Covalent Bonding

Section 4.2

Covalent Bonds

• Sharing Electrons

– Covalent bonds form when atoms share one or more pairs of electrons

•nucleus of each atom is attracted to electron cloud of other atom

•neither atom removes an electron from the other

Covalent Bonding

Covalent Bonds

• space where electrons move is called molecular orbital

• made when atomic orbitals overlap

Molecules (Space-filling Models)

Covalent Bonds

• Energy and Stability

– Noble gases are stable (full octet) (low P.E.)

– Other elements are not stable (high P.E.)

•covalent bonding decreases potential energy because each atom achieves an electron configuration like a noble gas

Covalent Bonds

• Energy and Stability

•because P.E. decreases when atoms bond, energy is released

–atoms lose P.E. when they bond

–loss of P.E. implies higher stability

Covalent Bonds

• Energy and Stability

• potential energy determines bond length

• at minimum P.E., distance between two bonded atoms' nuclei is called bond length

– bonded atoms vibrate

– therefore, bond length is an average length

Covalent Bonds

• bonds vary in strength

• bond energy is the amount of energy required to break the bonds in 1 mol of a chemical compound

• bond energy predicts reactivity

• bond energy is equal to loss of P.E. during formation

Multiple Bonds

• Consider 2 carbon atoms joined by a single, double, or triple bond

Structure Bond length (pm)

Bond strength (kJ mol-1

C * C 154 364

C = C 134 602

C ≡ C 120 835

Multiple Bonds• The shorter the bond, the greater the bond

strength.

• The more bonds between two covalently bonded atoms, the shorter the bond

• Length:

• triple bond < double bond < single bond

• Strength:

• single bond < double bond < triple bond

Another Example

• Consider the COOH group which is attached to certain acids

• The C is single bonded to one O and double bonded to the other

Structure Bond length (pm)

Bond strength (kJ mol-1

C * O 143 358

C = O 120 799

Covalent Bonds• Electronegativity

– Atoms share electrons equally or unequally

•nonpolar covalent bond: bonding electrons shared equally (when diff in EN value is zero)

•polar covalent bond: shared electrons more likely to be found around more electronegative atom (when diff. in EN value is between 0 and less than 1.8)

Covalent Bonds

• Differences in electronegativity can be used to predict type of bond (but boundaries are arbitrary)

• Subtract the electronegativities of the atoms in the bond to determine the difference

Bond Types

Practice: Calculate the bond type

• N and H

• F and F

• Ca and Cl

• C and O

Polar

Non-polar

Ionic

Polar

Covalent Bonds

• Polar molecules have positive and negative ends (poles)

•such molecules are called dipoles

•δ (“delta”) means “partial” in math and science

•positive end is designated as δ+

•negative end is designated asδ-

•example: Hδ+Fδ-

Drawing and Naming• Lewis Electron-Dot Structures

– Lewis structures represent valence electrons with dots

•position of electrons is symbolic (not literal)

•shows only the valence electrons of an atom

•dots around the atomic symbol represent electrons

Lewis Structures of Second-Period Elements

Drawing Lewis Structures for Compounds

Refer to extra notes on moodle

Coordinate Covalent (Dative) Bonds

• resulting bond when a shared pair of electrons originates from one atom

• Examples: CO (O contributes 2 electrons for one of the bonds), NH

4+ (N contributes 2), H

3O+ (O

contributes 2)

Drawing and Naming

• Polyatomic Ions: be sure to add or subtract the electrons

• use brackets [] to show overall charge

• Examples: NH4

+ and OH-

Naming Covalent Compounds

MOLECULAR GEOMETRY

VSEPR • Valence Shell Electron Pair Repulsion

theory.• Most important factor in determining

geometry is relative repulsion between electron pairs.

Molecules adopt the shape that

minimizes the electron pair repulsions.

Molecular Shapes

• Lewis structures show which atoms are connected where, and by how many bonds, but they don't properly show the 3-D shape of the molecules

• To find the actual shape, first draw the Lewis structure, and then apply the VESPR Theory

Molecular Shapes

• Determining Molecular Shapes

– Three-dimensional shape helps determine physical and chemical properties

– valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes• based on the idea that electrons repel one

another

VSEPR Rules• To apply VSEPR theory:

• 1: Draw the Lewis structure of the molecule and identify the central atom

• 2: Count the number of negative centers (lone and bonding pairs) surrounding the central atom.

• 3: Predict molecular shape by assuming that electrons orient themselves so they are as far away from one another as possible

VSEPR Shape Predictor Table

VSEPR Shape Predictor Table

Bond Angles

• Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected:

Bond Angles

• Methane, CH4, has a

perfect tetrahedral bond angle of about 109°, while the N-H bond angle of ammonia, NH

3, is

slightly less at 107°:

Bond Angles• The oxygen of water has two bonded electron pairs

and two non-bonded "lone" electron pairs. The lone pairs push the bond down so it is smaller

• The geometry is defined by the relationship between the H-O-H atoms and water is said to be "bent" or "angular" shape of 104.5°.

Shape Bond Angle Example

linear 180 CO2 and HCN

Trigonal planar 120 BF3 and SO3

V-shaped 117 SO2

Tetrahedral 109 CH4 and CCl4

Trigonal pyramidal

107 NH3 and NF3

Bent linear 104.5 H2O and H2S

Table of Common Shapes, Angles and Examples

Simple Shapes

Trigonal Planar

Tetrahedral

Bent

Molecular Shapes

• Determining Molecular Shapes and angles.

– Let’s try some.• CO

• NCl3

• BF3

• CH4

• H2O

• SO2

Polarity of Molecules

• Two atoms: bond polarity is the molecular polarity

• More than 2 atoms: the geometry of the molecule must be considered

• If the bonds are non-polar, the molecular is non-polar

• Some molecules with polar bonds can be non-polar

More

• Sometimes the partial charges cancel each other out because they are directly opposite each other

• Consider CO2 and CCl

4

• The symmetrical distribution of the bonds leads to cancellation of the charges

Polar Bonds

Giant Covalent Lattices

• Consist of a 3-D lattice of covalently bonded atoms

• Atoms can be either all of the same type as in silicon and carbon, or of two different elements such as silicon and dioxide

Allotropes of Carbon

• Allotropes are two or more crystalline forms of the same element in which the atoms are bonded differently

• Carbon has 3 allotropes

– Diamond, graphite, C60 (fullerene)

Diamond

• Each C atom is tetrahedrally bonded to 4 other C atoms by single covalent bonds (giant covalent lattice)

• Very rigid 3-D network with bond angles of 109.5°

• Very hard, poor conductor, very dense

Graphite• Each C atom is covalently bonded to 3

other C atoms (giant covalent lattice)

• 2-D network is formed consisting of hexagonal rings of carbon atoms

• Crystal structure is composed of many layers of rings

• Very soft and slippery, good

conductor along the plane of the layers, not as dense as diamond

Fullerene

• C atoms are arranged into the shape of a soccer ball

• Has 60 corners and 32 faces

• Bonding is a series of single and double bonds

• Soft, very poor conductor, less dense than graphite

Diamond Graphite Fullerene

Silicon Dioxide

• Exists as a giant covalent structure

• Most common form is quartz

• Structure is similar to diamond

• Hard, transparent, high melting and boiling point

• Impure form of silicon dioxide is sand

Silicon Dioxide

Silicon

• Has a giant covalent structure similar to diamond

• Less hard than diamond owing to the larger size of the silicon atoms (longer and weaker bonds)

• Is an insulator, but can be made to conduct electricity by adding small amounts of other atoms

Silicon structure is like the diamond structure