Chapter 5 Molecular Compounds

Chemical Bonds Three basic types of bonds: Ionic Covalent Metallic Electrostatic attraction between ions Covalent Sharing of electrons Metallic Metal atoms bonded to several other atoms

Attraction Between Two or More Ions Ionic Bonding Attraction Between Two or More Ions

The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6 [Ne] Li Li+ + e- e- + F - F - Li+ + Li+ 9.2

Sharing electrons within a Molecule! Covalent Bonding Sharing electrons within a Molecule!

Valence Electrons Valence electrons are the outer shell s and p electrons of an atom. The valence electrons are the electrons that participate in chemical bonding. We show the valence electrons in Lewis dot structures of atoms and molecules.

9.1

Molecular Compounds Molecular compound: a compound in which all bonds are covalent Naming binary molecular compounds the less electronegative element is named first prefixes “di-”, tri-”, etc. are used to show the number of atoms of each element; the prefix “mono-” is omitted when it refers to the first atom, and is rarely used with the second atom. Exception: carbon monoxide NO is nitrogen monoxide (nitric oxide) SF2 is sulfur difluoride N2O is dinitrogen monoxide (laughing gas)

Forming a Covalent Bond A covalent bond is formed by sharing one or more pairs of electrons the pair of electrons is shared by both atoms and, at the same time, fills the valence shell of each atom example: in forming H2, each hydrogen contributes one electron to the single bond

Why do two atoms share electrons? A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why do two atoms share electrons? 7e- 7e- 8e- 8e- F F + F Lewis structure of F2 lone pairs F single covalent bond single covalent bond F 9.4

Lewis structure of water single covalent bonds 2e- 8e- 2e- H + O + H O H or Double bond – two atoms share two pairs of electrons 8e- 8e- 8e- double bonds O C or O C double bonds Triple bond – two atoms share three pairs of electrons 8e- triple bond N 8e- or N triple bond 9.4

Lengths of Covalent Bonds Bond Type Bond Length (pm) C-C 154 CC 133 CC 120 C-N 143 CN 138 CN 116 Bond Lengths Triple < Double < Single 9.4

Comparison of Ionic and Covalent Compounds 9.4

Drawing Lewis Structures 1. Set aside the single bonders (H, F, Cl, Br, I). 2. Put together the 2, 3, and 4 bonders with single bonds (May require thinking). 3. Count the holes left A. If holes = single bonders, put them on. B. If holes > single bonders, make double or triple bonds, or a ring until A is satisfied.

Lewis Structures Examples draw a Lewis structure for hydrogen peroxide, H2O2 draw a Lewis structure for methanol, CH3OH draw a Lewis structure for acetic acid, CH3COOH

Lewis Structures Draw a Lewis Structure and Make a model for C2H6O BF3 HNO3 HCN C2H3Cl

Lewis Structures You will build models and draw Lewis Structures in Lab. You will have to identify the polarity in a molecule as well. Link to polarity idea

Polar Covalent Bonds Polar covalent bond or polar bond is a covalent bond with uneven electron density electron rich region electron poor region H F e- poor e- rich F H d+ d- 9.5

Electronegativity Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electronegativity - relative, F is highest 9.5

Electronegativity 9.5

Increasing difference in electronegativity Classification of bonds Difference Bond Type Covalent  2 Ionic 0.4 < and <2 Polar Covalent Increasing difference in electronegativity Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e- 9.5

Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond.

Solve The Following Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the CN bond in HCN; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic Polar Covalent C – 2.5 N – 3.0 3.0 – 2.5 = 0.5 N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent 9.5

Write the Lewis structure of nitrogen trifluoride (NF3). 9.6

Write the Lewis structure of the carbonate ion (CO32-). 9.6

Formal Charge and Lewis Structures Lewis structures with large formal charges are less plausible than those with small formal charges. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O 9.7

Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule: Ions or molecules with an odd number of electrons. Ions or molecules with less than an octet. Ions or molecules with more than eight valence electrons (an expanded octet).

Exceptions to the Octet Rule The Incomplete Octet BeH2 H Be F B BF3 9.9

Exceptions to the Octet Rule Odd-Electron Molecules NO N O The Expanded Octet (central atom with principal quantum number n > 2) S F SF6 9.9