Chemical Bonds Modern 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.

– BF 3 & AlCl 3 only have 6 electrons for B & Al – PF 5 & SF 6 have more than 8 electrons for atoms – H 2 & 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 pairs  F-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 (O 3 ) 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 – Tetragonal – Orthorhombic - Monoclinic - Hexagonal - Rhombohedral -Triclinic Lattice energy (kJ/mol) = energy released upon crystal formation from gaseous ions Polyatomic Ions: NH 4 + , MnO 4 , SO 4 = , etc… – Ions held together by covalent bonds.

VSEPR Theory

• • • • • • • Molecules form to lessen e pair repulsion Diatoms  form linear (180 o ) Group III/13  form trigonal-planar (120 o ) Group IV/14  form tetrahedral (109.5

o ) Group V/15  form trigonal-pyramidal (107 o ) Group VI/16  form bent or angular (105 o ) SF 6 types  form octahedral (90 o )

Hybridization

• Atomic orbitals mix & form equal hybrid orbitals on the same atom – – – s & p orbitals  sp orbital (180 o ) • BeF 2 s, p, & p orbitals 

sp

2 orbital (120 o ) • BF 3 s, p, p, & p orbitals  • CCl 4

sp

3 orbital (109.5

o )

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