Basic Concepts in Chemical Bonding

Basic Concepts in Chemical Bonding

Chapter 8

What is a Bond?

• A force that holds atoms together.

– Why?• We will look at it in terms of energy.• Bond energy the energy required to break a bond.

• Why are compounds formed?– Because it gives the system the lowest energy.

Lewis Symbols and the Octet RuleThere are three types of chemical

bonds (intramolecular force)

- Ionic Bond - electrostatic attraction between ions of opposite charge (MgO)

- Covalent Bond - sharing of electrons between two atoms (Br2)

- Metallic Bond - sharing of electrons between several atoms (Cu).

Keeping Track of Electrons

• The electrons responsible for the chemical properties of atoms are those in the outer energy level.

– Valence electrons - electrons in the outer energy level.

– Core electrons -those in the energy levels below.

Keeping Track of Electrons

• Atoms in the same column– Have the same outer electron configuration.– Have the same valence electrons.

• Easily found by looking up the group number on the periodic table.

• Group 2A - Be, Mg, Ca, etc.-– 2 valence electrons

Electron Dot diagrams (Lewis Symbols)

• A way of keeping track of valence electrons.

• How to write them– Write the symbol.– Put one dot for each valence

electron– Don’t pair up until you have

too (Hund’s rule).

X

Lewis Symbols and the Octet Rule

Octet rule – Atoms tend to gain, lose or share electrons until they are surrounded by eight valence electrons.

Ionic Bonding Energetics of Ionic Bonding

As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.

Energetics of Ionic Bonding

We get 349 kJ/mol back by giving electrons to chlorine.


• But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!


• There must be a third piece to the puzzle.

• What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

Lattice Energy

• This third piece of the puzzle is the lattice energy:The energy required to completely separate a mole of

a solid ionic compound into its gaseous ions.• The energy associated with electrostatic

interactions is governed by Coulomb’s law:

Eel = Q1Q2

d

Coulomb’s equation:

Q1, Q2 = charge on ionsk = 8.99 x 109 J-m/c2

d = distance between ions

Lattice energy is greater with more highly charged ions.

dQQkE 21

Energetics of Ionic Bond Formation

Lattice Energy

• Lattice energy, then, increases with the charge on the ions.

• It also increases with decreasing size of ions.


By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.


• These phenomena also helps explain the “octet rule.”

• Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

• These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital.

• Electron configurations can predict stable ion formation:• Mg: [Ne]3s2

• Mg+: [Ne]3s1 not stable• Mg2+: [Ne] stable• Cl: [Ne]3s23p5

• Cl-: [Ne]3s23p6 = [Ar] stable

Electron Configurations of Ions of the Representative Elements

Transition Metal Ions• Lattice energies compensate for the loss of up to three

electrons.• In general, electrons are removed from orbitals in order

of decreasing n (i.e. electrons are removed from 4s before the 3d).

Polyatomic Ions• Polyatomic ions are formed when there is an overall

charge on a compound containing covalent bonds. E.g. SO4

2-, NO3-.

Transition Metal Ions and Polyatomic Ions

Covalent bonds

• Nonmetals hold onto their valence electrons

• They can’t give away electrons to bond

• Still want noble gas configuration

• Get it by sharing valence electrons with each other

• By sharing both atoms get to count the electrons toward noble gas configuration.

Covalent Bonding

• There are several electrostatic interactions in these bonds:

– Attractions between electrons and nuclei

– Repulsions between electrons

– Repulsions between nuclei

Lewis Structures

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

Writing Lewis Structures

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

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

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

PCl3

5 + 3(7) = 26


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:

26 - 6 = 20


3. Fill the octets of the outer atoms.


26 - 6 = 20 - 18 = 2


4. Fill the octet of the central atom.


26 - 6 = 20 - 18 = 2 - 2 = 0


5. If you run out of electrons before the central atom has an octet…

…form multiple bonds until it does.

Multiple Bonds

• Sometimes atoms share more than one pair of valence electrons.

• A double bond is when atoms share two pair (4) of electrons.

• A triple bond is when atoms share three pair (6) of electrons.

• Generally, bond distances decrease as we move from single through double to triple bonds.

H H O O N N


• Then assign formal charges.

– HUH???? Formal What???

• It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms.

• To determine which structure is most reasonable, we use formal charge.

• Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity (i.e., if the electrons are shared equally between atoms).

Formal Charge

1) A Lewis structure in which all formal charges in a molecule are equal to zero is preferable to one in which some formal charges are not zero.

2) If a Lewis structure has non-zero formal charges, the one with the fewest nonzero formal charges is preferred.

3) A Lewis structure with one large formal charge is preferable to one with several small formal charges.

4) A Lewis structure with adjacent formal charges should have opposite signs.

5) When choosing between several Lewis structures, the structure with negative formal charges on the more electronegative atom is preferable.

Using Formal Charge

Writing Lewis Structures– 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• The best Lewis structure…

– …is the one with the fewest charges.

– …puts a negative charge on the most electronegative atom.

• Some molecules are not well described by Lewis Structures.

• Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms

• Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).

OO

O

Resonance Structures

Resonance

This is the Lewis structure we would draw for ozone, O3.

-

+

Resonance

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

Resonance• But this is at odds

with the true, observed structure of ozone, in which…

– …both O—O bonds are the same length.

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, but rather are delocalized.

• Benzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom.

• There are alternating double and single bonds between the C atoms.

• Experimentally, the C-C bonds in benzene are all the same length.

• Experimentally, benzene is planar.

Resonance in Benzene

• We write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring:

• Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor).

or

Resonance in Benzene

Example

• Draw the lewis dot structure for the following:– Nitrate ion– Sufite ion– Carbon dioxide– Thiocyanate ion– Carbonate ion– Formate ion

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).

Odd Number of Electrons

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

N O N O

Fewer Than Eight Electrons

• 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.


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


The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.

• This is the largest class of exceptions.

• Atoms from the 3rd period onwards can accommodate more than an octet.

• Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.

More than an Octet

More Than An Octet

• The only way PCl5 can exist is if phosphorus has 10 electrons around it.

More Than Eight Electrons

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

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

• The lesson is: When the central atom is on the 3rd row or below it can expand its octet to eliminate some formal charges.

Bond Polarity and Electronegativity

• Although 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.

• Electronegativity: The ability of one atoms in a molecule to attract electrons to itself.

• Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F).

• Electronegativity increases • across a period and • down a group.

Electronegativity

Bond Polarity and Electronegativity- A chemical bond between elements with large

differences in electronegativity will shift the electrons to the atom with the higher electronegativity.

- The positive end (or pole) in a polar bond is represented + and the negative pole -.

- This is called a polar covalent bond.

- If the electronegativity difference is small, the bond is nonpolar; if it is large, it is a polar bond.

• Difference in electronegativity is a gauge of bond polarity:• electronegativity differences around 0 result in non-polar

covalent bonds (equal or almost equal sharing of electrons);

• electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons);

• electronegativity differences around 3 result in ionic bonds (transfer of electrons).

Bond Polarity and Electronegativity

Example

• Which bond is more polar. Indicate in each case which atom has the partial negative charge.

• – B-Cl or C-Cl– P-F or P-Cl– S-Cl, S-Br, Se-Cl or Se-Br

Polar Covalent Bonds• 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:

= Qr• It is measured in debyes (D).

Polar Covalent Bonds

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

Strengths of Covalent Bonds

• 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,

D(Cl—Cl), is measured to be 242 kJ/mol.

Average Bond Enthalpies

• This table lists the average bond enthalpies for many different types of bonds.

• Average bond enthalpies are positive, because bond breaking is an endothermic process.

Average Bond Enthalpies

NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C—H bonds in methane, CH4, will be a bit different than theC—H bond in chloroform, CHCl3.

Enthalpies of Reaction

• Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed.

• In other words, Hrxn = (bond enthalpies of bonds broken) -

(bond enthalpies of bonds formed)


CH4(g) + Cl2(g)

CH3Cl(g) + HCl(g)

In this example, oneC—H bond and oneCl—Cl bond are broken; one C—Cl and one H—Cl bond are formed.


So,Hrxn = [D(C—H) + D(Cl—Cl) - [D(C—Cl) + D(H—Cl)

= [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)]= (655 kJ) - (759 kJ)= -104 kJ

• We know that multiple bonds are shorter than single bonds.

• We can show that multiple bonds are stronger than single bonds.

• As the number of bonds between atoms increases, the atoms are held closer and more tightly together.

Bond Enthalpy and Bond Length

Basic Concepts in Chemical Bonding

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