Lecture 7 & 8- Basic Concepts Of Chemical Bonding
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General Chemistry Course # 111, two credits Second Semester 2009
King Saud bin Abdulaziz University for Health Science Textbook: Principles of Modern Chemistry by David W. Oxtoby, H. Pat Gillis, and Alan Campion (6 edition; 2007)
Dr. Rabih O. Al-Kaysi Ext: 47247 Email: [email protected]
Lectures 7 & 8
Basic Concepts of Chemical Bonding
Chemical Bonds, Lewis Symbols, and the Octet Rule • Chemical bond: attractive force holding two or more atoms together. • Covalent bond results from sharing electrons between the atoms. Usually found between nonmetal elements. • Ionic bond results from the transfer of electrons from a metal to a nonmetal element. • Metallic bond: attractive force holding pure metal elements together.
Lewis Symbols • As a pictorial understanding of where the electrons are in an atom, we represent the electrons as dots around the symbol for the element. • The number of electrons available for bonding are indicated by unpaired dots. • These symbols are called Lewis symbols. • We generally place the electrons one four sides of a square around the element symbol.
Lewis Symbols
The Octet Rule • All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). • Caution: there are many exceptions to the octet rule.
Ionic Bonding • Na reacts with Chlorine: The reaction is violently exothermic. • We deduce that the NaCl is more stable than its constituent elements. • Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl−. Note: Na+ has an Ne electron configuration and Cl− has an Ar configuration. • That is, both Na+ and Cl− have an octet of electrons surrounding the central ion.
Ionic Bonding Arrangement
Energetics of Ionic Bond Formation • Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions. • Lattice energy depends on the charges on the ions and the sizes of the ions: Q1Q2 El = κ d κ
is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions.
Electron Configurations of Ions of the Representative Elements
• These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital. • Electron configurations can predict stable ion formation: • • • • •
Mg: [Ne]3s2 Mg+: [Ne]3s1 not stable Mg2+ : [Ne] stable Cl: [Ne]3s23p5 Cl−: [Ne]3s23p6 = [Ar] stable
Covalent Bonding • When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. • When similar atoms bond, they share pairs of electrons to each obtain an octet. • Each pair of shared electrons constitutes one chemical bond. • Example: H + H → H2 has electrons on a line connecting the two H nuclei.
Covalent Bonding
Lewis Structures • Covalent bonds can be represented by the Lewis symbols of the elements: Cl + Cl Cl Cl • In Lewis structures, each pair of electrons in a bond is represented by a single line: Cl Cl
H F
H O H
H N H H
H H C H H
Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): • One shared pair of electrons = single bond (e.g. H2); • Two shared pairs of electrons = double bond (e.g. O2); • Three shared pairs of electrons = triple bond (e.g. N2).
H H
O O
N N
• Generally, bond distances decrease as we move from single through double to triple bonds.
Bond Polarity and Electronegativity • In a covalent bond, electrons are shared. • Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. • There are some covalent bonds in which the electrons are located closer to one atom than the other. • Unequal sharing of electrons results in polar bonds.
Electronegativity • Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F).
• Electronegativity increases • Left to Right a period. • Up a group.
Electronegativity
Bond Polarity and Electronegativity • Difference in electronegativity is a gauge of bond polarity: • electronegativity differences around 0 result in non-polar covalent bonds (equal or almost equal sharing of electrons); • electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons); • electronegativity differences around 3 result in ionic bonds (transfer of electrons).
Bond Polarity and Electronegativity • There is no sharp distinction between bonding types. • The positive end (or pole) in a polar bond is represented δ + and the negative pole δ -.
Dipole Moments • Consider HF: • The difference in electronegativity leads to a polar bond. • There is more electron density on F than on H. • Since there are two different “ends” of the molecule, we call HF a dipole.
• Dipole moment, µ , is the magnitude of the dipole:
µ = Qr
where Q is the magnitude of the charges. • Dipole moments are measured in debyes, D.
Drawing Lewis Structures 1. Add the valence electrons of all atoms. 2. Write symbols for the atoms and show which atoms are connected to which. 3. Complete the octet for all atoms bounded to the central atom. 4. Place leftover electrons on the central atom even if it results in more than an octet of electrons around the atoms. 5. If there are not enough electrons to give the central atom an octet, try multiple bonds.
Formal Charge • It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. • To determine which structure is most reasonable, we use formal charge. • Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity (i.e., if the electrons are shared equally between atoms).
Formal Charge • To calculate formal charge: • •
All nonbonding electrons are assigned to the atom on which they are found. Half the bonding electrons are assigned to each atom in a bond.
• Formal charge is: valence electrons - number of bonds - lone pair electrons
Formal Charge • Consider:
C N • For C: • •
•
There are 4 valence electrons (from periodic table). In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge: 4 - 5 = -1.
Formal Charge • Consider:
C N • For N: • • •
There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge = 5 - 5 = 0.
• We write:
C N
Formal Charge • The most stable structure has: • •
the lowest formal charge on each atom (i.e., closest to zero), the most negative formal charge on the most electronegative atoms.
Resonance Structures • • •
Some molecules are not well described by Lewis Structures. Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).
O
O
O
Resonance Structures
Resonance Structures • Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities.
Resonance Structures • Example: in ozone the extreme possibilities have one double and one single bond. The resonance structure has two identical bonds of intermediate character.
O O
O
O
O O
• Common examples: O3, NO3-, SO42- , NO2, and benzene.
Exceptions to the Octet Rule • There are three classes of exceptions to the octet rule: – Molecules with an odd number of electrons; – Molecules in which one atom has less than an octet; – Molecules in which one atom has more than an octet.
Odd Number of Electrons • Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.
N O
N O
Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3. • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.
More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.
Bond Strength and Bond Length • We know that multiple bonds are shorter than single bonds. • We can show that multiple bonds are stronger than single bonds. • As the number of bonds between atoms increases, the atoms are held closer and more tightly together.
Strengths of Covalent Bonds
King Saud bin Abdulaziz University for Health Science Textbook: Principles of Modern Chemistry by David W. Oxtoby, H. Pat Gillis, and Alan Campion (6 edition; 2007)
Dr. Rabih O. Al-Kaysi Ext: 47247 Email: [email protected]
Lectures 7 & 8
Basic Concepts of Chemical Bonding
Chemical Bonds, Lewis Symbols, and the Octet Rule • Chemical bond: attractive force holding two or more atoms together. • Covalent bond results from sharing electrons between the atoms. Usually found between nonmetal elements. • Ionic bond results from the transfer of electrons from a metal to a nonmetal element. • Metallic bond: attractive force holding pure metal elements together.
Lewis Symbols • As a pictorial understanding of where the electrons are in an atom, we represent the electrons as dots around the symbol for the element. • The number of electrons available for bonding are indicated by unpaired dots. • These symbols are called Lewis symbols. • We generally place the electrons one four sides of a square around the element symbol.
Lewis Symbols
The Octet Rule • All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). • Caution: there are many exceptions to the octet rule.
Ionic Bonding • Na reacts with Chlorine: The reaction is violently exothermic. • We deduce that the NaCl is more stable than its constituent elements. • Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl−. Note: Na+ has an Ne electron configuration and Cl− has an Ar configuration. • That is, both Na+ and Cl− have an octet of electrons surrounding the central ion.
Ionic Bonding Arrangement
Energetics of Ionic Bond Formation • Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions. • Lattice energy depends on the charges on the ions and the sizes of the ions: Q1Q2 El = κ d κ
is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions.
Electron Configurations of Ions of the Representative Elements
• These are derived from the electron configuration of elements with the required number of electrons added or removed from the most accessible orbital. • Electron configurations can predict stable ion formation: • • • • •
Mg: [Ne]3s2 Mg+: [Ne]3s1 not stable Mg2+ : [Ne] stable Cl: [Ne]3s23p5 Cl−: [Ne]3s23p6 = [Ar] stable
Covalent Bonding • When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. • When similar atoms bond, they share pairs of electrons to each obtain an octet. • Each pair of shared electrons constitutes one chemical bond. • Example: H + H → H2 has electrons on a line connecting the two H nuclei.
Covalent Bonding
Lewis Structures • Covalent bonds can be represented by the Lewis symbols of the elements: Cl + Cl Cl Cl • In Lewis structures, each pair of electrons in a bond is represented by a single line: Cl Cl
H F
H O H
H N H H
H H C H H
Multiple Bonds • It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): • One shared pair of electrons = single bond (e.g. H2); • Two shared pairs of electrons = double bond (e.g. O2); • Three shared pairs of electrons = triple bond (e.g. N2).
H H
O O
N N
• Generally, bond distances decrease as we move from single through double to triple bonds.
Bond Polarity and Electronegativity • In a covalent bond, electrons are shared. • Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. • There are some covalent bonds in which the electrons are located closer to one atom than the other. • Unequal sharing of electrons results in polar bonds.
Electronegativity • Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F).
• Electronegativity increases • Left to Right a period. • Up a group.
Electronegativity
Bond Polarity and Electronegativity • Difference in electronegativity is a gauge of bond polarity: • electronegativity differences around 0 result in non-polar covalent bonds (equal or almost equal sharing of electrons); • electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons); • electronegativity differences around 3 result in ionic bonds (transfer of electrons).
Bond Polarity and Electronegativity • There is no sharp distinction between bonding types. • The positive end (or pole) in a polar bond is represented δ + and the negative pole δ -.
Dipole Moments • Consider HF: • The difference in electronegativity leads to a polar bond. • There is more electron density on F than on H. • Since there are two different “ends” of the molecule, we call HF a dipole.
• Dipole moment, µ , is the magnitude of the dipole:
µ = Qr
where Q is the magnitude of the charges. • Dipole moments are measured in debyes, D.
Drawing Lewis Structures 1. Add the valence electrons of all atoms. 2. Write symbols for the atoms and show which atoms are connected to which. 3. Complete the octet for all atoms bounded to the central atom. 4. Place leftover electrons on the central atom even if it results in more than an octet of electrons around the atoms. 5. If there are not enough electrons to give the central atom an octet, try multiple bonds.
Formal Charge • It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. • To determine which structure is most reasonable, we use formal charge. • Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity (i.e., if the electrons are shared equally between atoms).
Formal Charge • To calculate formal charge: • •
All nonbonding electrons are assigned to the atom on which they are found. Half the bonding electrons are assigned to each atom in a bond.
• Formal charge is: valence electrons - number of bonds - lone pair electrons
Formal Charge • Consider:
C N • For C: • •
•
There are 4 valence electrons (from periodic table). In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge: 4 - 5 = -1.
Formal Charge • Consider:
C N • For N: • • •
There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge = 5 - 5 = 0.
• We write:
C N
Formal Charge • The most stable structure has: • •
the lowest formal charge on each atom (i.e., closest to zero), the most negative formal charge on the most electronegative atoms.
Resonance Structures • • •
Some molecules are not well described by Lewis Structures. Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms Example: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).
O
O
O
Resonance Structures
Resonance Structures • Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities.
Resonance Structures • Example: in ozone the extreme possibilities have one double and one single bond. The resonance structure has two identical bonds of intermediate character.
O O
O
O
O O
• Common examples: O3, NO3-, SO42- , NO2, and benzene.
Exceptions to the Octet Rule • There are three classes of exceptions to the octet rule: – Molecules with an odd number of electrons; – Molecules in which one atom has less than an octet; – Molecules in which one atom has more than an octet.
Odd Number of Electrons • Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.
N O
N O
Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3. • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.
More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.
Bond Strength and Bond Length • We know that multiple bonds are shorter than single bonds. • We can show that multiple bonds are stronger than single bonds. • As the number of bonds between atoms increases, the atoms are held closer and more tightly together.
Strengths of Covalent Bonds
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