AP Ch 8 Bonding (2005).pptx

Ch. 8 Chemical Bonding

•_{Chemical bonds hold atoms together.} •_{There are 3 types of chemical bonds:}

-_{Ionic bonds (electrostatic forces that hold ions}

together…)

-_{Example: Na}+Cl-, K+Br

--_{Covalent bonds (result from sharing electrons}

between atoms…)

-Example: H₂, NH₃

-_{Metallic bonds (refers to metal nuclei floating in a}

sea of electrons…)

Lewis Symbols: Electron dots

•_{Valence Electrons: outer-most electrons}

–_{determined by the “Group A #” from the periodic table}

–_{Exceptions: d or f-block = 2 valence electrons & Helium =2}

-Octet Rule

•_{Atoms often gain or lose or share electrons to fill their}

valence shell with 8 electrons to achieve a noble gas configuration.

–_{Exceptions: * Hydrogen needs only 2 e}- _{to be filled.}

* Some nonmetals can have more or less than 8.

Ionic Bonding

• _{Ionic bonds—transfer of electrons}

- _{Form between an element of low ionization energy (not}

much energy required to pull off an electron) and an element of high electron affinity (lots of energy is released when an electron is added to its outer shell).

Ionic Bonding Energies

• _{Consider the reaction between sodium and chlorine:}

Na_(s) + ½Cl_2(g)  NaCl_(s) ∆Hº_f = –410.9 kJ/mol - The reaction is violently exothermic.

- We infer that the NaCl is more stable than its constituent elements.

• _{Here’s another way to look at the energy of ionic bond formation:}

–_{Sodium loses 1 electron…Na  Na}+ _{+ 1 e- Requires 5.1 eV of energy}

– _{Chlorine gains l electron…Cl + 1 e-  Cl}- _{Releases 3.6 eV of energy}

– _{NaCl forms… Na}+ _{+ Cl}- _{ [Na+][ Cl-] Releases 5.2 eV of energy}

[1 eV (electron volt) = 1.602 x 10-19 _J]

•_{The energy released is greater than the energy required, therefore}

Lattice Energy

•

separate a mole of a solid ionic compound into its ions.

Ions

•

Metals lose electrons to form smaller (+)

cations

.

•

Nonmetals gain e- to form larger (-)

anions

.

–_{The # of e- gained or lost depends on how many they need to}

gain or lose to get to a noble gas configuration. Only then will they become stable.

•_{Groups of atoms can have charges too. They are called}

polyatomic ions.

–_{The atoms share electrons (covalent bonds) but the group} still has an overall charge.

Examples: [NH₄+] , [CO

Covalent Bonds

•_{Atoms share electrons to fill their valence shell.} •_{Usually form between 2 nonmetals}

–_{Lewis Structures: represent covalent bonds as 2 dots between the atoms}

- A line can also be used to represent 2 shared e-_{’s (or one covalent bond.)}

Cl + Cl Cl Cl

Cl Cl H F H O

H

H N H

H H C

H

Multiple Covalent Bonds

•_{Single bond = 2 electrons shared …(1 pair)} •_{Double bond= 4 electrons shared…(2 pairs)} •_{Triple bond= 6 electrons shared…(3 pairs)}

Bond Lengths & Bond Strengths

•_{Single bonds are the longest and weakest covalent}

bonds.

•_{Triple are the shortest and strongest covalent bonds.} •_{Ionic bonds are much stronger than covalent bonds.}

Rules For Drawing Lewis Structures

Rule 1) How many electrons are possible around an atom?

• _{For hydrogen, only 2 electrons are possible, therefore only one bond!}

• _{Second row elements usually try to get 8 e}-_{. Notable Exception: Boron}

needs only 6.

• _{Third & Fourth Row usually have 8 e}- _{but can expand to get 10 or more.}

Rule 2) Drawing the Lewis Structure:

• _{First arrange the atoms around the central atom (usually the least} electronegative one, but never hydrogen!

• _{Count the total # of valence electrons in the molecule. If it is an ion, add 1} for each (-) charge or subtract 1 for each (+) charge.

• _{Distribute the electrons keeping Rule #1 in mind. If you have too many} electrons, look for double or triple bonds, or place the extras around the “3rd

Rules For Drawing Lewis Structures

Rule 3) For odd-numbered valence electrons:

• _{If you must cheat an element out of 8 e}- _{and only give it 7, then cheat}

the least electronegative element.

Rule 4) Resonance Structures:

• _{If there is more than one way to draw the Lewis structure, show them all.}

Rule 5) Nature doesn’t know anything about Rules1-4.

Resonance

Resonance Structures:

• _{The ability to draw more than one “correct” Lewis structure.}

Benzene (C₆H₆)

• _{The true structure for the molecule is somewhere “in between” the}

resonance structures.

Bond Polarity

•_{Bond polarity helps to describe the sharing of the}

electrons between atoms. There are 3 possibilities…

–_{Nonpolar covalent: equal sharing of the e}- pair

–_{Polar covalent: unequal sharing of the e}- _pair

–_{Ionic: transfer of valence e}- from the metal to the nonmetal

•_{A molecule that has one side slightly positive and one}

side slightly negative is said to be a “dipole.”

•_{The positive end (or pole) in a polar bond is represented}

+ and the negative pole -. Arrow can also show dipoles.

-Bond Polarity & Electronegativity

(How can you tell what type of bond will form?)

•_{Electronegativity: describes an atom’s attraction to the e}

-pair in a bond…(It’s a number from 0 to 4.0)

•_{The difference between electronegativities indicate}

whether a bond will be nonpolar, polar or ionic.

•_{There is no sharp distinction between bonding types.}

•_{In General: Nonpolar = 0-0.4 Polar= 0.5-2.0 Ionic= Above 2.0}

Lattice Energy and Polarity

Lattice energy increases as the electronegativity

between the atoms in an ionic compound

Formal Charge

• _{Formal Charge: The formal charge of an atom is the}

charge that an atom (in a molecule) would have if all of the atoms had the same electronegativity.

• _{To calculate formal charge:}

(valence e

- # of bonds - lone pair e

)

Practice: Determine the formal charge on C and N.

C= 4 – 3 – 2 = -1 N= 5 – 3 – 2 = 0

Formal Charge

• _{The most stable Lewis structure has the smallest formal}

charge on each atom and the most negative formal charge on the most electronegative atoms.

Hybridization

•

Hybridization is the idea that atomic

orbitals fuse to form newly hybridized

orbitals.

•

Example: CH

•

The hybridization corresponds to the

Sigma & Pi Bonds

•

Sigma Bonds: are the first bonds formed

between atoms, use s orbitals

•

Pi Bonds: are the 2

and 3

bonds

formed between atoms, use p-orbitals

– _{Double bond: 1 sigma 1 pi bond}

Strengths of Covalent Bonds

•

The energy required to break a covalent bond is

called the

bond dissociation enthalpy

,

D

.

That is, for the Cl₂ molecule, D(Cl-Cl) is given by H for the reaction:

Cl_2(g)  2Cl_(g) H = 242 kJ When more than one bond is broken…

CH_4(g)  C_(g) + 4H_(g) H = 1660 kJ …the bond enthalpy is a fraction of H for the “atomization reaction”:

Strengths of Covalent Bonds

The bond enthalpy

(energy) for a given set of atoms depends on the rest of the molecule of

which it is a part. An average

bond enthalpy is therefore shown.

The energy

Using “D” to determine ∆H

∆H

=



D

-



D

• _{Practice Problem}

CH_4(g) + Cl_2(g)  CH₃Cl_(g) + HCl_(g) H_(rxn) = ?

- _{In this reaction one C-H bond and one Cl-Cl bond is broken}

while one C-Cl bond and one H-Cl bond is formed.

So… ∆H_(rxn) = [D(C-H) + D(Cl-Cl)] - [D(C-Cl) + D(H-Cl)]

= 413kJ + 242 kJ – 328 kJ – 431 kJ = – 104 kJ

• _{The overall reaction is exothermic which means that the bonds}

formed are stronger than the bonds broken.