Chapter 5

Basic Concepts of Chemical Bonding

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Types of Bonds

Chemical Bond – a strong attraction between atoms

Ionic Bond – an electrostatic interaction where one atom donates an electron to the other

Covalent Bond – atoms share electrons

Metallic Bond – atoms bonded to several others and electrons can move freely.

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Types of Bonding

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• In a Lewis structure, we represent the valence electrons of main-group elements as dots surrounding the symbol for the element – aka electron dot structures

• the symbol of the element is used to represent the nucleus and inner electrons

• dots around the symbol represent valence electrons

• first four dots are for the two s orbital electrons and two p orbital electrons, the final three dots represent paired p orbital electrons

Lewis Structures of Atoms

•Li •Be• •B••

•C••• •N•

••

• •O•••

•• •F•••

••• •Ne•

••

•••• 4

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Lewis Dot Symbols

Ionic Bonding

• lattice energy is a measure of stabilization

• lattice energy – the energy required to completely separate 1 mol of solid ionic components into its gaseous ions.

• NaCl (s) Na+ (g) + Cl– (g) H = +788kJ/mol

• the reverse

• Na+ (g) + Cl– (g) NaCl (s) H = –788kJ/mol

ionic substances are brittle, crystalline, with high melting points

•Na •Cl• ••••

•+

[Ne] 3s1 [Ne] 3s2 3p5

Na+ •Cl•••

••••

–1

[Ne] [Ar]

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Ionic Bonding

• The magnitude of the lattice energy of an ionic solid depends on the charges of the ions, their sizes, and their arrangement in the solid.

• The potential energy of two interacting charges is given by

Eel = Q1Q2

d

= 8.99 x 109 J-mC2

Q1Q2 = are the charges on atom 1 and atom 2, d is the distance between them, and

The Crystal Structure of Sodium Chloride

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Lattice Energy Diagram for Sodium Chloride

lattice energy Born – HaberCycle

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5.2

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Covalent Bondingmost substances have covalent type bonding – which means they share electrons.

• Electrons that are shared by atoms are called bonding pairs

• Electrons that are not shared by atoms but belong to a particular atom are called lone pairs – aka nonbonding pairs

....OSO.. .... ..

.. ....Bonding pairs Lone pairs

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Electronegativity

• electronegativity – the ability of an atom in a molecule to attract electron density toward itself

increase

increase

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• follows electron affinity

Electronegativities of the Elements

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Electronegativity Difference and Bond Type• pure covalent - electronegativity between bonded atoms is

0 – equal sharing

“100%”

0 0.4 2.0 4.0

4% 51%Percent Ionic Character

Electronegativity Difference

• nonpolar covalent – electronegativity between bonded atoms is 0.1 to 0.4

• polar covalent – electronegativity between bonded atoms is 0.5 to 1.9

• ionic – electronegativity between bonded atoms is larger than or equal to 2.0

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Polar Covalent Bonds

• Though atoms often form compounds by sharing electrons, the electrons are not always shared equally.

• Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.

• Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

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• When two atoms share electrons unequally, a bond dipole results.

• The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:

• It is measured in debyes (D).

Polar Covalent Bonds

= Qr

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The greater the difference in electronegativity, the more polar is the bond.

Polar Covalent Bonds

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Lewis Structures

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

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Lewis Structures

2. Determine the central atom3. Write symbols for the atoms and show which atoms are attached to each other and connect them with a single bond (a dash representing two atoms) – subtract from total (2 per dash) 4. Complete octet around all the atoms bonded to the central atom Subtract from the total

5. Place any leftover electrons around the central atom.

6. for any atom lacking an octet form double or triple bonds

1. Sum valence electrons from all atoms – add or subtract e–’s for charge

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Writing Lewis Structures

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.

PCl3

Keep track of the electrons:

5 + 3(7) = 26

• If it is an anion, add one electron for each negative charge.

• If it is a cation, subtract one electron for each positive charge.

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2. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

Keep track of the electrons:

PCl

Cl

Cl

26 − 6 = 20

1 ••=

Writing Lewis Structures

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3. Fill the octets of the outer atoms.

Keep track of the electrons:

26 − 6 = 20; 20 − 18 = 2

Writing Lewis Structures

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4. Put left over electrons around the central atom.

Keep track of the electrons:

26 − 6 = 20; 20 − 18 = 2 place extra e– around central atom;2 − 2 = 0

Writing Lewis Structures

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5. If you run out of electrons before the central atom has an octet form multiple bonds until it does.

Writing Lewis Structures

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• To determine the best Lewis structure assign formal charges.

• For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.

• Subtract that from the number of valence electrons for that atom: the difference is its formal charge.

Writing Lewis Structures

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• The best Lewis structure is the one with the fewest charges.

• It is best to put a negative charge on the most electronegative atom.

Writing Lewis Structures

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ResonanceThis is the best Lewis structure that can be drawn for ozone, O3.

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Resonance• But this is at odds with the true, observed structure of ozone.

• both O—O bonds are the same length.

• both outer oxygens have a charge of −1/2.

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Resonance• One Lewis structure cannot accurately depict a molecule

like ozone.

• We use multiple structures, resonance structures, to describe the molecule.

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Resonance

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

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Resonance• In truth, the electrons that form the second C—O bond in

the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

• They are not localized; they are delocalized.

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Resonance• The organic compound benzene, C6H6, has two resonance

structures.

• It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

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Exceptions to the Octet Rule

• There are three types of ions or molecules that do not follow the octet rule:

• ions or molecules with an odd number of electrons,

• ions or molecules with less than an octet,

• ions or molecules with more than eight valence electrons (an expanded octet).

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Odd Number of Electrons

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

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• Consider BF3:

• Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.

• This would not be an accurate picture of the distribution of electrons in BF3.

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Fewer Than Eight Electrons

• Therefore, structures that put a double bond between boron and fluorine are much less reasonable than the one that leaves boron with only 6 valence electrons.

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More Than Eight Electrons• The only way PCl5 can exist is if phosphorus has 10

electrons around it.• It is allowed to expand the octet of atoms on the third row

or below. Presumably d orbitals in these atoms participate in bonding.

• The octet rule works really well for C, N, O, and F. Atoms after period 2 may or may not have an octet.

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Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

More Than Eight Electrons

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• This eliminates the charge on the phosphorus and the charge on one of the oxygens.

More Than Eight Electrons

• The lesson is: When the central atom is on the third row or below and expanding its octet eliminates some formal charges, do so.

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Covalent Bond Strength

• Most simply, the strength of a bond is measured by determining how much energy is required to break the bond.

• This is the bond enthalpy.

• The bond enthalpy for a Cl—Cl bond, (Cl— Cl), is measured to be 242 kJ/mol.

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