Covalent Bonding

Mr. SolsmanChapter 8

• Apply the octet rule to atoms that form covalent bonds.

• Describe the formation of single, double, and triple covalent bonds.

• Contrast sigma and pi bonds.

• Relate the strength of a covalent bond to its bond length and bond dissociation energy.

• Atoms gain stability when they share electrons and form covalent bonds.

• Lower energy states make an atom more stable.• Gaining or losing electrons makes atoms more

stable by forming ions with noble-gas electron configurations.

• Sharing valence electrons with other atoms also results in noble-gas electron configurations.

• Atoms in non-ionic compounds share electrons.

• The chemical bond that results from sharing electrons is a covalent bond.

• A molecule is formed when two or more atoms bond.

• Diatomic molecules (H2, F2 for example) exist because two-atom molecules are more stable than single atoms.

The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

• When only one pair of electrons is shared, the result is a single covalent bond.

• The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in an electron configuration like helium.

• In a Lewis structure dots or a line are used to symbolize a single covalent bond.

• The halogens—the group 17 elements—have 7 valence electrons and form single covalent bonds with atoms of other non-metals.

• Atoms in group 16 can share two electrons and form two covalent bonds.

• Water is formed from one oxygen with two hydrogen atoms covalently bonded to it .

• Atoms in group 15 form three single covalent bonds, such as in ammonia.

• Atoms of group 14 elements form four single covalent bonds, such as in methane.

• Sigma bonds are single covalent bonds.• Sigma bonds occur when the pair of shared

electrons is in an area centered between the two atoms.

• Double bonds form when two pairs of electrons are shared between two atoms.

• Triple bonds form when three pairs of electrons are shared between two atoms.

• A multiple covalent bond consists of one sigma bond. The pi bond is formed when parallel orbitals overlap and share electrons.

•

• The strength depends on the distance between the two nuclei, or bond length.

• As length increases, strength decreases.

• The amount of energy required to break a bond is called the bond disassociation energy. The shorter the bond length, the greater the energy required to break it.

• An endothermic reaction is one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products.

• An exothermic reaction is one where more energy is released than is required to break the bonds in the initial reactants.

Naming Binary Molecular Compounds

• The first element is always named first using the entire element name.

• The second element is named using its root and adding the suffix –ide.

• Prefixes are used to indicate the number of atoms of each element in a compound.

• Name each of the covalent compounds below:• CO2

• SO2

• NF3

• CCl4

• N2O4

• Write the formula for diarsenic trioxide

Naming Acids

• The first word has the prefix hydro- followed by the root of the element plus the suffix –ic.

• The second word is always acid (hydrochloric acid is HCl in water).

• HBr• HI• HF• HCN

• An oxyacid is an acid that contains both a hydrogen atom and an oxyanion.

• Identify the oxyanion present.• The first word is the root of the oxyanion and

the prefix per- or hypo- if it is part of the name, plus the suffix -ic if the anion ends in -ate or -ous if the oxyanion ends in -ite.

• Name the following oxyacids:• HClO• HClO2

• HClO3

• HClO4

Structural Formulas

• A structural formula uses letter symbols and bonds to show relative positions of atoms.

Lewis Structures

– Predict the location of certain atoms.– Determine the number of electrons available

for bonding.– Determine the number of bonding pairs.– Place the bonding pairs.– Determine the number of bonding pairs

remaining.– Determine whether the central atom satisfies

the octet rule.

• Determine and draw the Lewis structure:• LiH

•

• BeH2

• BH3

• CH4

• NH3

• OH2

• FH

• CO2

• N2

• NO3

Resonance Structures

• Two or more correct Lewis structures that represent a single ion or molecule are resonance structures.

• The molecule behaves as though it has only one structure.

• The bond lengths are identical to each other and intermediate between single and double covalent bonds.

• Try drawing resonance structures for:• SO2

• NO2-

• O3

• SO32-

Exceptions to the Octet Rule

• Some molecules do not obey the octet rule.• A small group of molecules might have an odd

number of valence electrons.• NO2 has five valence electrons from nitrogen

and 12 from oxygen and cannot form an exact number of electron pairs.

• A few compounds form stable configurations with less than 8 electrons around the atom—a suboctet.

• A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.

• A third group of compounds has central atoms with more than eight valence electrons, called an expanded octet.

• Elements in period 3 or higher have a d-orbital and can form more than four covalent bonds.

• Draw the Lewis structure for:• ClF3

• XeF4

• SF6

Shape: VSEPR Model

• The shape of a molecule determines many of its physical and chemical properties.

• Molecular geometry (shape) can be determined with the Valence Shell Electron Pair Repulsion model, or VSEPR model which minimizes the repulsion of shared and unshared atoms around the central atom.

• Electron pairs repel each other and cause molecules to be in fixed positions relative to each other.

• Unshared electron pairs also determine the shape of a molecule.

• Electron pairs are located in a molecule as far apart as they can be.

HOW????

• Hybridization is a process in which atomic orbitals mix and form new, identical hybrid orbitals.

• Carbon often undergoes hybridization, which forms an sp3 orbital formed from one s orbital and three p orbitals.

• Lone pairs also occupy hybrid orbitals.

Electron Sites

• A pair of electrons in a sigma bond count as one electron site.

• A non-bonding (unshared) pair of electrons count as one electron site.

• Electrons in a π – bond do NOT count.

Electron Sites

• Sigma's + unshared pairs = 4 Then sp3

• Sigma's + unshared pairs = 3 Then sp2

• Sigma's + unshared pairs = 2 Then sp1

• Sigma's + unshared pairs = 1 Then s

Geometries

• s or sp = linear• sp2 = planar triangular• sp3 = tetrahedral

Shapes

• If there are only two atoms the shape is always linear.

• If there are NO NON-BONDING pairs, the shape is always the same as the geometry.

Shapes

• sp2 ( no non-bonding pairs) = planar triangle• sp2 ( 1 non-bonding pair) = bent• sp2 ( 2 non-bonding pairs) = linear• sp3 (no non-bonding pairs) = tetrahedral• sp3 (1 non-bonding pair) = pyramidal• sp3 (2 non-bonding pairs) = bent• sp3 (3 non-bonding pairs) = linear

• Predict the geometry and shape of the following:

• CS2

• CH2O

• H2Se

• CCl2F2

• NCl3

• BF3

• OCl2

• BeF2

• CF4

• NH4+

• PH3

• CH2Cl2

• C2H4

• C2H2F2

• (CH3)2CO

ELECTRONEGATIVITY

A measure of the attraction an atom has for a pair of electrons in a chemical bond.

a. Electronegativity is a combination of ionization energy and electron affinity.

b. Increases across each period to a maximum in the 17th (or 18th) family.

c. Decreases down each family of elements.

Electronegativities of the Representative Elements

0.0

0.5

1.0

1.5

2.0

2.5

3.0

3.5

4.0

4.5

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 31 32 33 34 35 36 37 38 49 50 51 52 53 54 55 56 81 82 83 84 85 86 87 88

Atomic Number

Relat

ive E

lectro

nega

tivity

Percent of ionic character = the percent the electrons are transferred in a chemical bond.

If ∆ EN is 0; bonding is nonpolar covalent.If Δ EN is < 0.4; bonding is mostly covalent.If Δ EN is 0.4 – 1.7; bonding is polar covalentIf Δ EN is > 1.7 ;bonding is ionic.

• Unequal sharing of electrons results in a polar covalent bond.

• Bonding is often not clearly ionic or covalent.

Polar Covalent Bonds

• Polar covalent bonds form when atoms pull on electrons in a molecule unequally.

• Electrons spend more time around one atom than another resulting in partial charges at the ends of the bond called a dipole.

• Covalently bonded molecules are either polar or non-polar.

• Non-polar molecules are not attracted by an electric field.

• Polar molecules align with an electric field

Molecular Geometry and Shape

• The shape of a molecule helps determine the polarity of the molecule.

• The shape of the molecule is determined by the geometry of the electron pairs.