Chemical Bonding Compounds are formed from chemically bound atoms or ions. Bonding involves only the valence electrons.

Chemical Bonding Ionic Compounds Ionic Radii Lattice Energy Molecular Compounds Covalent bonds Bond Order Bond Strength Lewis Structures

Lewis Symbols Lewis symbols show the valence electrons as dots arranged around the atomic symbol. hydrogen: sodium: chlorine: H · Na · Cl · · ·

The Octet Rule Atoms tend to gain, lose, or share electrons until they have eight valence electrons.

Ionic and Molecular Compounds Formation of sodium chloride: Cl · · · ® Na+ [ ] Cl · · Na · + Formation of hydrogen chloride: Cl · · · ® Cl · · H H · + A metal and a nonmetal transfer electrons to form an ionic compound. Two nonmetals share electrons to form a molecular compound.

Ionic Compounds Ionic compounds consist of a lattice of positive and negative ions. NaCl:

Ionic Bonds An ionic bond is simply the electrostatic attraction between opposite charges. · Ions with charges Q1 and Q2: Q2 Q1 d d Q E 2 1 µ The potential energy is given by:

Ionic Sizes

The Lattice Energy The lattice energy is the enthalpy change required to separate one mole of an ionic compound into its ions: NaCl(s) ® Na+(g) + Cl-(g) DH = 788 kJ Lattice energy increases with increasing ionic charges and decreasing ionic sizes. KCl(s) ® K+(g) + Cl-(g) DH = 701 kJ MgCl2(s) ® Mg+2(g) + 2 Cl(g) DH = 3795 kJ

The Born-Haber Cycle The Lattice energy can be calculated with Hess’s law and the following steps: Na+(g) + e- + Cl(g) H I1(Na) E(Cl) Na+(g) + Cl-(g) Lattice energy Na(g) + Cl(g) DHf°(Cl,g) Na(g) + ½Cl2(g) DHf°(Na,g) Na(s) + ½Cl2(g) -DHf°(NaCl,s) NaCl(s)

Estimating Lattice Energy Arrange with increasing lattice energy: KCl NaF MgO KBr NaCl d Q E 2 1 µ 701 kJ 910 kJ · K+ Cl 3795 kJ d · K+ Br 671 kJ 788 kJ d

Molecular Compounds The simplest molecule is H2: Increased electron density draws nuclei together The pair of shared electrons constitutes a covalent bond.

Lewis Structures Covalent bonding in a molecule is repre-sented by a Lewis structure. A valid Lewis structure should have an octet for each atom except hydrogen. H · + H2: ® H H or H H Cl · · · + ® Cl Cl · · Bonding electrons Cl2: or Cl Cl · · Nonbonding electrons

Lewis Structures Draw Lewis structures for: or H F HF: H F H O H · · or H F · · HF: H F · · · · · · H O H · · or H O H · · H2O: H N H H · · or H N H H · · NH3: H C H H · · or H C H H CH4:

Double and Triple Bonds Atoms can share four electrons to form a double bond or six electrons to form a triple bond. = O O · · O2: · · N N  N2: The number of electron pairs is the bond order.

Electronegativity Polarity refers to a separation of positive and negative charge. In a nonpolar bond, the bonding electrons are shared equally: H2, Cl2: In a polar bond, electrons are shared unequally because of the difference in Zeff. HCl:

Electronegativity Electronegativity refers to the ability of an atom in a molecule to attract shared electrons. The Pauling scale of electro- negativity:

Bond Polarity H Cl A polar bond can be pictured using partial charges: +  H Cl  = 0.9 2.1 3.0 Electronegativity Difference Bond Type 0 - 0.5 Nonpolar 0.5 - 2.0 Polar 2.0  Ionic

Drawing Lewis Structures Sum the valence electrons from all atoms. Add one for each negative charge and subtract one for each positive charge. Draw a skeleton structure with atoms attached by single bonds. Complete the octets of atoms bound to the central atom. Place extra electrons on the central atom. If the central atom doesn’t have an octet, try forming multiple bonds.

Drawing Lewis Structures · · Cl C Cl O · · · · COCl2 24 ve’s · · · · · · · · · · · · · · · · HOCl 14 ve’s H O Cl · · · · · ·  · · O Cl O O · · · · · · · · ClO3 26 ve’s · · · · · · · · · · H C O H H · · CH3OH 14 ve’s · ·