

**P

d dE e−∝ / , d* = 34.5 pm, a constant

Total potential energy

*Total P PE E E= +

Energy at the minimum

2 *A 1 2

P,min0

| | 1 4

N z z e dE Ad dπε

⎛ ⎞= − −⎜ ⎟

⎝ ⎠

Born-Meyer equation

▶ Great stabilization Highly charged, Small ions

Fig. 2.7 PE of an ionic solid

♦ Our skeleton ionic solid, calcium phosphate

Doubly charged Ca2+ ions

Triply charged ions 34PO−

rigid, insoluble solid

Fig. 2.8 A micrograph of bone, calcium phosphate.

2009년도 제1학기 화 학 1 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 2 COVALENT BONDS 1916 Lewis discovered “Covalent bonds” a pair of electrons shared between two atoms

2.5 Lewis Structures

Octet Rule:

In covalent bond formation, atoms go as far as possible

toward completing their octet by sharing electron pairs.

Valence (原子價) of an element is the number of bonds that its atom can form.

Lone pairs (孤立雙) of electrons are pairs of valence electrons that do not take part in bonding.

Lewis structure of a molecule shows atoms by their chemical symbols, covalent bonds by lines,

And lone pairs by pairs of dots.

F F F F : : :⎯⎯→ −⋅ + ⋅ii ii ii ii

ii ii ii ii:

⋅

2.6 Lewis Structures for Polyatomic Species

Mathane: 8 valence electrons C H H H H: ⋅ ⋅ ⋅i

i

Arrange dots so that C atom has an octet and each H atom has a duplet

C is tetravalent

Single, double, and triple bonds C::O C=O C:::O C≡O

Bond order : Number of bonds that links a specific pair of atoms

H2(1), C=C (2), C≡C (3)


▶ Rule of thumb for predicting the Lewis structure of compounds:

(1) Choose as the central atom the element with the lowest ionization energy.

(2) Arrange the atoms symmetrically around the central atom.

H2SO4 (NH4)2CO3

Ex. 2.4 Write the Lewis structure for acetic acid.

1) Count the number of all valence electrons C(4)H3(3)C(4)O(6)O(6)H(1) 24 valence electrons

2) Arrange atoms methyl group, CH3 – and carboxyl group, –COOH

H O

░ ░

H ░ C ░ C

░ ░

H O ░ H

3) Connect the atoms with bonding electron pairs.

H O.. ..

H : C : C.. ..H O : H

4) Complete the octets with remaining electrons.

H : O :.... ..

H : C : C.. ..H : O : H

..

5) Represent the bonds with single or double, or triple lines leaving the lone pairs. H : O : I II H – C – C I I H : O – H ..

2009년도 제1학기 화 학 1 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 2 2.7 Resonance (共鳴)

◆ Resonance hybrid (共鳴混成) nitrate ion, NO3–

All bond lengths in a nitrate ion are the same : 124 pm

cf. N=O bond length (120 pm), N–O bond length (140 pm)

Electrons are delocalized (非偏在化) over N and three O atoms

[ resonance structure 1 resonance structure 2 ↔ ↔ ······ ↔ resonance structure n ]

◆ Resonance hybrid of acetate ion

◆ Resonance hybrid of benzene, C6H6

August Kekulé (獨,1829-1896) ▶ Reactivity : Benzene does not react with bromine.

1-hexene + Br2 1,2-dibromohexene ⎯⎯→

Fig. 2.9 Addition of bromine to 1-hexene

▶ Bond length: C–C bonds in benzene are all the same length 139 pm

cf. C–C (154 pm), C=C (134 pm)

▶ Structural evidence : Only one dichlorobenzene exists.


2.8 Formal Charge (形式電荷)

Formal charge on an atom in a given Lewis structure is the charge it would have

if the bonding were perfectly covalent. exactly half-sharing of bonding electrons

12Formal charge = ( )V L B− +

V : number of valence electrons in the free atom L : number of lone pairs on the bonded atom

B : number of bonding electrons on the atom Used for prediction of most favorable arrangement of atoms in a molecule

Lewis structure with lowest formal charges are likely to have the lowest energy

Ex. CO2 N2O

EXCEPTIONS TO THE OCTET RULE 2.9 Radicals and Biradicals Radicals (free radicals) :

Species having electrons with unpaired spins, • CH3

Very reactive

Antioxidants (Vitamin C, E) delay the damage (human aging) by radicals NO•

Involved in atmospheric chemical reactions (decomposition of ozones)

NO• neurotransmitter, supply of blood to organs

Biradicals : two unpaired electrons, O⋅ ⋅ii

ii ,


Box 2.1 WHAT HAS THIS TO DO WITH …. STAYING ALIVE? Human aging Radicals

Oxidation of lipids in cell membranes

DNA, RNA damaged by radicals

Antioxidants: Vitamin A,C,E, Antioxidant enzymes, Coenzyme Q (조효소 Q)

Fish oils, Wheat grass (갯보리), Ginkgo biloba leaves (銀杏잎), Green vegetables, Orange juice,

Chocolate, Coffee, Tea

Oxidative stress : from suntan, smoking

2.10 Expanded Valence Shells

Large atoms with empty d-orbitals in the valence shell

Hypervalent compound : Octet rule expanded

Element showing variable covalence: P

4 2P (s) + 6 Cl (g) 4 PCl (l)⎯⎯→ 3

Fig. 2.10 PCl3(l)

3 2PCl (l) + Cl (g) PCl (s)⎯⎯→ 5

☺ Phosphorus pentachloride is an ionic solid consisting of 4PCl+ and 6PCl

−

At 160oC it vaporizes to a gas of 5PCl

6PCl− : expanded valence shell of P to 12 electrons, making use of two 3d-orbitals

5PCl : expanded valence shell of P to 10 electrons, making use of one 3d-orbital

2009년도 제1학기 화 학 1 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 2 Ex. 2.7 Writing a Lewis structure with an expanded valence shell : sulfur tetrafluoride, SF4

(1) Count the number of valence electrons

6 from sulfur S⋅ ⋅ii

ii

7 from each fluorine atom F: ⋅iiii

F S F

F F

: : / \

: :

− −ii ii ii

ii ii

ii ii

ii ii

(2) Find the number of electron pairs

6 + (4 x 7) = 34 electrons 17 electron pairs

(3) Construct the Lewis structure

Give each F atom 3 lone pairs and 1 bonding pair shared with the central S atom.

Place 2 extra electrons as a lone pair on the S atom

Ex. 2.8 Selecting the dominant resonance structure for a molecule

Determine the dominant resonant structure of a sulfate ion by calculating

the formal charges on the atoms in each structure in (26)

Step 1 Count the valence electrons (V) (O: 6) x 4 + (S: 6) x 1

Total 30 electrons, which provide 15 pairs of electrons

Step 2 Draw the Lewis structures

2

O |O-- S -- O

| O

: :

: :

: :

−ii

ii ii

ii ii

ii

2

O |O== S -- O | O

: :

:

: :

−ii

ii ii

ii ii

ii

2

O |O== S == O | O

: :

: :

−ii

ii ii

ii ii

ii

Step 3 Assign electron ownership, 12( )L B+

27

4

O |O-- S -- O

| O

: :

: :

: :

−ii

ii ii

ii ii

ii

27

65

O |

O == S -- O | O

: :

:

: :

−ii

ii ii

ii ii

ii

27

66

O |

O == S == O | O

: :

: :

−ii

ii ii

ii ii

ii

Step 4 Find the formal charge, 12( )V L B− +

21

1+2

1

1

O |

O -- S -- O | O

: :

: :

: :

−−

−

−

−

ii

ii ii

ii ii

ii

21

0+1

1

1

O |

O == S -- O | O

: :

:

: :

−−

−

−

ii

ii ii

ii ii

ii

21

00

0

1

O |

O == S == O | O

: :

: :

−−

−

ii

ii ii

ii ii

ii

2009년도 제1학기 화 학 1 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 2 2.11 The Unusual Structures of Some Group 13/III Compounds

◆ Incomplete octet, BF3 Resonant hybrid of (29) and (30) (29) major contribution

↔

◆ Coordinate covalent bond both electrons come from one of the atoms

2

F-- B -- F | F

F |

: :ii

: :

: :

−

ii ii

ii ii

ii

31 Tetrafluoroborate, 4BF−

IONIC VERSUS COVALENT BONDS 2.12 Correcting the Covalent Model: Electronegativity

Cl2

+

: Cl Cl: : Cl : Cl : : Cl : Cl :

⋅ ⋅⋅ ⋅ ⋅ ⋅ ⋅ ⋅ ⋅ ⋅ ⋅ ⋅− + −

⋅ ⋅ ⋅ ⋅ ⋅ ⋅⋅ ⋅ ⋅ ⋅ ⋅ ⋅− ↔ ↔

Small contribution from ionic structures

Two ionic structures equally but oppositely contribute (average charge on each atom is 0)

Cl2 almost purely covalent

HCl

+

H Cl: H : Cl : H : Cl :

⋅ ⋅ ⋅ ⋅ ⋅ ⋅− + −

⋅ ⋅ ⋅ ⋅ ⋅ ⋅− ↔ ↔

Unequal contributions from ionic structures

has lower energy contribute more !

+

H : Cl :

⋅ ⋅−

⋅ ⋅

Partial charges on each atoms: + H Clδ δ −− cf. zero formal charge for both H and Cl H and Cl form an electric dipole (a partially positive charge next to a partially negative charge)

Dipole is represented by an arrow pointing toward the negative charge

Polar covalent bond ◈ Electric dipole moment, μ = q r, in units of debye (D) Size of an electric dipole

Measure of the magnitude of the partial charges

Peter Debye (和,1884-1966) Nobel Prize ’36 Chemistry


★ Definition of the dipole moment:

A single negative charge separated by 100 pm from a single positive charge

has a dipole moment of 4.8 D

☆ Dipole moment of the Cl – H bond is 1.1 D.

Arising from a partial charge of about 23% of an electron’s charge on the Cl atom and an

Equivalent positive charge on the H atom ◈ Electronegativity, χ

1932 Linus Pauling proposed a quantitative measure of electron distribution in bonds.

Electron-pulling power of an atom when it is part of a molecule

Fig. 2.11 More electronegative atom (B) pulls shared electrons.

◆ Pauling’s electronegativity scale:

[ ]2 1A B 2(A B) (A A) (B B)D D Dχ χ− = − − − + − , D: dissociation energy

◆ Robert Mullikan’s electronegativity scale:

(1 a2 )I Eχ = + I : Ionization energy Ea : Electron affinity

Fig. 2.12 Variation in the electronegativity Fig. 2.13 Dependence of the percentage

of the main-group elements. ionic character of the bond on χ∆


Linus Pauling (美,1901-1994) Robert Mullikan (美,1896-1986)

Nobel Prize ’54 Chemistry Nobel Prize ’66 Chemistry

Nobel Prize ’62 Peace "for his fundamental work concerning chemical

"for his research into the nature of the chemical bonds and the electronic structure of molecules

bond and its application to the elucidation of the by the molecular orbital method"

structure of complex substances"

“The Nature of the Chemical Bond,” Cornell University Press (1939). 2.13 Correcting the Ionic Model: Polarizability

Covalent character of ionic bond

Distortion of electron cloud of anion toward the cation

Sharing of electrons of anion with cation

◈ Polarization

Highly polarizable anions ––– large anions, I− Highly polarizing cations ––– small, highly charged, Al3+

Polarizing power of cation increases across a period,

decreases down a group diagonal relationship

Li+, Mg2+ similar power

Fig. 2.14 Polarization of

electron cloud

THE STRENGTHS AND LENGTHS OF COVALENT BONDS 2.14 Bond Strengths ◈ Dissociation energy, D :

Energy required to separate the bonded atoms, Measure of strength of a chemical bond

Fig. 2.15 Dissociation energy

2009년도 제1학기 화 학 1 담당교수: 신국조 Textbook: P. Atkins / L. Jones, Chemical Principles, 4th ed., Freeman (2008) Chapter 2 2.15 Variation in Bond Strength

◈ Average bond dissociation energy

● Bond strength vs. Lewis structure

Bond strengths: N2 > O2 > F2Bond oders: N2(3), O2(2), F2(1)

Fig. 2.16 Bond dissociation energies

Fig. 2.17 Strengths of bonds between two C atoms

● Stabilization due to resonance hybridization

C≈C bond in benzene D = 518 kJㆍmol–1

C—C bond in alkanes D = 348 kJㆍmol–1

C==C bond in alkenes D = 612 kJㆍmol–1

● Influence of lone pairs of electrons : D(H2) > D(F2)

Repulsion between lone pairs weaken the bond


● Trend in bond strength vs. trend in atomic radii

Fig. 2.18 D(HX) Fig. 2.19 D(HY) Y: p-block element

Increase in atomic radii Decrease in bond strength

Hydrides: D(methane,CH4) > D(silane,SiH4) > D(stannane,SnH4) > D(plumbane,PH4)

● Adenosine triphosphate (ATP, 35)

ATP ADP releases energy !

Breaking P-O bond in ATP 276 kJㆍmol–1 required

New P-O bond formation in 2 4H PO− releases 350 kJㆍmol–1

Net energy released : 74 kJㆍmol–1


2.16 Bond Lengths

◈ Bond length

Distance between the centers of two atoms joined by a covalent bond

Measured by spectroscopy or X-ray diffraction

Bonds between heavy atoms are longer than those between light atoms

Multiple bonds are shorter than single bonds between the same two elements

The stronger the bond, the shorter the bond

Approximately the sum of the covalent radii of the two atoms

Fig. 2.21 Covalent radii of H and p-block elements


BOX 2.2 HOW DO WE KNOW……THE LENGTH OF A CHEMICAL BOND

Measurement of the bond length : X-ray (solids), Microwave spectroscopy (gases) Diatomic molecule, AB

Rotation energy 2

2 2

( 1) , 0,1,2,...8

h J JE JRπ µ+

= =

J : rotational quantum number

μ : reduced mass

A B

A B

m mm m

µ =+

Minimum energy needed to excite a molecule

into rotation from rest: 2

2 24hE

Rπ µ∆ =

Heavy molecules are easy to excite!

Energy of incident microwave radiation 2

2 24hE h

Rν

π µ∆ = =

Rotational energy levels (a) a heavy diatomic molecule (b) a light diatomic molecule

MAJOR TECHNIQUE 1 INFRARED SPECTROSCOPY ◆ Infrared radiation

~ 1000 nm or 3 x 1014 Hz

typical frequency of molecular vibration

excites the vibrational degree of freedom of a molecule

◆ Vibration

Hooke’s law : Force = – k x displacement

k : force constant measure of stiffness of a bond

▷ Strength of a bond measure of the depth of the potential well

▷ Stiffness of a bond

determined by the steepness of the potential curve

◆ Vibrational frequency : 1

2kν

π µ=

High frequency for stiff bonds and low atomic masses


Vibrational energy : 12( ) ( 0, 1, 2, ...)vibE n h nν= + =

vibE hν∆ = independent of vibrational quantum number, n

◈ Normal modes of vibration (基準振動方式) of polyatomic molecules with N atoms

3N – 6 vibrational degrees of freedom (nonlinear molecules)

3N – 5 vibrational degrees of freedom (linear molecules)

▶ Water, H2O 3 normal modes vibrations

▶ Carbon dioxide, CO2 4 normal modes vibrations (two degenerate bending modes)

◆ IR spectrometer

IR absorption spectrum

Fingerprint region

Fig. 3. IR spectrum of an amino acid

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CHAPTER 2 CHEMICAL BONDSrhbestsh.x-y.net/.../Chemistry/Lecture_Note/Ch.02.pdf · 2010. 4. 12. · Ionic bond : Na+Cl– Covalent bond : H 2 Metallic bond : Fig. 2.1 Sodium Chloride