Covalent Bonding Chapter 9

What do the following have in common? Oil and Vinegar They are covalent compounds.

Review Ionic bonds are formed by a transfer of e- Metals + Nonmetals  Ionic Compound FU = Formula unit smallest part of an ionic compound

Apply the octet rule to atoms that bond covalently Octet Rule: Atoms will gain e-, lose e- or share e- in order to get 8 valence e- to be stable

Covalent Bond Sharing of e- to be stable Usually occurs when elements are close on periodic table NM + NM  covalent compound aka: molecular compound

molecule 2 or more atoms combine covalently Ex: carbohydrates, proteins, DNA, fibers

Describe the formation of single, double and triple covalent bonds Single bond –formed because only 1 e- pair is shared between 2 elements These elements need only one additional e- to be stable –Ex: hydrogen and the halogens (Group 7A) H·+·H  H– H  H 2

Double Bonds Form because each element needs 2 e- pairs ·:O:·+ ·:O:·  ::O=O::  O 2

Triple Bonds Formed because 3 e- pairs are shared between two elements :N:· + :N:·  :N Ξ N:  N 2

Diatomic molecule 2 atoms that bond together Featuring “H and the Sensational 7”

Covalent bonds are formed by diatomic molecules Different nonmetals can share e- to be stable and form covalent bonds By drawing e- dot structures of each atom, you can put them together so each has 8 valence e- making a LEWIS STRUCTURE

Draw Lewis Structures PH 3 H 2 S CO 2

Draw Lewis Structures PH 3 ¨ H—P—H I H H 2 S CO 2

Draw Lewis Structures PH 3 H 2 S ֵ H—S—H ¨ CO 2

Draw Lewis Structures PH 3 H 2 S CO 2 :O=C=O: ¨ ¨

Relate the strength of covalent bonds to bond length The more bonds located between 2 atoms, the shorter the bonds are The shorter a bond is, the stronger it is H – H single bond, not too strong O=O double bonds, stronger NΞN triple bonds, strongest

Review 9.1 When 2 Nonmetals form a compound, they SHARE e- This is a covalent bond The compound formed is a molecular compound The smallest part of a molecular compound is a molecule

Bond TypeIonic BondCovalent Bond ElementsM + NM Polyatomic Ion NM + NM Atoms become stable by Gain/Lose e-SHARE e- Name of Compound Ionic compoundMolecular Compound Smallest Particle Formula Unit (FU) Molecule

9.2 Naming molecules Identify the names of binary molecular compounds from their formulas Molecules are formed when nonmetallic atoms share e- They can combine in different ratios, such as CO and CO 2 ; therefore, PREFIXES are used in the name

PREFIXES KNOW the prefixes on table 9-1 (p. 248) 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10 deca-

Using prefixes The prefix “mono” is NEVER used for the first element, but all prefixes are When the element begins with a vowel (O, I), you drop the a or o from the prefix The second element always has a prefix and ends in “-ide”

Examples CCl 4 Carbon tetrachloride As 2 S 3 Diarsenic trisulfide CO Carbon monoxide CO 2 Carbon dioxide

Writing Formulas When writing formulas for molecular compounds, use prefixes to tell how many atoms are in the molecule Sulfur dioxide SO 2 Diphosphorous pentoxide P 2 O 5

Name acidic solutions Acid formulas begin with an H H becomes hydro- The acid name comes from the second element or polyatomic ion name

Naming acid 2 nd name ends in -ide HCl (chloride) HF (fluoride) -ate -ite Acid name Hydro ____ic acid Hydrochloric acid Hydrofluoric acid ___ic acid ___ous acid

Naming acid 2 nd name ends in -ide -ate H 2 SO 4 (sulfate) HNO 3 (nitrate) -ite Acid name Hydro ____ic acid ___ic acid sulfuric acid nitric acid ___ous acid

Naming acid 2 nd name ends in -ide -ate -ite HClO 3 (chlorite) H 3 PO 3 (phosphite) Acid name Hydro ____ic acid ___ic acid ___ous acid Chlorous acid Phosphorous acid

Naming acid 2 nd name ends in -ide HCl (chloride) HF (fluoride) -ate H 2 SO 4 (sulfate) HNO 3 (nitrate) -ite HClO 2 (chlorite) H 3 PO 3 (phosphite) Acid name Hydro ____ic acid Hydrochloric acid Hydrofluoric acid ___ic acid sulfuric acid nitric acid ___ous acid Chlorous acid Phosphorous acid

Writing Acid Formulas The first element in an acid is ALWAYS H + To find the second, use the same chart, work backwards: -idehydro__ic acid -ate___ic acid -ite____ous acid

Example: Phosphoric acid ___ ic acid means the second name ends in –ate Phosphate is the second name Hydrogen phosphate (now use charges and criss cross) H + PO4 3- H 3 PO 4

9.3 Molecular Structure Structural formulas = uses letter symbols and bonds to show relative position of atoms (same as Lewis Structure) Chem I: –Don’t need definition of resonance or coordinate covalent bond –Only need step 1 of drawing Lewis structures –Don’t need obj. 2, 3, of 9.3

Rules *1. Predict location of certain atoms a) Hydrogen is always terminal (end) atom b) Atom with least e- affinity (furthest to left on periodic table) is central atom

9.4 Molecular Shape VSEPR Valence Shell Electron Pair Repulsion Theory States: Repulsion of shared and unshared pairs (lone pairs) of e- around the central atom shapes the molecule

Table 9-3 Look at a correctly drawn structural formula Count the number of shared pairs and lone pairs on the CENTRAL ATOM Compare to the table Examine table 9-3 examples on p. 260

Atoms in Molecule Lone pairs on central atom Molecule Shape 2Linear 3NoneLinear 3Lone pairBent 4NoneTrigonal planar 4Lone pairTrigonal pyramidal 5tetrahedral

9.5 Electronegativity & Polarity Electron Affinity is the ability of an atom to accept an e- Excluding noble gases, e- affinity increases to the right of the periodic table Electron affinity increases going up in a group

Electronegativity Assigned values for elements that compare the ability of an atom to attract shared e- to itself to the ability of Fl to do the same Has the same trend as e- affinity

Bond Polarity To determine if a bond is polar or not: (polar means having a negative and positive side) –Compare the electronegativity of each atom connected by the bond –The atom that is further to the right (or in the same group: further up) will hold the e- closer and be σ- –The other atom will be σ+

Examples H 2 O CH 3 Cl An atom may be + with respect to one bond and – with respect to another in the same molecule

More info Electronegativity Difference Bond Type 0Nonpolar Covalent 0.1 – 1.7Polar Covalent Above 1.7Ionic Bond

Molecule Polarity To determine if the entire molecule is polar (has a negative and positive side) or not, you must look at 2 things 1. molecule shape 2. bond polarity If the molecule contains only nonpolar bonds, then the molecule is also NONPOLAR

If the molecule has polar bonds, it may be a NONPOLAR or a POLAR molecule, depending on the shape Nonpolar molecules are usually linear (with same charge on all sides); tetrahedral (with same charges on all sides); planar (with same charge on all sides)

Polar molecules are bent or pyramidal (with opposite charges on sides) Polar molecules can be linear, tetrahedral, or planar (with opposite charges on sides)

Properties of Molecular Compounds REVIEW: Ionic compounds were crystalline solids at room temp with high melting points and high boiling points Molecular compounds may be gases, liquids or solids at room temp Molecular compounds have LOW melting points and LOW boiling points

Comparison IonicMolecular State at room temp Crystalline solidGas, solid, liquid Boiling PointHighLow Melting PointHighLow