2012 general chemistry i chapter 2. chemical bonds 2012 general chemistry i ionic bonds covalent...

2012 General Chemistry I

Chapter 2.Chapter 2.CHEMICAL BONDSCHEMICAL BONDS


IONIC BONDS

COVALENT BONDS

2.1 The Ions That Elements Form2.2 Lewis Symbols2.3 The Energetics of Ionic Bond Formation2.4 Interactions Between Ions

2.5 Lewis Structures2.6 Lewis Structures of Polyatomic Species2.7 Resonance2.8 Formal Charge

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Key Ideas and Definitions

Chemical bond is the link between atoms.

- ionic bond i.e. Na+, Cl-

- covalent bond i.e. NH3

- metallic bond i.e. Cu

Bond formation is accompanied by a lowering of energy: it isa favorable process. Formation of chemical bonds results in ionic lattices, molecules,and metallic lattices.It is fundamental to the production of compounds and so is central to the science of chemistry.Energy lowering is due to attractions between oppositely charged ions (ionic bonding), or between nuclei and shared electron pairs (covalent bonding).

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IONIC BONDS (Sections 2.1-2.4)

Ionic model: the description of bonding in terms of ions

An ionic solid is an assembly of cations and anions stacked together in a regular pattern, called a crystal lattice.

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2.1 The Ions That Elements Form2.1 The Ions That Elements Form

Cations are formed by removal of outermost electrons in the order np, ns, (n-1)d.

Main-group metal atoms lose their valence s- and/or p-electrons and acquire the electron configuration of the preceding noble gas atom.


Anions: Add electrons until the next noble-gas configuration is reached

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2012 General Chemistry I 5

Self-Tests 2.1B and 2.2B

2.1B Write the electron configurations of (a) the Manganese (II) ion and (b) the lead (IV) ion.

Solution: (a) Mn is [Ar] 3d5 4s2, hence Mn2+ is [Ar] 3d5

(b) Pb is [Xe] 5d10 4f14 6s2 6p2, hence Pb4+ is [Xe] 5d10 4f14

2.2B Write the chemical formula and electron configuration of the iodide ion.

Solution: I needs to gain only one electron to achieve Xe electron structure, hence I- and [Xe] or [Kr] 4d10 4s2 4p6


2.2 Lewis Symbols2.2 Lewis Symbols

Valence electrons – depicted as dots; a pair of dots for paired electrons.

- Cations and anions

6

Atoms and ions are conveniently represented by Lewis dot symbols, showing the element symbol, the valence electrons and charge, if any.

- Atoms

Here, we can use Lewis dot symbols to show electron transfer in the formation of cations and anions.


Gilbert N Lewis


Na(g) → Na+(g) + e- (g)

Cl(g) + e- (g) → Cl-(g)

Na+(g) + Cl- (g) → NaCl(s)

494 kJ·mol-1

-349 kJ·mol-1

-787 kJ·mol-1

Na(g) + Cl(g) → NaCl(s) -642 kJ·mol-1

2.3 The Energetics of Ionic Bond Formation2.3 The Energetics of Ionic Bond Formation

The formation of a typical ionic compound, such as sodium chloride can be broken down into three simple steps:1. Ionization of gaseous metallic atom to give the cation.2. Formation of gaseous non-metallic anion.3. Condensation of the gaseous ions into a crystal lattice.

The energies involved in these processes illustrate the favorability of ionic compound formation (see Fig. 2.4, next slide):


Energetics of ionic compound formation.The difference in energybetween the ions in thelattice and separated gaseous ions is calledthe lattice energy.


2.4 Interactions Between Ions2.4 Interactions Between Ions

- In an ionic solid, each cation is attracted to all the anions to a greater or lesser extent. This is a “global” characteristic of the entire crystal

Lattice energy: the difference in energy between the ions packed together in a solid and the ions widely separated as a gas

- Strong electrostatic interactions in ionic solids → high melting points and brittleness

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- Coulomb potential energy of the interactions of two individual ions:

Here, e is the fundamental charge; z1 and z2 are the charge numbers of thetwo ions; r12 is the distance between the centers of the ions; 0 is the vacuum permittivity constant.

- Molar potential energy of a three-dimensional crystal:

The factor A is the Madelung constant, dependent on how the ions are arranged about one another in the 3-dimensional lattice.

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d = distance between centers of neighboring ions (rcation + ranion)NA is the Avogadro number


- In a one-dimensional crystal in which cations and anions alternate along a line:

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One-Dimensional Crystal Model


This implies Ep is most favorable for small ions with large charges.Also for a one dimensional model crystal, the Madelung constant A is 2 ln2 = 1.386. This compares well with values of A for real ioniccrystals (Table 2.2).For multi-charged ions (z1 and z2), z2 is replaced by the absolutevalue of z1z2 (its value without the negative sign).


Self-Test 2.3A

The ionic solids KCl and CaO crystallize to form structures of the same type. In which compound are the interactions between the ions stronger?

Solution: The ions in CaO are both smaller and more highly charged, hence CaO has the stronger interactions.


Attraction, Repulsion and the Born-Mayer Equation

15

The previous discussion does not take into account cation-anion repulsion – the real potential energy of an ionic solid is a balance between attractive and repulsive interionic interactions.If a cation and anion are brought together (see Fig. 2.7, opposite), potential energy decreases to a minimum value. Further decrease in d, leads to serious unfavorable repulsive interaction.


Born-Mayer equation

with d* = 34.5 pm

repulsive effect

This equation allows for repulsive interactions at small values of d: EP,min increases (becomes less favorable) when d approaches d* and is actually positive when d* > d.


COVALENT BONDS (Sections 2.5-2.8)

2.5 Lewis Structures2.5 Lewis Structures

Covalent bond - a pair of electrons shared between two atoms

- Octet rule: in covalent bond formation, atoms go as far as possible toward completing their octets by sharing electron pairs.

- Valence of an atom is the number of bonds it can form.

- A line (-) represents a shared pair of electrons.

- Lone pairs of electrons – electron pairs not involved in bonding.

- Lewis structure – atoms are indicated by chemical symbols, covalent bonds by lines, and lone pairs by pairs of dots.

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Some Definitions According to Lewis Theory


Self-Test 2.4A

Write the Lewis structure for the “interhalogen” compound chlorine monofluoride, ClF, and state how many lone pairs each atom possesses in the compound.

Solution

F Cl: :.. .... ..

. .+ F: .... .Cl:..

... or F:..

..Cl:....

Three lone pairs onboth atoms


2.6 Lewis Structures of Polyatomic Species2.6 Lewis Structures of Polyatomic Species

– A Lewis structure does not portray the 3D shape of a molecule or ion, but simply displays which atoms are bonded together.

-

- bond order: the number of bonds that link a specific pair of atoms.

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Writing a Lewis Structure 1. Count the total number of valence electrons, from the group

numbersE.g. CO2 4 + 2 x 6 = 16 (C in group IVA; O in group VIA)

NOTE: if an anion (-), add the value of the charge; if a cation (+), subtract that value.

2. Calculate the total number of electrons that are needed if each atom had its own noble gas shell of electrons (2 for H and 8 for all others).

E.g. for CO2 there are 3 atoms (no H) and hence 24 noble gas shell

electrons.3. Subtract the number in 1 from the number in 2: this gives the

number of shared (bonding) electrons present (and number of bonds = 1/2 that number).

4. Add electron pairs to “complete the octets”, as necessary.

5. Represent each bond by a line.


E.g. for CO2, the figure is 24 – 16 = 8 (4

bonds: two “double bonds”)Hence the Lewis diagram is

CO O

CO O

or

: :

: :

::::....

.. ..

The above works well only for molecules that obey the octet rule.

For certain molecules (like BF3, radicals, and high valence

compounds like SF6), this rule is ignored.


Writing a Lewis Structure: Rules of Thumb

- Terminal atom: bonded to only one other atom- Central atom: bonded to at least two other atoms

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– Usually, element with lowest I1 or lowest electronegativityis central atom. E.g., in HCN, carbon has lowest I1 and is leastelectronegative: hence it is the central atom.

– Usually, there is symmetrical arrangement around central atom. E.g., in SO2, OSO is symmetrical, with S central and O terminal.

– Oxoacids have H atoms mostly bonded to O atoms. E.g., H2SO4

is actually (HO)2SO2, with two O atoms and two OH groups bonded to S.

– For organic compounds, the atoms are arranged into groups, as suggested by the standard molecular formula. E.g., CH3COOH = one CH3 group and one COOH group.


Examples

PHO OH

OH

..

Phosphorous acid,P(OH)3

Acetone, (CH3)2CO

C

O

H3C CH3

..:

N

H

H H

H

N

H

H H

H

+ +C

O

O O.. ..

..

.... ::

: 2-

Ammonium carbonate, (NH4)2CO3

:..

..

......


Self-Test 2.5AWrite a Lewis structure for the cyanate ion, NCO-

(sometime written CNO-).

Solution

N C O.. ..

: :

_or N C O

.. ..:

_

..


Self-Test 2.6A

Write a Lewis structure for the urea molecule, (NH2)2CO

Solution

C

O

N N

H

H

H

H.. ..

..:

from skeleton C O

N

N

H

H

H

H


2.7 Resonance2.7 Resonance

– Multiple Lewis structures: many compounds can be represented bydifferent Lewis structures in which the location of electrons (but not nuclei) differ. They are known as resonance structures, each making a contributionto the real structure of the molecule (called a “resonance hybrid”).

– Resonance implies delocalization: in which a shared electron pair is distributed over several pairs of atoms and cannot be identified withjust one pair of atoms.

– The resonance symbol is a double-headed arrow (↔), indicating a blend of the contributing structures:

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Benzene, C6H6

- No reactions typical of compounds with double bonds

- All the carbon-carbon bonds with the same length

- Only one 1,2-dichlorobenzene exists.

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Self-Test 2.7B

Write Lewis structures contributing to the resonance hybrid for the nitrite ion, NO2

O N O..:.. ..

:.._

ONO .. :....

:.. _

O N O_

The resonance hybrid is

Solution

_


Additional Self-Test

Write Lewis (resonance structures for the ozone molecule (O3).Comment on the predicted bond lengths.

Solution

O O O.... ..

..: : OOO

.. ......

::

O O O

Resonance hybrid

Suggests bonds are equal in length:

r(O=O) < r(O O) < r(O O)

(Without formal charges: see 2.8)

)


2.8 Formal Charge2.8 Formal Charge

Formal charge – the charge an atom would have if the bonding were perfectly covalent in the sense that the atom had exactly a half-share of the bonding electrons.

V = the number of valence electrons in the free atomL = the number of electrons present on the bonded atom as lone pairsB = the number of bonding electrons on the atom

- A Lewis structure in which the formal charges of the individual atoms are closest to zero typically represents the lowest energy arrangement of the atoms and electrons.

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– Formal charge exaggerates the covalent character of bonds by assuming that the electrons are shared equally.

– Oxidation number (state) exaggerates the ionic character of bonds. It represents the atoms as ions, and all the electrons in a bond are assigned to the atom with the lower ionization energy.

O C O0 0 0

O C O-2 +4 -2

formal charge oxidation state

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Self-Test 2.8B

Suggest a likely structure for the oxygen difluoride molecule. Write its Lewis structure and formal charges.

OF F: :..

..

..

..

..

..

Solution

0 00

(not F F O :

:

..

.. ..

.. ..

..: )

0 +1 -1


Chapter 2. CHEMICAL BONDSChapter 2. CHEMICAL BONDS


EXCEPTIONS TO THE OCTET RULE

IONIC VERSUS COVALENT BONDS

2.9 Radicals and Biradicals2.10 Expanded Valence Shells2.11 The Unusual Structures of Some Group 13/III Compounds

2.12 Correcting the Covalent Model: Electronegativity2.13 Correcting the Ionic Model: Polarizability

THE STRENGTH AND LENGTHS OF COVALENT BONDS

2.14 Bond Strength2.15 Variation in Bond Strength2.16 Bond Lengths

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2.9 Radicals and Biradicals2.9 Radicals and Biradicals

34

There are three types of molecules for which the octet rule of Lewis has to be dropped:

1.Odd-electron molecules (radicals)2.High valence molecules (hypervalent

compounds) (See 2.10)3.Low valence molecules (especially of group IIA

and IIIA elements) (See 2.11)


Radicals are species with at least one unpaired electron. They are often highly reactive.

CH3 OH NO NO2

Biradicals: molecules with two unpaired electrons

N O..

::.

Examples


Self-Test 2.9A

Write a Lewis structure for the hydrogenperoxyl radical, HOO. (etc).

Solution

O

H

O::.

....

The unpaired electron isbest accommodated onthe singly bonded O


2.10 Expanded Valence Shells2.10 Expanded Valence Shells

High valence compounds of elements in and beyond period 3 require Lewis structures that disobey the octet rule - more than8 electrons are associated with the central atom in a Lewis structure.These are often called hypervalent compounds.Examples:

37

Sulfuric acid

SO O

OH

OH

:

: :

:

..

..

..

..

P

O

HO OH

OH

:

:

..

..

..

..

..

..

Phosphoric acid

SCl Cl

Cl Cl

xx....

.. ..

..

.... ..

:

: :

:

Sulfur tetrachloride Triiodide anion

II Ixxxx

xx: :.. ..

....

_


– Variable covalence: the ability of certain elements to form different numbers of covalent bonds. Prevalent in elements of the p block inand beyond period 3.

PCl Cl

Cl:

:

:

:..

..

....

..

Obeys octetrule for P

..PCl Cl

Cl:

:

:

:..

..

....

..Cl Cl: :.. ..

....

Disobeys octetrule for P

Phosphorus trichloride

Phosphorus pentachloride

PCl3(l) + Cl2(g) PCl5(s)


Self-Test 2.10A

Write a Lewis structure for xenon tetrafluoride, XeF4, and give the number of electrons in the expanded valence shell.

XeF F

F

F

xx

xx

: :

::

::

..

..

.. ..

..

..12 electrons around Xe

Solution


Self-Test 2.11A

Calculate the formal charge for the two Lewis structures of the phosphate ion shown in (27).

P

Solution

O

O O

O

+1-1-1

-1

-13

_:

:

: :

:

:..

..

.. ......

P

O

O O

O

-1-1

-1

3_:

:

: :

:

..

..

.. ......

0

0

Bigger contribitor toresonance hybrid


2.11 The Unusual Structures of Some Group 2.11 The Unusual Structures of Some Group IIIA/13 CompoundsIIIA/13 Compounds

– Incomplete octet: many of the compounds of group IIIA are characterized by fewer than eight valence electrons around the central atom:

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E.g. Compounds of B and Al BF F

F

Boron trifluoride

: :

: :..

..

..

..

..Al

Cl

Cl Cl

Aluminum trichloride(gaseous)

:..

..:

..

..

: :..


– Their chemistry is dominated by completing octets using a coordinate covalent bond, in which both electrons come from a terminal atom.

BF3 + F- BF4-

BF F

F

: :

: :..

..

..

..

..

F: :..

_

NH3 + BF3 NH3.BF3 B

F

F

F

::

::

.... ..

..

..

N

H

H

H

+ _

AlCl

AlCl

Cl

Cl

Cl

Cl

Aluminum chloride(anhydrous solid)

..

..

..

..:

:

:

:

..

..

..

..

..

..

..

..


IONIC VERSUS COVALENT BONDS(Sections 2.12-2.13)

2.12 Correcting the Covalent Model: 2.12 Correcting the Covalent Model: ElectronegativityElectronegativity

– Many covalent compounds have polar covalent bonds (with partial ionic character).

– Polar covalent bond: a bond in which ionic contributions to resonance result in partial (+ and -) charges (the actual charges on the atoms in a molecule). Bond electrons in the resonance hybrid are shared unevenly.

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H Cl H +_..

..: Cl

..

..::

H Cl+

_

(not Cl..

..: +H:

_)


- electric dipole moment (): size of an electric dipole

Debye

For a Cl-H bond: = ~1.1 D, indicating considerable ionic character (~23%)

– An electric dipole results when a partial positive charge is next to an equal but opposite partial negative charge. It can be represented in two ways:


Electronegativity

45

Electronegativity () was first defined by Linus Pauling (1932) as the electron-pulling power of an atom when it is part of a molecule.

Pauling’s electronegativity scale is based on the difference in bonddissociation energies (in eV) between A-A, B-B, and A-B.

The electronegativity difference between two elements A and B is:

A – B = {D(A-B) – ½[D(A-A) + D(B-B)}1/2 In time, the electronegativity of fluorine was chosen as 4.0, andelectronegativities of other elements are determined relative tothis (maximum) value. See next slide for electronegativities of main group elements (Fig. 2.12)


Electronegativities of Main Group elements (Fig. 2.12)

increases from left to right andfrom bottom to top


Linus Pauling


- Mulliken scale: = ½(I1 + Eea)

(average of the ionizationenergy and electron affinity)

-Rough rules of thumb based on electronegativity difference (Fig. 2.13):

ionicpolar covalent

covalent

Examples: NaCl or KF : ionicC-O : polar covalentCa-O : ionic

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Robert Mulliken devized an electronegativity scale based on ionization energy and electron affinity:

Is a bond, covalent, polar covalent or ionic?


Self-Test 2.12A

In which of the following compounds do the bonds have greater ionic character: (a) P4O10 or (b) PCl3?

Solution

Difference in for P and O = 1.2

Difference in for P and Cl = 1.0

Hence the bonds in (a) have greater ionic character.


2.13 Correcting the Ionic Model: Polarizability2.13 Correcting the Ionic Model: Polarizability

– All ionic bonds have some covalent character.

– Highly polarizable atoms and ions readily undergo a large distortion of their electron cloud.

i.e. large anions and atoms such as I-, Br-, and Cl-

– Polarizing power is the property of ions (and atoms) that cause large distortions of electron clouds.

It increases with decreasing size and increasing charge of a cation

i.e. small and/or highly charged cations Li+, Be2+, Mg2+, and Al3+

50


Self-Test 2.13A

In which of the compounds NaBr and MgBr2 do the bonds have greater covalent character?

Solution

Mg2+ is a smaller cation than Na+, hence MgBr2 has bonds with greater covalent character.


THE STRENGTHS AND LENGTHS OF COVALENT BONDS (Sections 2.14-2.16)

2.14 Bond Strength2.14 Bond Strength Dissociation

energy (D): energy required to separate the bonded atoms

(Fig. 2.15).

- The bond breaking is homolytic, which means that each atom retainsone of the electrons from the bond.

52


- Average dissociation energy is for one type of bond found in many different molecules (Tables 2.3 and 2.4).

E.g. the average C-H single bond dissociation energy gives

average strength of bonds in a selection of organic molecules, such as methane (CH4), ethane (C2H6), and ethene (C2H4)


2.15 Variation in Bond Strength2.15 Variation in Bond Strength

Factors influencing bond strength

54

Figs. 2.16, 2.17

Figs. 2.18, 2.19, 2.20

Bond dissociation energies of N2,O2, and F2 (Fig. 2.16) (in kJ/mol)


Strengths of single and multipleC-C bonds (Fig. 2.16) (in kJ/mol)

Strengths of H-halogen bonds (Fig. 2.18) (in kJ/mol)


2.16 Bond Length2.16 Bond Length

Bond length: the distance between the centers of two atoms joined by a covalent bond

- It corresponds to the internuclear distance at the potential energy minimum for the two atoms

- It affects the overall size and shape of a molecule, evaluated by using spectroscopy or x-ray diffraction methods

Factors influencing bond length

56


Covalent radius is the contribution an atom to the length of a covalent bond.

- Bond length is approximately the sum of the covalent radii of the two atoms.

- Decreases from left to right(increasing Zeff)

- Increases in going down a group(size of valence shells and bettershielding by inner core electrons)

57


INFRARED SPECTROSCOPY

Infrared radiation: electromagnetic radiation with longer wavelengths (lower frequencies) than red light ~ 1000 nm or ~ 3×1014 Hz

-Molecules become vibrationally excited when they absorbinfrared radiation of correct energy.

- “stretching” mode: the atoms moving closer and away again. “bending” mode: bond angles periodically increase and decrease.

Vibrational frequencies

- The stiffness of a bond measured by its force constant, k

Force = -k × displacement by Hooke’s law

- Vibrational frequency, , of a bond between two atoms A and B of mass mA and mB:

m = effective mass (or reduced mass)

58


Normal (fundamental) modes of vibration

A nonlinear molecule consisting of N atoms→ 3N-6 normal modes

A linear molecule → 3N-5 normal modes

Characteristic frequencies of functional groups detectable in a spectrum

- fingerprint region: a complex series of absorptions

i.e. H2O, n = 3 → 3 normal modes

CO2, n = 3 → 4 normal modes

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2012 general chemistry i chapter 2. chemical bonds 2012 general chemistry i ionic bonds covalent...

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