Chemical BondsModern Chemistry: Chapter 6

Why?How?What?

Where?

Why?• Electromagnetism is 1 of the 4 universal forces• Balance between repulsion & attraction– Protons repel protons; electrons repel electrons– VSEPR = valence shell electron pair repulsion– Positive nuclear charges attract electrons

• Electrons are from self & nearby atom(s)

• Compounds are more stable than free atoms/ions– Lower potential energy at the optimal bond length– Energy is released upon bond formation, typically

• Heat &/or Light• Sound• Movement

Octet Rule answers “Why?”

• Bonds typically form to give 8 e- in outer shell.• This provides a “noble gas configuration.”• There are, however, exceptions.– BF3 & AlCl3 only have 6 electrons for B & Al

– PF5 & SF6 have more than 8 electrons for atoms

– H2 & He only have 2 electrons in outer shell

What?• Types of bonds– Ionic chemical bonds• Oppositely charged ions attract one another

– Covalent chemical bonds• Nonpolar = Equal sharing of electrons• Polar = Unequal sharing of electrons

– Metallic bonds• Mobile sea of electrons surround cations

How?• Chemical bonds are determined by differences

in the degrees of electronegativity (0 – 4.0)– Ionic = difference > 2.0• Alkali & alkaline earth metals with halogens/ nonmetals• Some transition metals with nonmetals

– Polar covalent = difference of 0.6 – 1.9• Nonmetals with one another• Some transition metals with nonmetals

– Nonpolar covalent = difference of 0 – 0.5• Diatoms (I, Br, Cl, F, O, N, H)• Nonmetals with very similar nonmetals

How else?• Molecular orbitals form around atoms• Size of atoms/ ions influence bond formation• Unshared valence electrons affect the shape

of the entire molecule (molecular geometry)• Bond length: minimum potential energy– Single bonds are longest• Two electrons are involved• Smallest bond energy (least repulsion)

– Triple bonds are shortest• Six electrons (3 pairs) are shared

What about metallic bonds?

• Vacant orbitals in outer energy levels overlap• Overlapping orbitals allow outer electrons to

roam freely throughout the entire metal– Malleability, Ductility, & Conductivity

• Many orbitals spaced by incremental energy levels allow absorption of many frequencies– Luster

Where?

• Covalent bonds form in the electron cloud– s bonds: symmetrical along nuclei’s axis– p bonds: side by side overlap of p orbitals

in sausage-shaped regions above & below axis• Ionic bonds form in the charged space of the

electron cloud• Metallic bonds form between cations of the

same element as a mobile sea of electrons

Electron-Dot Notation

• Illustrates only the valence electrons– Nucleus & inner-shell electrons = element symbol– Valence electrons shown as dots: E, N, W, & S

• Compounds shown as Lewis structures– Shared valence electrons = dot-pairs or dash(es)– Unshared valence electrons = dot(s)

• Structural formulas don’t show unshared pairsF-F, H-Cl, K-I, etc…

Resonance Structures

• A single representation is inadequate• Molecule may constantly alternate between

bonding structures• Molecule may form an average of 2 structures– Ozone (O3) forms identical O-O bonds that are

between a single & a double bond

Ionic Compounds• Formula unit = simplest collection of atoms• Crystal lattice = 3-D arrangement of ions– Cubic - Monoclinic -Triclinic– Tetragonal - Hexagonal– Orthorhombic - Rhombohedral

• Lattice energy (kJ/mol) = energy released upon crystal formation from gaseous ions

• Polyatomic Ions: NH4+, MnO4

-, SO4=, etc…

– Ions held together by covalent bonds.

VSEPR Theory

• Molecules form to lessen e- pair repulsion• Diatoms form linear (180o)• Group III/13 form trigonal-planar (120o)• Group IV/14 form tetrahedral (109.5o)• Group V/15 form trigonal-pyramidal (107o)• Group VI/16 form bent or angular (105o)• SF6 types form octahedral (90o)

Hybridization

• Atomic orbitals mix & form equal hybrid orbitals on the same atom– s & p orbitals sp orbital (180o)• BeF2

– s, p, & p orbitals sp2 orbital (120o)• BF3

– s, p, p, & p orbitals sp3 orbital (109.5o)• CCl4

Intermolecular Forces

• Dipole-dipole forces– Separated equal but opposite charges Dipole– Represented by arrow with head toward (-) pole• Forces of attraction between polar molecules

• Hydrogen bonding– H of 1 molecule pulled to (-) charge on another

• London dispersion forces– Attractions from creation of instantaneous dipoles