chapter 2 chemical bonds

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.Ener

General Chemistry I 1Chapter 2.

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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 cationsand anions

stacked together in a regular pattern, called a crystal lattice.

2.1 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.

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

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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

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2.2 Lewis Symbols

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

- Cations and anions

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.

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Gilbert N Lewis

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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 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):

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

the lattice energy.

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2.4 Interactions Between Ions

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

➢ Lattice energy: the difference in energy between the ions packedtogether in a solid and the ions widely separated as a gas

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

vacuum permittivity constant.

- Molar potential energy of a three-dimensional crystal:

➢ The factor A is the Madelung constant,dependent on how the ions are arrangedabout one another in the 3-dimensionallattice.

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

General Chemistry I 13This 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), z2is replaced by the absolutevalue of z1z2 (its value without the negative sign).

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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.

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Attraction, Repulsion and the Born-Mayer EquationThe 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.

General Chemistry I 16➢ Born-Mayer equation

with d* = 34.5 pmrepulsive 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.

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COVALENT BONDS (Sections 2.5-2.8)

2.5 Lewis Structures

➢ Covalent bond - a pair of electrons shared between two atoms

- Octet rule: in covalent bond formation, atoms go as far as possibletoward 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.

Some Definitions According to Lewis Theory

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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

Three lone pairs onboth atoms

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2.6 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 numbers

E.g CO24 + 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.

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.

– 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., H2SO4is 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 CH3group and one COOH group.

Phosphorous acid,P(OH)3

Acetone, (CH3)2COC

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Self-Test 2.5A

Write a Lewis structure for the cyanate ion, NCO

-(sometime written CNO-).

_

H

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2.7 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 contribution

to 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:

_The resonance hybrid is

_

General Chemistry I 29Additional Self-Test

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

::

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2.8 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 atom

L = 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.

are assigned to the atom with the lower ionization energy.

formal chargeoxidation state

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General Chemistry I 32

Self-Test 2.8B

Suggest a likely structure for the oxygen difluoride

molecule Write its Lewis structure and formal charges.

:

General Chemistry I 33Chapter 2 CHEMICAL BONDS

2012 General Chemistry I

EXCEPTIONS TO THE OCTET RULE

IONIC VERSUS COVALENT BONDS

2.9 Radicals and Biradicals2.10 Expanded Valence Shells

2.11 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 Strength

2.15 Variation in Bond Strength2.16 Bond Lengths

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2.9 Radicals and Biradicals

➢ 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)

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➢ Radicals are species with at least one unpaired electron

They are often highly reactive.

Biradicals: molecules with two unpaired electrons

Examples

The unpaired electron isbest accommodated onthe singly bonded O

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2.10 Expanded Valence Shells

High valence compounds of elements in and beyond period 3 require Lewis structures that disobey the octet rule - more than

8 electrons are associated with the central atom in a Lewis structure.

These are often called hypervalent compounds.

Sulfuric acidS

Phosphoric acid

_

General Chemistry I 38– 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.

Obeys octetrule for P

Disobeys octetrule for P

Phosphorus trichloride

Phosphorus pentachloridePCl3(l) + Cl2(g) →PCl5(s)

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Self-Test 2.10A

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

12 electrons around XeSolution

Bigger contribitor toresonance hybrid

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2.11 The Unusual Structures of Some Group IIIA/13 Compounds

– Incomplete octet: many of the compounds of group IIIA

are characterized by fewer than eight valence electrons around the central atom:

E.g Compounds of B and AlFBFF

Boron trifluoride

: :

Aluminum trichloride(gaseous)

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– Their chemistry is dominated by completing octets using a

coordinate covalent bond, in which both electrons come from a terminal atom.

: :

NH3+ BF3NH3.BF3BF

Aluminum chloride(anhydrous solid)

– Polar covalent bond: a bond in which ionic contributions to

resonance result in partial (d+ and d-) charges (the actual charges on the atoms in a molecule) Bond electrons in the resonance hybrid are shared unevenly.

)

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➢ Electronegativity (c)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:

cA –cB = {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)

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Electronegativities of Main Group elements (Fig 2.12)

increases from left to right and

from

bottom to top

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Linus Pauling

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- Mulliken scale: c = ½(I1 + Eea)

(average of the ionizationenergy and electron affinity)

-Rough rules of thumb based on

electronegativity difference (Fig 2.13):

polar covalentcovalentExamples:

Difference in c for P and O = 1.2

Difference in c for P and Cl = 1.0

Hence the bonds in (a) have greater ionic character

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2.13 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+

(Fig 2.15).

- The bond breaking is homolytic, which means that each atom retains

one of the electrons from the bond.

General Chemistry I 53- 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)

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2.15 Variation in Bond Strength

Factors influencing bond strength

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)

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2.16 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 energyminimum 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

(increasing Zeff)-Increases in going down a group

(size of valence shells and bettershielding by inner core electrons)

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-“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, n, of a bond between two atoms A and Bof mass mAand mB:

m = effective mass (or reduced mass)

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➢ Normal (fundamental) modes of vibration

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

A linear molecule → 3N-5 normal modes

General Chemistry I 60MOLECULAR SHAPE AND STRUCTURE

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THE VSEPR MODEL

VALENCE-BOND THEORY

3.1 The Basic VSEPR Model

3.2 Molecules with Lone Pairs on the Central Atom3.3 Polar Molecules

3.4 Sigma and Pi Bonds

3.5 Electron Promotion and the Hybridization of Orbitals3.6 Other Common Types of Hybridization

3.7 Characteristics of Multiple Bonds