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What to learn in this topic? Chemical Bonding and Structures: Please read P.5-20 for your preparation for the lecture. Much information on these pages has been covered in your high school chemistry syllabus.

 A quick review on  covalent, ionic and metallic bonding  Lewis structures of molecules and ions  electronegativity and bond polarity  bond order, bond lengths and bond energies

CHEM 1042 General Chemistry I

Topic 3 Chemical Bonding and Structures

Part I

 Shapes of Molecules  Bonding Theories: 1. VSEPR theory 2. VB theory

Dr. A P L Tong

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Part 2 2

Your reading: Chapter 10) Chemical Bonding I: Basic Concepts

Chemical Bonding I: Basic Concepts (Part 1)

11-1 – 11-4, Chapter 11) Chemical Bonding II: Valence Bond and Molecular Orbital Theories Petrucci; Herring; Madura; Bissonnette: General Chemistry: Principles and Modern Applications, 11th edition, Pearson

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Lewis Theory: An Overview

Why study Chemical Bonding? • The properties of many materials can be understood in terms of their microscopic properties.

• Chemical bonds are the forces that hold groups of atoms together in forming a compound.

• Microscopic properties of molecules include:

• Purpose of chemical bonding: To achieve a more stable electron configuration. Maximum stability results when an atom is isoelectronic with a noble gas (one with eight outer-shell electrons) – an octet.

 

The bonding between atoms The 3D shape of molecules

• Electrons, especially those of the valence shell, play a fundamental role in chemical bonding. • The type of bond to be formed depends on the electrostatic properties of bonding atoms. Diamond and graphite

View of DNA

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Lewis Theory: An Overview

Lewis Symbols and Lewis Structures • In a Lewis symbol, the element symbol represents the nucleus with core (non-valence) electrons. Dots surrounding the symbol represent the valence (outer-shell) electrons.

• Three types of chemical bonding: 

Covalent bond – one of more pairs of electrons are shared between atoms.

• •



Ionic bond – electrons are transferred from one atom to another; +ve and –ve ions are formed and attract each other through electrostatic forces.



Metallic bond – delocalization of electrons over a lattice of ions.

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• Si •

1s2 2s2 2p6 3s2 3p2 [Ne]

• A Lewis structure shows the arrangement of valence electrons (both bonding and non-bonding) and nuclei for the transfer or sharing of electrons in a chemical bond.

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Lewis Symbols and Lewis Structures

Lewis Symbols for Ionic Compounds

E.g.

• For an ionic compound of a main-group element, (1) the Lewis symbol of the metal ion has no dots if all the valence electrons are lost, and (2) the ionic charges of both cations and anions are shown

E.g.

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Lewis Symbols for Ionic Compounds

Lewis Symbols for Ionic Compounds

E.g. (Con’t)

• The structures of ionic compounds are much more complicated than is suggested by Lewis structure (see below for the structure of NaCl). So we seldom write Lewis structures for ionic compounds. MgCl2

Al2O3

Alternating Na+ and Cl ions extend in all directions and involves countless numbers of ions. CHEM1042 Dr. A P L Ton g 201 7-18

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Crystal Structures of Ionic Compounds

Ionic Bonding • Ionic bonds are formed by the transfer of e- between atoms. • Positive and negative ions joined together by electrostatic forces of attraction to form an ionic compound. • In forming an ionic compound,  

• In crystalline solids, rigid and highly regular arrangement of their components (atoms, molecules, or ions) is observed. • The patterns of continuous three dimensional arrangement are outlined against a framework called a lattice.

Metallic atoms: tend to lose one or more e-s → become cations Nonmetal atoms: tend to gain one or more e-s → become anions

=

= replicated unit cell

• After electron transfer, the ions achieve the octet structures. • To achieve an electrically neutral state, ions in ionic compound are in specific ratio :  

• Crystal lattice is formed by the repetition of unit cell, which is the simplest structural unit of a crystalline solid.

Sodium chloride: Na+ + Cl ion [+1 – 1 = 0] → NaCl Magnesium chloride: Mg2+ + 2 Cl ions [+2 – (1 x 2) = 0] → MgCl2 13

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Strength of Ionic Bond

Metallic Bonding

• Ionic charge:

• Metallic bonding is the electrostatic attractive force between the metal cations and the sea of delocalized electrons.

Compound



Lattice Energy (kJmol-1)

NaCl

786

MgCl2

2526

• The electron sea model of metal:

As no. of ionic charge ↑, charge density ↑, ionic bond strength ↑, lattice energy more negative.

• Size of metallic radius:



Compound

Lattice Energy (kJmol-1)

NaCl

786

KCl

715

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

Valence electrons of a metal atom are easily removable.



Removal of valence electrons leads to the formation of metallic cations.



Metallic cations are closely packed in a crystal lattice (the metal lattice).



In the metal lattice, valence electrons are delocalized from nuclei and free to move around the whole lattice.



A network of positive ions is immersed in a “sea of electrons”.

As size of metallic radius ↓, charge density ↑, ionic bond strength ↑, lattice energy more negative.

(Lattice energy = the negative of energy required to completely separate one mole of a solid ionic compound into gaseous ionic constituents.) CHEM1042 Dr. A P L Ton g 201 7-18

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Strength of Metallic Bond

Covalent Bonding

• No. of valence electrons:

• In a covalent compound, electrons are shared between bonding atoms via covalent bond.

Metal



No. of valence

e-

Melting point ( )

Na

1

98

Mg

2

650

Al

3

660

As no. of valence electrons ↑,metallic bond strength ↑, melting point ↑.

• Size of metallic radius: Metal



• Bonding atoms (H, O and Cl) attain noble gas electron configuration (He, Ne and Ar) by sharing e-.

Metallic radius (nm) Melting point ( )

Li

0.152

181

Na

0.186

98

K

0.231

64

For metals with the same no. of valence electrons, as metallic radius ↓, charge density ↑, metallic bond strength ↑, melting point ↑

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• Electron pair in a covalent bond is the bond pair, while the nonbonding pair is called the lone pair. 18

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Coordinate Covalent Bonds

Multiple Covalent Bonds

• Not necessary for each atom in a compound to contribute an electron to form a covalent bond.

• Often, more than one pair of electrons has to be shared in order for the bonding atoms to attain the noble gas electron configuration.

• N

N

N

••

N •

•• O ••

C

••

• N •

••

••

•• • N •

•• O ••

• O ••

••

• •N •

• C •

••

••

• N• •

• O ••

••

• In the reaction of NH3 with HCl, the lone pair electrons on N atom of NH3 attacks the H atom of HCl. This breaks the H–Cl bond and forms a new coordinate covalent bond [H–NH3]+.

• •O ••

••

• • C• •

••

• O• ••

••

• A covalent bond in which a single atom contributes both of the electrons to a shared pair is called a coordinate covalent bond.

(A half arrow is used to represent the movement of a single electron)

• A double covalent bond is formed when two pairs of electrons (a total of four electrons) are shared. When three pairs of electrons are shared, a triple covalent bond is formed.

(A full arrow is used to represent the movement of an electron pair) CHEM1042 Dr. A P L Ton g 201 7-18

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Polar Covalent Bonds

Polar Covalent Bonds

Electrostatic Potential Maps

• Except homonuclear covalent bonds, most of covalent bonds do not have electron pairs equally shared between bonding atoms.

•

The electrostatic potential is the work done in moving a unit of positive charge at a constant speed from one region of a molecule to another. The electrostatic potential map is obtained by hypothetically probing an electron density surface with a positive point charge.

•

An electrostatic potential map is a way to visualize the charge distribution within a molecule.

• When two atoms with different affinity for electrons are bonded covalently, a polar covalent bond is formed (i.e. a covalent bond in which electrons are not shared equally between the two atoms). •• δ-

Cl

••

H

••

δ+

••

• In a polar covalent bond, electrons are displaced toward the atom having larger electron affinity, leading to a partial negative charge on it (signified by δ). A corresponding partial positive charge is resulted on the other bonding atom (signified by δ+). • An atom’s ability to compete for electrons with other atoms to which it is bonded is expressed as the electronegativity (EN). CHEM1042 Dr. A P L Ton g 201 7-18

Electrostatic Potential Maps

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Polar Covalent Bonds

•

E.g.) In a neutral molecule, if the potential at a point is +ve, it is likely that an atom at this point carries a net +ve charge.

•

Blue is used to colour regions of the most +ve potential; red for most –ve potential. Intermediate colours represent intermediate values of the electrostatic potential.

The colours show the distribution of charge in the molecule.

Fig. 10-4) Determination of the electrostatic potential map for ammonia The electrostatic potential at any point on the charge density surface of a molecule is defined as the change in energy that occurs when a unit positive charge is brought to this point, starting from another point that is infinitely far removed from the molecule. The surface encompassing the ammonia molecule is analogous to the 95% surface of electron charge density for atomic orbitals. CHEM1042 Dr. A P L Ton g 201 7-18

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Electronegativity • It refers to the ability of an atom to compete for electrons with other atoms to which it is bonded. • In the periodic table, electronegativity generally increases across a period and decreases down a group.

Fig. 10-5) The electrostatic potential maps for sodium chloride, hydrogen chloride, and chlorine The dark red and dark blue on the electrostatic potential map correspond to the extremes of the electrostatic potential, negative to positive, for the particular molecule for which the map is calculated. To get a reliable comparison of different molecules, the values of the extremes in electrostatic potential (in kJmol–1) must be the same for all of the molecules compared. In the maps shown here the range is –157 to 157 kJ mol–1.

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The values in this table are from L. Pauling. The range is from 4.0 for fluorine (the most EN) to 0.7 for francium (the least EN). CHEM1042 Dr. A P L Tong 201 7-18

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Electronegativity

Electronegativity

• EN is related to ionization energy (I) and electron affinity (EA).

• The absolute value of the difference in electronegativity (∆EN) of the bonded atoms determines the amount of polar character in a covalent bond.

• The ionization energy (I) is the quantity of energy a gaseous atom must absorb to be able to expel an electron. Mg (g)  Mg+ (g) + e-

I1 = 738 kJ/mol

Mg+ (g)  Mg2+ (g) + e-

I2 = 1451 kJ/mol

Fig. 10-7) Percent ionic character of a chemical bond as a function of electronegativity difference

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Class Practice 1

ionic polar covalent

• Electron affinity (EA) can be defined as the enthalpy change (Hea) that occurs when an atom in the gas phase gains an electron. F (g) + e-  F (g)

covalent

Hea = 328 kJ/mol

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Electronegativity

• For ∆EN is very small for the two bonded atoms (e.g. IBr), the bond between them is essentially covalent. • For ∆EN is large for the two bonded atoms (e.g. LiF), the bond between them is essentially ionic. CHEM1042 Dr. A P L Tong 201 7-18

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Writing Lewis Structures • Essential features of Lewis structures: 

All the valence electrons of the atoms must appear in the structure.



Usually, all the electrons in a Lewis structure are paired.



Usually, each atom requires an outer-shell octet of electrons (H only requires 2 outer-shell e-).



Sometimes, multiple bonds are needed (most readily formed by C, N, O, P, and S).

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Skeletal Structures

Which of the following is the correct Lewis structure for chloric acid, HOClO2?

• Starting point in writing a Lewis structure is to designate the skeletal structure. • In the skeletal structure:  

A central atom is bonded to two or more atoms. A terminal atom is bonded to just one other atom.

C, C, O here are central atoms; others are terminal.



Hydrogen atoms are always terminal atoms.



Central atoms are generally those with the lowest electronegativity.

Which of the following is the correct Lewis structure for phosphoric acid, H3PO4?



Carbon atoms are always central atoms.



Generally structures of molecules and polyatomic ions are compact and symmetrical.

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A Strategy for Writing Lewis Structures

A Strategy for Writing Lewis Structures Summary scheme for drawing Lewis Structures

(the previous figure) valence

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Class Practice 2: Writing a Lewis Structure for a Polyatomic Ion

Class Practice 2 (Con’t)

Petrucci Example 10-7

Write the Lewis structure for the nitronium ion, NO2+.

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Formal Charge

Formal Charge

• 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.

E.g.) For the structure for the nitronium ion, NO2+ (Class Practice 2) +

+

••

O≡N—O ••

••

Improbable structure: FC = no. of valence e- – no. of lone pair e- – ½ no. of bond pair e-

•• -

+

• General rules: 

All of the lone pair electrons are assigned to the atom on which they are found.



Half of the bonding electrons are assigned to each atom in the bond.



The sum of FCs is the overall charge of the compound.



FCs should be as small as possible.



Negative FCs usually on most electronegative atoms; positive FCs on the least electronegative atoms.



Structures having FC of same sign on adjacent atoms are unlikely.

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The structure should be: This has only one formal charge, +1, on the central atom. It is the most satisfactory Lewis structure. 35

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Class Practice 3: Using Formal Charges in Writing Lewis Structures

Class Practice 3 (Con’t)

Petrucci Example 10-8

Write the most plausible Lewis structure of nitrosyl chloride, NOCI.

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Formal Charge

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Resonance • Some molecules can have more than one valid Lewis structure. • For instance, two Lewis structures can be drawn for ozone molecule (O3).

• From experiment, only one bond length for O3 is found (127.8 pm)  O3 has two equivalent bonds  intermediate value between the O–O single-bond length of 147.5 pm in HO–OH and the double-bond length of 120.74 pm in O=O. CHEM1042 Dr. A P L Ton g 201 7-18

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Resonance

Resonance

=

• Neither of the two proposed Lewis structures above can represent the true structure of O3. Actually, O 3 exists as of the hybrid of both Lewis structures. This phenomenon is called resonance, which is caused by the delocalization of electrons. • Resonance structures of O3:

Bond order = 1.5

• As the charge density is spread out, resonance helps stabilize the whole molecule.

= (The two structures are equivalent, contributing equally to the resonance hybrid structure.) We use a double-headed arrow to show that individual structures are related by resonance.

• By resonance, the lone-pair electrons of one O atom and the bond-pair electrons in the double bond (π bond) delocalize over the three atoms of the molecule.

• Resonance structures differ only in the assignment of electron pair positions, the positions of atoms are unchanged. • The ...