Chapter 5 Compounds and Their Bonds

1

Covalent Bonds

Naming and Writing Formulas of Covalent Compounds

Bond Polarity

Chapter 5 Compounds and Their Bonds

2

In hydrogen, two hydrogen atoms share their electrons to form a covalent bond.

Each hydrogen atom acquires a stable outer shell of two (2) electrons like helium (He).

H + H H : H = HH = H2

hydrogen molecule

H2, A Covalent MoleculeIonic Bond

3

A covalent bond between two hydrogen atoms is A covalent bond between two hydrogen atoms is shown in this picture.shown in this picture.

Fig 5.1 A covalent bond is the result of attractive and repulsive forces between atoms.

4

Spherical 1S orbital of two individual hydrogen atoms bends together and overlap to give an egg shaped region in the hydrogen molecule. The shared pair of electrons in a covalent bond is often represented as a line between atoms.

5

Bond length: The optimum distance between nuclei involved in a covalent bond. If the atoms are too far apart, the attractive forces are small and no bond exists. If the atoms are too close, the repulsive interaction between the nuclei is so strong that it pushes the atoms apart, Fig 5.2.

6

When two chlorine atoms approach each other, the unpaired 3p electrons are shared by both atoms in a covalent bond. Each chlorine atom in the Cl2 molecule now have 6 electrons in its own valence shell and sharing two giving each valence shell octet.

7

Diatomic Elements

As elements, the following share electrons to form diatomic, covalent molecules.

8

In addition to H2 and Cl2, five other elements always exist as diatomic molecule.

9

What is the name of each of the following diatomic molecules?

H2 hydrogen

N2 nitrogen

Cl2 _______________

O2 _______________

I2 _______________

Learning Check

10

What are the names of each of the following diatomic molecules?

H2 hydrogen

N2 nitrogen

Cl2 chlorine

O2 oxygen

I2 iodine

Solution

11

The compound NH3 consists of a N atom and three H atoms.

N and 3 H

By sharing electrons to form NH3, the electron dot structure is written as

H Bonding pairs

H : N : H Lone pair of electrons

Covalent Bonds in NH3

12

Number of Covalent Bonds

Often, the number of covalent bonds formed by a nonmetal is equal to the number of electrons needed to complete the octet.

13

Dot Structures and Models of Some Covalent Compounds

14

Multiple Bonds

Sharing one pair of electrons is a single bond.

X : X or X–X In multiple bonds, two pairs of electrons are

shared to form a double bond or three pairs of electrons are shared in a triple bond.

X : : X or X =X

X : : : X or X ≡ X

15

In nitrogen, octets are achieved by sharing three pairs of electrons.

When three pairs of electrons are shared, the multiple bond is called a triple bond.

octets N +

N N:::N

triple bond

Multiple Bonds in N2

16

5.4 Coordinate Covalent BondsCoordinate Covalent Bond: The covalent bond that forms when both electrons are donated by the same atom.

17

Fig 5.7 Electronegativities and the periodic table

18

1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put the unique element ( or least electronegative atom) in the center.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Complete an octet for all atoms except hydrogen

4. If structure contains too many electrons, form double and triple bonds on central atom as needed.

5.6 Drawing Lewis Structure

UNKNOWN_PARAMETER_VALUE.mp413408K   Download  

19

Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

20

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

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

21

Two possible skeletal structures of formaldehyde (CH2O)

H C O HH

C OH

An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

formal charge on an atom in a Lewis structure

=1

2

total number of bonding electrons( )

total number of valence electrons in the free atom

-total number of nonbonding electrons

-

The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

22

H C O H

C : 4 e-

O : 6 e-

2H :2x1 e-

12 e-

2 single bonds (2x2) = 41 double bond = 4

2 lone pairs (2x2) = 4Total = 12

formal charge on C = 4 -2 - ½ x 6 = -1

formal charge on O = 6 -2 - ½ x 6 = +1

formal charge on an atom in a Lewis structure

=1

2

total number of bonding electrons( )

total number of valence electrons in the free atom

-total number of nonbonding electrons

-

-1 +1

23

C – 4 e-

O – 6 e-

2H – 2x1 e-

12 e-

2 single bonds (2x2) = 41 double bond = 4

2 lone pairs (2x2) = 4Total = 12

HC O

H

formal charge on C = 4 -0 - ½ x 8 = 0

formal charge on O = 6 -4 - ½ x 4 = 0

formal charge on an atom in a Lewis structure

=1

2

total number of bonding electrons( )

total number of valence electrons in the free atom

-total number of nonbonding electrons

-

0 0

24

Formal Charge and Lewis Structures

1. For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.

2. Lewis structures with large formal charges are less plausible than those with small formal charges.

3. 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 H

-1 +1 HC O

H

0 0

25

VSEPR

The shape of a molecule is predicted from the geometry of the electron pairs around the central atom.

In the valence-shell electron-pair repulsion theory (VSEPR), the electron pairs are arranged as far apart as possible to give the least amount of repulsion of the negatively charged electrons.

5.7 Shape of Molecules

26

Two Electron Pairs In a molecule of BeCl2, there are two bonding

pairs around the central atom Be. (Be is an exception to the octet rule.)

The arrangement of two electron pairs to minimize their repulsion is 180° or opposite each other.

The shape of the molecule is linear.

27

Two Electron Pairs with Double Bonds

The electron-dot structure for CO2 consists of two double bonds to the central atom C.

Because the electrons in a double bond are held together, a double bond is counted as a single unit.

Repulsion is minimized when the double bonds are placed opposite each other at 180° to give a linear shape.

28

Three Electron Pairs

In BF3, there are 3 electron pairs around the central atom B. (B is an exception to the octet rule.)

Repulsion is minimized by placing three electron pairs in a plane at angles of 120°, which is a trigonal planar arrangement.

The shape with three bonded atoms is trigonal planar.

29

Two Bonding Pairs and A Nonbonding Pair

In SO2, there are 3 electron units around the central atom S.

Two electron units are bonded to atoms and one electron pair is a nonbonding pair.

Repulsion is minimized by placing three electron pairs in a plane at angles of 120°, which is trigonal planar.

The shape with two bonded atoms is bent.

30

Four Electron Pairs

In CH4, there are 4 electron pairs around the central atom C.

Repulsion is minimized by placing four electron pairs at angles of 109°, which is a tetrahedral arrangement.

The shape with four bonded atoms is called tetrahedral.

31

Three Bonding Atoms and One Nonbonding Pair

In NH3, there are 4 electron pairs around the N. Three pairs are bonded to atoms and one is a

nonbonding pair. Repulsion is minimized by placing four electron

pairs at angles of 109°, which is a tetrahedral arrangement.

The shape with three bonded atoms is pyramidal.

32

Two Bonding Atoms and Two Lone Pairs

In H2O, there are 4 electron pairs around O. Two pairs are bonded to atoms and two are

nonbonding pairs. Repulsion is minimized by placing four

electron pairs at angles of 109° called a tetrahedral arrangement.

The shape with two bonded atoms is called bent.

33

Some Steps Using VSEPR to Predict Shape

Draw the electron dot structure. Count the charged clouds around the central

atom. Arrange the charged clouds to minimize

repulsion. Determine the shape using the number of

bonded atoms in the electron arrangement.

34

Summary of Electron Arrangements and Shapes

Number of atoms bonded to the central atom

35

The shape depends on the number of charged clouds surrounding the atom as summarized in Table 5.1

36

Learning Check

Use VSEPR theory to determine the shape of the following molecules or ions.

1) tetrahedral 2) pyramidal 3) bent

A. PF3

B. H2S

C. CCl4

D. PO43-

37

Solution

Use VSEPR theory to determine the shape of the following molecules or ions.

1) tetrahedral 2) pyramidal 3) bent

A. PF3 2) pyramidal

B. H2S 3) bent

C. CCl4 1) tetrahedral

D. PO43- 1) tetrahedral

38

Comparing Nonpolar and Polar Covalent Bonds

39

Electronegativity is the attraction of an atom for shared electrons.

The nonmetals have high electronegativity values with fluorine as the highest.

The metals have low electronegativity values.

Electronegativity

40

Fig 5.7 Electronegativities and the periodic table

41

42

The atoms in a nonpolar covalent bond have electronegativity differences of 0.4 or less.

Examples: Atoms Electronegativity Type of

Difference Bond N-N 3.0 - 3.0 = 0.0 Nonpolar covalentCl-Br 3.0 - 2.8 = 0.2 Nonpolar covalentH-Si 2.1 - 1.8 = 0.3 Nonpolar covalent

Nonpolar Covalent Bonds

43

The atoms in a polar covalent bond have electronegativity differences of 0.5 to 1.9.

Examples: Atoms Electronegativity Type of

Difference BondO-Cl 3.5 - 3.0 = 0.5 Polar covalentCl-C 3.0 - 2.5 = 0.5 Polar covalentO-S 3.5 - 2.5= 1.0 Polar covalent

Polar Covalent Bonds

44

Comparing Nonpolar and Polar Covalent Bonds

45

Polar, Nonpolar and Ionic Bond

46

Ionic Bonds

The atoms in an ionic bond have electronegativity differences of 2.0 or more.

Examples: Atoms Electronegativity Type of

Difference BondCl-K 3.0 – 0.8 = 2.2 IonicN-Na 3.0 – 0.9 = 2.1 Ionic

47

Predicting Bond Type

48

Identify the type of bond between the following as

1) nonpolar covalent

2) polar covalent

3) ionic A. K-N

B. N-O

C. Cl-Cl

Learning Check

49

A. K-N3) ionic

B. N-O2) polar covalent

C. Cl-Cl1) nonpolar covalent

Solution

50

5.9 Polar Molecules

Entire molecule can be polar if electrons are attracted more strongly to one part of the molecule than to another.

Molecule’s polarity is due to the sum of all individual bond polarities and lone-pair contribution in the molecule.

51

Molecular polarity is represented by an arrow pointing at the negative end and is crossed at the positive end to resemble a positive sign.

52

Molecular polarity depends on the shape of the molecule as well as the presence of polar covalent bonds and lone-pairs.

53

Would a linear water molecule be Polar?

Why is water not linear?

OO........HH

HH

54

55

Learning Check

Identify each of the following molecules as

1) polar or 2) nonpolar. Explain.

A. PBr3

B. HBr

C. Br2

D. SiBr4

56

Solution

Identify each of the following molecules as

1) polar or 2) nonpolar. Explain.

A. PBr3 1) polar; pyramidal

B. HBr 1) polar; polar bond

C. Br2 2) nonpolar, nonpolar bond

D. SiBr4 2) nonpolar; dipoles cancel

57

In the name of a covalent compound, the first nonmetal is named followed by the name of the second nonmetal ending in –ide.

Prefixes indicate the number of atoms of each element.

Naming Covalent Compounds

58

Complete the name of each covalent compound:

CO carbon ______oxide

CO2 carbon _______________

PCl3 phosphorus ___________

CCl4 carbon _______________

N2O ______________________

Learning Check

59

Complete the name of each covalent compound:

CO carbon monoxide

CO2 carbon dioxide

PCl3 phosphorus trichloride

CCl4 carbon tetrachloride

N2O dinitrogen monoxide

Solution

60

Formulas and Names of Some Covalent Compounds

61

Select the correct name for each compound.A. SiCl4 1) silicon chloride

2) tetrasilicon chloride3) silicon tetrachloride

B. P2O5 1) phosphorus oxide2) phosphorus pentoxide3) diphosphorus pentoxide

C. Cl2O7 1) dichlorine heptoxide2) dichlorine oxide3) chlorine heptoxide

Learning Check

62

Select the correct name for each compound.A. SiCl4 3) silicon tetrachloride

B. P2O5 3) diphosphorus pentoxide

C. Cl2O7 1) dichlorine heptoxide

Solution

63

Chapter Summary

Covalent bond: Bond formed by sharing of electrons between the atoms.

Molecule: A group of atoms held together by covalent bonds.

Coordinate covalent bond: Bond formed when a filled orbital containing lone pair of electrons on one atom overlaps a vacant orbital on another atom.

Molecular formula: Formula that shows the numbers and kinds of atoms in a molecule.

Lewis structure: shows how atoms are connected in a molecule.

64

Molecules have specific shapes that depend on the number of electron charge clouds surrounding the various atoms (VSEPR model)

Bonds between atoms are polar covalent if the bonding electrons are not shared equally between the atoms.

The ability of an atom to attract electrons in a covalent bond is the atom’s electronegativity.

Molecular compounds have lower melting points and boiling points than ionic compounds.

Chapter Summary Contd.Chapter Summary Contd.