Bonding Exam

Ionic Bonds

• Metal + nonmetal 

• Soluble in polar liquids

• Hard but brittle 

• High boiling/melting point (ex: salt) 

Ionic Bonds

• Anions and cations are attracted to each other

• Bond formed through electron transfer 

• Called salts → simplest ratio is the formula unit 

Covalent

Covalent 

• Nonmetal + nonmetal share electrons

• Not soluble in water → only in nonpolar 

• Weak and soft 

• Low boiling/melting points (ex: CO2)

Giant Molecular

• Nonmetal + metal 

• Not soluble

• Hard and strong (ex: diomond)

• Very high melting/boiling points 

• Insulator (except graphite)

Metallic

• Metal + metal

• Not soluble 

• Hard and strong 

• Very high melting/boiling points

• Conduct electricity as liquids and solids

Electrons

• Valence electrons: those in the outermost energy level → responsible for chemical properties

• Core electrons: those in the energy levels below the valence shell

Dot diagrams

• Dots around an element symbol 

  • Around then double up 

Bonding

• A single bond is one line and one pair of electrons shared between two atoms

• A double bond is 2 lines and 2 pairs of electrons shared between two atoms

  • H, Cl, F, Br, and I rarely form double bonds

• A triple bond is 3 lines and 3 pairs of electrons shared between two atoms

  • C, O, and N often form double/triple bonds

σ and 𝜋 bonds 

• Single bonds have one σ bond 

• Double bonds have one σ bond and one 𝜋 bond 

• A triple bond has one σ and two 𝜋 bonds 

  • As the bond increases, the distance becomes shorter, and the bond is stronger

    • 𝜋 bonds pull atoms closer and make the bond stronger

Molecular structure 

• The central atom will be the least electronegative 

• Hydrogen will not typically be a central atom 

  • Electronegativity increases from left to right 

• The formula may also indicate structure 

Diagraming molecular compounds

• Count up the valence electrons contributed by each atom

• Draw basic structure and bonds

  • Subtract 2 e- for each bond

  • Distribute additional e- around central and peripheral atoms 

• If the octet/duet rule is not met, try double bonds

VSEPR

valence shell electron pair repulsion theory 

• The repulsion between electron pairs causes molecular geometries 

• Both bonded e- and “lone pairs” repel

Linear

• 2 elements surrounding 

• 0 lone pairs 

• 2 domains 

  • 180*

Trigonal Planar 

• 3 elements surrounding 

• 0 lone pairs

• 3 domains 

  • 120* 

Bent 

• 2 elements surrounding 

• 1 lone pair 

• 3 domains 

  • 117.5*

Tetrahedral 

• 4 elements surrounding 

• 0 lone pairs

• 4 domains 

  • 109.5*

Trigonal pyramidal 

• 3 elements surrounding 

• 1 lone pair 

• 4 domains 

  • 107* → subtract 2.5 for each lone pair from the parent shape (same domain)

Bent 

• 2 elements surrounding 

• 2 lone pairs 

• 4 domains 

  • 104.5* → subtract 5 from tetrahedral 

Formula Charge 

• Valence electrons, you start with, then subtract the dots and lines added together 

  • Preferred lewis structure is the one with the formula charges closest to 0

Resonance structure

• There is more than one Lewis structure for many molecules → only differ in double bond positions 

  • Different structures, called resonance structures

• Structure is described as a hybrid of the individual resonance structures 

• Shown by the double-headed arrows ←→

  • Ex: O3 has 2 resonance structures 

    • One structure ←→ second structure 

• For polyatomics 

  • Draw model

  • Then put in brackets and put the ionic charge outside 

    • Ex: [structure of CO3]-2

  • Then, to show all three resin structures: 

    •  [CO3]-2 ←→ [CO3]-2 ←→ [CO3]-2