1. Lecture 37 Covalent Bonds Ozgur Unal 1

2.  What type of bond exist between the ions?  NaClMgCl2Ca3(PO4)2 2  Are the following compounds ionic compounds?  CO2  H2O  NH3  The bonds between CO2, H2O and NH3 are called covalent bonds.

3.  Ions lose or gain electrons to gain chemical stability (octet rule) forming ionic compounds.  Nonmetals share electrons with other nonmetals to gain chemical stability  Formation of covalent bonds  The chemical bond that results from sharing valence electrons is a covalent bond.  In a covalent bond, the shared electrons are considered to be part of the outer energy levels of both atoms involved.  A molecule is formed when 2 or more atoms bond covalently. 3

4.  Example: H2, N2, O2, F2, Br2, I2 when two atoms of each element share electrons.  These 2 atom molecules are chemically more stable than the individual atoms. Consider Fluorine: 1s 2 2s 2 2p 5  7 valence electrons  Check out Figure 8.2  Check out Figure 8.3  Two F atoms share 2 electrons (a pair of electrons)  The pair of electrons shared are called bonding pair.  The unshared valence electrons are called lone pairs. 4

5.  When only one pair of electrons is shared, it is a single covalent bond.  Example: F2, H2, Cl2  Check out Figure 8.4 for H2.  Lewis dot structure can be used to show the bonding pairs and lone pairs in a molecule. 5

6. Group 17 and single bonds  Halogens have 7 valence electrons  1 electron is needed to form an octet.  Group 17 elements form single covalent, bonds with atoms of other nonmetals.  They also form covalent bonds with identical atoms. 6 Group 16 and single bonds  Atoms of group 16 can share 2 electrons  Two covalent bonds (either single or double)  Example: H2O (water)

7. Group 15 and single bonds  Group 15 elements form 3 covalent bonds.  Example: NH3, NF3, NCl3 7 Group 14 and single bonds  Group 14 elements form 4 covalent bonds.  Example: CH4 (methane)  http://bcs.whfreeman.com/thelifewire/content/chp02/0202 0.html

8.  Example: The pattern on the glass shown in Figure 8.6 was made chemically etching its surface with hydrogen fluoride (HF). Draw the Lewis dot structure for a molecule of HF. 8  Example: Draw a generic Lewis dot structure for a molecule formed between atoms of Group 1 and Group 16 elements.

9.  Single covalent bonds are also called sigma bonds, represented by the Greek letter sigma (σ).  When 2 atoms share electrons, their valence atomic orbitals overlap end-to-end, concentrating the electrons in a bonding orbital between the 2 atoms.  Sigma bonds can form when an s orbital overlaps with another s orbital or a p orbital, or two p orbitals overlap.  Check out Figure 8.7 for examples 9

10. Lecture 38 Multiple Covalent Bonds Ozgur Unal 10

11.  Some nonmetals share more than one pair of electrons to have the electron configuration of a noble gas.  Sharing multiple pairs of electrons forms multiple covalent bonds.  A double covalent bond and a triple covalent bond are examples of multiple covalent bonds.  Example: C, O, N and S can form multiple covalent bonds. 11

12. Double bonds:  A double covalent bond forms when two pairs of electrons are shared between two atoms.  Example: O2 12 Triple bonds:  A triple covalent bond forms when three pairs of electrons are shared between two atoms.  Example: N2

13.  A multiple covalent bond consists of one sigma bond and at least of one pi bond.  A pi bond, represented by the Greek letter pi (π), forms when parallel orbitals overlap and share electrons.  The shared electron pair of a pi bond occupies the space above and below the line that represents where the two atoms are joined together.  Check out Figure 8.9 13

14.  In a molecule, nuclei and electrons attract each other, electrons repel each other and nuclei also repel each other.  When this balance of forces is upset, the covalent bond may be broken.  Some covalent bonds are easy to break (weak), some are hard to break (strong). 14

15. 15  The strength of a bond depends on the bond length – the distance between the two bonded nuclei at the position of maximum attraction.  Bond length is determined by the sizes of the two bonding atoms and the number of electron pairs shared.  The shorter the bond length, the stronger the bond strength.  Check out Table 8.1

16.  Forming bonds release energy, breaking bonds require energy.  The amount of energy required to break a specific covalent bond is called bond-dissociation energy  Always positive.  The smaller the bond length, the greater the bond dissociation energy. 16

17. 17  A chemical reaction involves reactants and products.  The total energy change in a chemical reaction is determined from the energy of the bonds broken and formed.  An endothermic reaction occurs when a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the products.  Example: Separating H2O into H2 and O2  An exothermic reaction occurs when more energy is released during product bond formation than is required to break bonds in the reactants.  Example: Burning C with O2 to have CO2