1TYPES OF CHEMICAL BONDING Ionic Bonding - Bond between ions whose charges attract each other - One atom gives electrons and one atom takes electrons. Example Covalent Bonding - two atoms each sharing electrons within a molecular orbital F – F - both fluorine atoms “own” electrons within bond OCTET RULE – Generally atoms prefer electron configurations with 8 valence electrons. - Filled s and p subshells Atoms bond with each other so that every atom has 8 electrons in its outer shell. - Atoms may take electrons from each other or they may share electrons. Octet rule is able to explain a lot of chemistry. I. e., it is able to explain why certain elements combine together in specific proportions. Exceptions to the octet rule are plentiful. We will consider these later.

Na+ Cl-

ionic bond

covalent bond

2LEWIS STRUCTURES Only valence electrons are important in bonding. Lewis dot structures show valence electrons surrounding atom. We visualize the four valence orbitals of an atom as the sides of a box. Electrons are put

into orbitals according to Hund’s rule. Examples Be has 2 valence electrons. Therefore Lewis structure is Be N has 5 valence electrons.

Therefore Lewis structure is N

Br has 7 valence electrons.

Therefore Lewis structure is Br

C has 4 valence electrons.

Therefore Lewis structure is C

IONIC BONDING (again) - Oppositely charged ions attract each other. - Metal atoms lose e- and nonmetal atoms gain e-. - Ions attract each other to form ionic lattice (crystal). Lewis structures can be used to illustrate ionic bonding. Consider Potassium and Bromine

BrK + → Br-K+

Consider Calcium and Fluorine

F+ → Ca2+ F-Ca

FF-

Lewis structures are much more illuminating when we consider the sharing of electrons (covalent bonding).

3Lattice Energy – energy of released when positive and negative ions form crystal lattice due

to their attraction for each other. Creation of ionic compound can be decomposed into many small steps. I. e., lattice energy can be decomposed into many smaller energy steps. COULOMB’S LAW - fundamental law of physics (hugely important in chemistry)

1 2

12

q qE kr

=

Three items of note about Coulomb’s law 1. If the energy is negative, ions must be oppositely charged. 2. The higher the charge, the greater the energy. 3. The higher the distance, the smaller the energy. Coulomb’s Law and Lattice Energy - Coulomb’s law can be used to make relative assessments of lattice energy. Example: Which compound has greater lattice energy: LiBr or CaO? Li+ – Br- Ca2+ – O2- ← higher charges means greater lattice energy - Distance between ions is also important but to a much lesser degree. - Smaller distance means higher lattice energy. - Note that since ions are oppositely charged, the lattice energy is a negative value. - Greater lattice energy means a more negative value. Example: Which compound has higher lattice energy: NaCl or CsCl? Charges are identical; however, cesium ion is much larger than sodium ion. Na+ --- Cl- Cs+-------------Cl- ← greater distance between charges means smaller lattice energy Therefore, NaCl has higher lattice energy.

q1 – charge of ion 1 q2 – charge of ion 2

r12 – distance between ion 1 and ion 2

E – energy of interaction

k – proportionality constant

4COVALENT BONDING - Atoms that are bonded share valence electrons. - Sharing is what creates covalent bond. - When atoms bond covalently, new entity termed molecule is formed. Covalent Bonding in Diatomic Molecules Hydrogen – H2

H + →H H H - each atom contributes e- to the bond - e- in bond belongs to both atoms (e- are shared) - when sharing e-, each atom has full shell (like He) - note hydrogen is an exception to octet rule Fluorine – F2

F + →F FF

- Since both atoms share electrons in bond, both atoms have 8 valence electrons, octet rule

is satisfied. Hydrogen fluoride (hydrofluoric acid)

H + →F FH

Oxygen

O + →O OO

- note that a molecule can have more than one covalent bond Nitrogen

N + →N NN

5Bonding in Polyatomic Molecules Water

HH

O+ →OH

H

Ammonia

HH

N+ →H

NH

H

H Carbon Dioxide

O →OC CO O

Useful reminders in polyatomic bonding 1. Carbon will form 4 covalent bonds per atom. 2. Oxygen will form 2 covalent bonds per atom. 3. Nitrogen will form 3 covalent bonds per atom. 4. Hydrogen will form 1 covalent bond per atom. 5. Halogens will form 1 covalent bond per atom. DRAWING LEWIS DOT STRUCTURES FOR MOLECULES Do the following steps in order. 1. Identify whether compound is molecule or polyatomic ionic compound. -I. e., Identify the charge of the molecule or ion. 2. Sum the number of valence electrons from all atoms in the molecule. 3. Add or subtract appropriate number of electrons for polyatomic ion. 4. Identify the central atom and decide how other atoms are bonded to it. - least electronegative atom is usually central atom - hydrogen is never central atom - oxygen is rarely central atom 5. Draw bonds between atoms. (Subtract 2 e- for each bond from total number of valence e-) 6. Complete octets of peripheral (outer) atoms. (Subtract each electron used from total

number of valence e-) 7. Put remaining electrons in pairs on central atom. 8. If central atom has too few electrons to complete octet, change lone pairs on peripheral

atoms into a double bond between central atom and peripheral atom.

6Example: Draw the Lewis structure for PCl3. # of valence e- of P = 5 # of valence e- of Cl = 7 # of e- = 5 + 3(7) = 26 e- Central atom is P. Example: Draw the Lewis structure for NaOH. OH- is the polyatomic ion # of valence e- of O = 6 # of valence e- of H = 1 # of valence e- added = 1 # of e- = 6 + 1 + 1 = 8

O H Including the sodium ion we could write

Na O H+ -

Example: Draw the Lewis structure of SO3. 1. # of e- = 6 + 3(6) = 24 e- 3. Central atom is S. 7. Central atom has too few electrons, use multiple bond. - doesn’t matter which oxygen atom makes double bond.

Note: In NaOH, there is ionic and covalent bonding.

P Cl

Cl

Cl

S O

O

O

S O

O

O

Used in the manufacture of pesticides and flame retardants.

NaOH is also known as lye and is the essential ingredient in drain cleaners.

SO3 is a component of pollution from burning coal. When mixed with water it makes a component of acid rain.

7 Example: Draw the Lewis Structure of Fe3(PO4)2. PO4

3- is the polyatomic ion # of e- = 5 + 4(6) + 3 = 32 e- Note: three electrons are added. Central atom is P.

O OP

O

O

O OP

O

O

3-

Fe2+

Fe2+

Fe2+

O OP

O

O

3-

Example: Draw the Lewis Structure of NH4Cl. # of e- = 5 + 4(1) – 1 = 8 e- Note: one electron is subtracted. Central atom is N. RESONANCE - Sometimes more than one correct Lewis structure can be drawn. - In that case, actual structure is a blend of correct structures. Example: Draw a Lewis structure for the nitrate ion, NO3

- # of e- = 5 + 3(6) + 1 = 24 e- N is the central atom. 6N + 2 rule = 26 ⇒ double bond is expected

N H

H

H

H +

N O

O

O

Ammonium chloride can be used as an expectorant, an agent to loosen mucus in the respiratory tracts.

Iron (II) phosphate is used by organic farmers to control snails and slugs.

8 Change lone pair into double bond to satisfy octet rule. - The oxygen atom on the left was chosen arbitrarily. Any of the three atoms could

have been used to form the double bond. Choosing the oxygen atom on the

right would yield the structure: Choosing the oxygen atom on the

bottom would yield the structure: The actual structure is a blend of resonance structures. - Double-headed arrows indicate resonance structures. - Double bond is not confined to a single O – N pair. - Double bond is distributed over all three O – N pairs. - Resonance is good, resonance makes a molecule more stable.

N O

O

O

N O

O

O

N O

O

O

N O

O

O N O

O

O

N O

O

O

9BOND LENGTHS The length of the covalent bond depends on two items. 1. The atomic radii of the atoms involved in the bond. Example: d(C – H) = 109 pm r(H) = 37 pm d(C – F) = 133 pm r(F) = 72 pm d(C – Cl) = 177 pm r(Cl) = 100 pm d(C – Br) = 194 pm r(Br) = 114 pm d(C – I) = 213 pm r(I) = 133 pm r(C) = 77 pm 2. The number of bonds between the atoms. Example: d(N – N) = 146 pm d(C – C) = 154 pm d(N = N) = 122 pm d(C = C) = 134 pm d(N ≡ N) = 110 pm d(C ≡ C) = 121 pm Generally, the shorter the bond length, the higher the bond dissociation energy. d(C – C) = 154 pm Edis = 348 kJ/mol d(C = C) = 134 pm Edis = 614 kJ/mol d(C ≡ C) = 121 pm Edis = 839 kJ/mol SHAPES OF MOLECULES (VSEPR MODEL) Valence Shell Electron-Pair Repulsion model - Electron pairs surrounding atom spread out as to minimize repulsion. - Electron pairs can be bonding pairs or nonbonding pairs (multiple bonds not included). - arrangement of all the atoms surrounding central atom depends on electron pairs

surrounding central atom. Two similar, but different geometries 1. Electron set geometry - arrangement of e- sets around central atom - remember: e- set from multiple bonds is only one set. 2. Molecular geometry - arrangement of atoms around central atom **A molecular geometry is decided only after an electron set geometry has been determined.** *- need to write Lewis structure to determine number of electron sets.*

10Geometries with two e- sets about central atom. 1. electron set geometry – linear - angle between e- pairs is 180 ° 2. possible molecular geometries a) Linear - only linear geometry is possible with two electron pairs Example: BeCl2 Example: CO2 - Note: Only two electron sets around central atom since double bonds only count as one. Geometries with three e- sets about central atom. 1. electron sets geometry – trigonal planar - angle between e- pairs is 120 ° 2. possible molecular geometries a) Trigonal Planar - all three electron sets are bonding pairs Examples: BF3 and NO3

- b) Bent (V-shaped) - two bonding pairs and one nonbonding pair Example: SO2 *Nonbonding e- pairs take up more room than bonding pairs. Therefore bond angle between

oxygen atoms is slightly less than 120°.*

A A – generic atom

Be Cl Cl

C O O

A

N

O

O

O B

F

F F

A – generic atom

S O O

11Geometries with four e- sets about central atom. 1. electron set geometry – tetrahedral - tetrahedron is three dimensional object - angle between electron pairs is 109.4° 2. possible molecular geometries a.) Tetrahedral - all four electron sets are bonding pairs

Example: CH4 Example: PO4

3- b) Trigonal Pyramidal - 3 bonding pairs and 1 nonbonding pair Example: NH3 - Bond angle is 107°. - Bond angle is less than 109.4° because nonbonding pair takes more room than

bonding pair.

A Caution! Representation is 2-D, not 3-D, bond angle between e- pairs is not 90°

H C H

H

H

2-D picture 3-D picture

CHH

H H

P O

O

O

O

PO O

O O

N H

H

H NH HH

12 Example: ClO3

-

c) Bent (V-shaped) - two bonding pairs and two nonbonding pairs Example: H2O - Bond angle is 104.5°. Redraw H2O to show tetrahedral angle. Example: SF2

F F S

O H

H

Cl O

O

O

O

H H

ClO OO

13ELECTRONEGATIVITY In covalent bonding, sharing of electrons is rarely perfect. One atom will draw electrons closer to itself than the other atom. Consider Methanol

→H OC

H

H

H

H OC

H

H

H

= H OC

H

H

H

- in the C—O bond, oxygen draws e- to itself Electronegativity – ability of an atom in a molecule to attract electrons to itself. - symbol is χ (Greek chi) Periodicity of Electronegativity POLAR COVALENT BONDS - bond between two atoms where electrons are imperfectly shared - more electronegative atom has a greater electron density about it - bond has poles, regions that are have more positive or negative charge - more electronegative atom is slightly negative relative to less electronegative atom

δ – small change in charge density (lower case Greek delta)

perfect sharing imperfect sharing

C O C O

increasing

C O

δ+ δ-

14ELECTRONEGATIVITY AND BONDING When the electronegativities of the two atoms involved in a bond are known, the type of bonding can be predicted. Consider two atoms, A and B.

Atom A Atom B Bonding Less than 2.20 Less than 2.20 metallic Greater or = 2.20 Less than 1.70 ionic Greater or = 2.20 Greater or = 1.70 covalent

Examples 2.1 H 1.0 Li

1.5 Be

2.0B

2.5C

3.0N

3.5O

4.0F

1. Characterize the bonding of C2F4 χC = 2.5 χF = 4.0 ⇒ Bonding is covalent

- Both nonmetal atoms want the electrons, so the compromise is a covalent bond 2. Characterize the bonding of Li2O2 χLi = 1.0 χO = 3.5 ⇒ Bonding is ionic

- Metal atom doesn’t want electrons and nonmetal atom want electron, so ionic bonding is a perfect result

3. Characterize the bonding of BeH2 χBe = 1.5 χH = 2.1 ⇒ Bonding is metallic

- Both metals atoms don’t want the electrons, so the valence electrons become delocalized in a metallic bond

C2F4 is used in Teflon for non-stick cookware.

Li2O2 is used in subs to convert CO2 to O2.

BeH2 forms polymer [long chain] rather than molecule.

Is one ofthe atomsgreater orequal to 2.20?

Is theother atomgreater orequal to 1.70?

Bonding is metallic

Bonding is covalent

No

No

Yes

Yes

Bonding is ionic

15POLAR MOLECULES A polar molecule has one side slightly positive and the other slightly negative. Two conditions must be met in a polar molecule. 1.) Polar covalent bonds 2.) Correct geometry To emphasize necessity of correct geometry, compare two examples. GeO2 H2O χGe = 2.0 χH = 2.1 χCl = 3.5 χO = 3.5 |χGe – χO| = 1.5 |χH – χO| = 1.4 δ+ δ- δ+ δ- Ge – O H – O Ge – O bond is more polar than H – O bond; however, GeO2 is nonpolar molecule and H2O is a polar molecule. Question: How can this be? Answer: GeO2 has a linear geometry and H2O has a bent geometry. Example: Is either ammonia or methane a polar molecule? Answer: Ammonia is a polar molecule, but methane is a nonpolar molecule. Scheme: Chemical formula → Lewis structure → e- set geometry → molecular geometry

→ polarity

indicates positive end of bond

total polarity is nonzero

O

H H δ+

δ-

NH HH

δ+

δ-

CHH

H H

Ge O O

total polarity adds to zero