Cheat sheet: Chemical Bonding and Molecular Structure | Chemistry Class 11 - NEET PDF Download

 Page 1


Chemical Bonding and Molecular Structure Cheat
Sheet (Class 11 CBSE)
1 Introduction to Chemical Bonding
• Chemical Bond : Attr active force holding atoms, ions, etc., together in a molecule or com-
pound.
• Wh y Atoms Combine : T o achieve stability b y attaining noble gas electron configur ation
(lower energy state).
• K ey Theories :
– K össel-Lewis Approach
– V alence Shell Electron Pair Repulsion (VSEPR) Theory
– V alence Bond (VB) Theory
– Molecular Orbital (MO) Theory
2 K ö ssel-Lewis Approach
• Lewis S ymbols : Dots around element symbol represent valence electrons (e.g., Na· , Cl:· ).
• Octet Rule : At oms combine to achieve 8 electrons in their valence shell (noble gas config-
ur ation).
– Ionic Bond : Electron tr ansfer (e.g., Na? Na
+
+ e
-
; Cl + e
-
? Cl
-
; forms NaCl).
– Covalent Bond : Electron sharing (e.g., Cl
2
: Cl:Cl, sharing one electron pair).
• Limitations of Octet Rule :
– Incomplete Octet : e.g., LiCl, BeH
2
, BCl
3
(centr al atom < 8 electrons).
– Odd-Electron Molecules : e.g., NO , NO
2
(odd number of electrons).
– Expanded Octet : e.g., PF
5
, SF
6
(centr al atom > 8 electrons due to d-orbitals).
3 Types of Bonds
• Ionic (Electrovalent) Bond :
– F ormed b y electron tr ansfer between electropositive (low ionization enthalp y) and
electronegative (high electron gain enthalp y) elements.
– Stabilized b y lattice enthalp y (energy to separ ate 1 mole of ionic solid into gaseous ions,
e.g., NaCl: 788 kJ/mol).
• Covalent Bond :
– F ormed b y sharing electron pairs (single, double, or triple bonds).
– Single Bond : One shared pair (e.g., H
2
).
– Double Bond : Two shared pairs (e.g., O
2
, CO
2
).
1
Page 2


Chemical Bonding and Molecular Structure Cheat
Sheet (Class 11 CBSE)
1 Introduction to Chemical Bonding
• Chemical Bond : Attr active force holding atoms, ions, etc., together in a molecule or com-
pound.
• Wh y Atoms Combine : T o achieve stability b y attaining noble gas electron configur ation
(lower energy state).
• K ey Theories :
– K össel-Lewis Approach
– V alence Shell Electron Pair Repulsion (VSEPR) Theory
– V alence Bond (VB) Theory
– Molecular Orbital (MO) Theory
2 K ö ssel-Lewis Approach
• Lewis S ymbols : Dots around element symbol represent valence electrons (e.g., Na· , Cl:· ).
• Octet Rule : At oms combine to achieve 8 electrons in their valence shell (noble gas config-
ur ation).
– Ionic Bond : Electron tr ansfer (e.g., Na? Na
+
+ e
-
; Cl + e
-
? Cl
-
; forms NaCl).
– Covalent Bond : Electron sharing (e.g., Cl
2
: Cl:Cl, sharing one electron pair).
• Limitations of Octet Rule :
– Incomplete Octet : e.g., LiCl, BeH
2
, BCl
3
(centr al atom < 8 electrons).
– Odd-Electron Molecules : e.g., NO , NO
2
(odd number of electrons).
– Expanded Octet : e.g., PF
5
, SF
6
(centr al atom > 8 electrons due to d-orbitals).
3 Types of Bonds
• Ionic (Electrovalent) Bond :
– F ormed b y electron tr ansfer between electropositive (low ionization enthalp y) and
electronegative (high electron gain enthalp y) elements.
– Stabilized b y lattice enthalp y (energy to separ ate 1 mole of ionic solid into gaseous ions,
e.g., NaCl: 788 kJ/mol).
• Covalent Bond :
– F ormed b y sharing electron pairs (single, double, or triple bonds).
– Single Bond : One shared pair (e.g., H
2
).
– Double Bond : Two shared pairs (e.g., O
2
, CO
2
).
1
– Triple Bond : Three shared pairs (e.g., N
2
, C
2
H
2
).
• Hydrogen Bond :
– W eak bond between H (bonded to F , O , N) and another electronegative atom (F , O , N).
– Intermolecular : Between different molecules (e.g., H
2
O , HF).
– Intr amolecular : Within the same molecule (e.g., o-nitrophenol).
– Stronger than van der W aals forces but weak er than covalent/ionic bonds.
4 Bond Par ameters
• Bond Length : Equilibrium distance between nuclei of bonded atoms (e.g., H
2
: 74 pm, C=C:
133 pm).
– Covalent r adius = half the distance between two similar atoms in a covalent bond.
• Bond Angle : Angle between orbitals containing bonding electron pairs (e.g., H
2
O: 104.5
?
).
• Bond Enthalp y : Energy to break 1 mole of a bond (e.g., H
2
: 435.8 kJ/mol, N
2
: 946 kJ/mol).
– Higher bond order? higher bond enthalp y , shorter bond length.
• Bond Order : Number of bonds between atoms (e.g., H
2
: 1, O
2
: 2, N
2
: 3).
• Dipole Moment (?? ) : Product of charge and distance (?? = ??× ?? , in Deb ye, 1D = 3.33564 ×
10
-30
C· m).
– Polar Covalent Bond : Unequal electron sharing (e.g., HF: H
+??
–F
-??
).
– Nonpolar Covalent Bond : Equal sharing (e.g., H
2
, O
2
).
– Net dipole depends on molecular geometry (e.g., CO
2
: linear , ?? = 0 ; H
2
O: bent, ?? = 1.85
D).
5 VSE PR Theory
• Principle : Electron pairs (bond pairs, lone pairs) around centr al atom repel each other ,
minimizing repulsion to determine molecular geometry .
• Repulsion Order : Lone pair-lone pair (lp-lp) > Lone pair-bond pair (lp-bp) > Bond pair-
bond pair (bp-bp).
• Geometries (No Lone Pairs) :
– 2 electron pairs: Linear (e.g., BeCl
2
, 180
?
).
– 3 electron pairs: Trigonal planar (e.g., BCl
3
, 120
?
).
– 4 electron pairs: T etr ahedr al (e.g., CH
4
, 109.5
?
).
– 5 electron pairs: Trigonal bip yr amidal (e.g., PCl
5
, 120
?
/90
?
).
– 6 electron pairs: Octahedr al (e.g., SF
6
, 90
?
).
• Geometries (With Lone Pairs) :
– AB
2
E: Bent (e.g., SO
2
,~ 119.5
?
).
– AB
3
E: Trigonal p yr amidal (e.g., NH
3
, 107
?
).
– AB
2
E
2
: Bent (e.g., H
2
O , 104.5
?
).
– AB
4
E: See-saw (e.g., SF
4
).
– AB
3
E
2
: T-shaped (e.g., ClF
3
).
2
Page 3


Chemical Bonding and Molecular Structure Cheat
Sheet (Class 11 CBSE)
1 Introduction to Chemical Bonding
• Chemical Bond : Attr active force holding atoms, ions, etc., together in a molecule or com-
pound.
• Wh y Atoms Combine : T o achieve stability b y attaining noble gas electron configur ation
(lower energy state).
• K ey Theories :
– K össel-Lewis Approach
– V alence Shell Electron Pair Repulsion (VSEPR) Theory
– V alence Bond (VB) Theory
– Molecular Orbital (MO) Theory
2 K ö ssel-Lewis Approach
• Lewis S ymbols : Dots around element symbol represent valence electrons (e.g., Na· , Cl:· ).
• Octet Rule : At oms combine to achieve 8 electrons in their valence shell (noble gas config-
ur ation).
– Ionic Bond : Electron tr ansfer (e.g., Na? Na
+
+ e
-
; Cl + e
-
? Cl
-
; forms NaCl).
– Covalent Bond : Electron sharing (e.g., Cl
2
: Cl:Cl, sharing one electron pair).
• Limitations of Octet Rule :
– Incomplete Octet : e.g., LiCl, BeH
2
, BCl
3
(centr al atom < 8 electrons).
– Odd-Electron Molecules : e.g., NO , NO
2
(odd number of electrons).
– Expanded Octet : e.g., PF
5
, SF
6
(centr al atom > 8 electrons due to d-orbitals).
3 Types of Bonds
• Ionic (Electrovalent) Bond :
– F ormed b y electron tr ansfer between electropositive (low ionization enthalp y) and
electronegative (high electron gain enthalp y) elements.
– Stabilized b y lattice enthalp y (energy to separ ate 1 mole of ionic solid into gaseous ions,
e.g., NaCl: 788 kJ/mol).
• Covalent Bond :
– F ormed b y sharing electron pairs (single, double, or triple bonds).
– Single Bond : One shared pair (e.g., H
2
).
– Double Bond : Two shared pairs (e.g., O
2
, CO
2
).
1
– Triple Bond : Three shared pairs (e.g., N
2
, C
2
H
2
).
• Hydrogen Bond :
– W eak bond between H (bonded to F , O , N) and another electronegative atom (F , O , N).
– Intermolecular : Between different molecules (e.g., H
2
O , HF).
– Intr amolecular : Within the same molecule (e.g., o-nitrophenol).
– Stronger than van der W aals forces but weak er than covalent/ionic bonds.
4 Bond Par ameters
• Bond Length : Equilibrium distance between nuclei of bonded atoms (e.g., H
2
: 74 pm, C=C:
133 pm).
– Covalent r adius = half the distance between two similar atoms in a covalent bond.
• Bond Angle : Angle between orbitals containing bonding electron pairs (e.g., H
2
O: 104.5
?
).
• Bond Enthalp y : Energy to break 1 mole of a bond (e.g., H
2
: 435.8 kJ/mol, N
2
: 946 kJ/mol).
– Higher bond order? higher bond enthalp y , shorter bond length.
• Bond Order : Number of bonds between atoms (e.g., H
2
: 1, O
2
: 2, N
2
: 3).
• Dipole Moment (?? ) : Product of charge and distance (?? = ??× ?? , in Deb ye, 1D = 3.33564 ×
10
-30
C· m).
– Polar Covalent Bond : Unequal electron sharing (e.g., HF: H
+??
–F
-??
).
– Nonpolar Covalent Bond : Equal sharing (e.g., H
2
, O
2
).
– Net dipole depends on molecular geometry (e.g., CO
2
: linear , ?? = 0 ; H
2
O: bent, ?? = 1.85
D).
5 VSE PR Theory
• Principle : Electron pairs (bond pairs, lone pairs) around centr al atom repel each other ,
minimizing repulsion to determine molecular geometry .
• Repulsion Order : Lone pair-lone pair (lp-lp) > Lone pair-bond pair (lp-bp) > Bond pair-
bond pair (bp-bp).
• Geometries (No Lone Pairs) :
– 2 electron pairs: Linear (e.g., BeCl
2
, 180
?
).
– 3 electron pairs: Trigonal planar (e.g., BCl
3
, 120
?
).
– 4 electron pairs: T etr ahedr al (e.g., CH
4
, 109.5
?
).
– 5 electron pairs: Trigonal bip yr amidal (e.g., PCl
5
, 120
?
/90
?
).
– 6 electron pairs: Octahedr al (e.g., SF
6
, 90
?
).
• Geometries (With Lone Pairs) :
– AB
2
E: Bent (e.g., SO
2
,~ 119.5
?
).
– AB
3
E: Trigonal p yr amidal (e.g., NH
3
, 107
?
).
– AB
2
E
2
: Bent (e.g., H
2
O , 104.5
?
).
– AB
4
E: See-saw (e.g., SF
4
).
– AB
3
E
2
: T-shaped (e.g., ClF
3
).
2
6 V a lence Bond (VB) Theory
• Principle : Coval ent bond forms b y overlap of half-filled atomic orbitals with opposite spins.
• Types of Overlap :
– Sigma (?? ) Bond : Head-on overlap (s-s, s-p, p-p).
– Pi (?? ) Bond : Sidewa ys overlap (p-p).
– Sigma bonds are stronger due to greater overlap.
• Hybridisation : Mixing of atomic orbitals to form equivalent h ybrid orbitals.
– sp : Linear , 180
?
(e.g., BeCl
2
).
– sp
2
: Trigonal planar , 120
?
(e.g., BCl
3
).
– sp
3
: T etr ahedr al, 109.5
?
(e.g., CH
4
, NH
3
, H
2
O).
– sp
3
d : Trigonal bip yr amidal (e.g., PCl
5
, axial bonds longe r due to more repulsion).
– sp
3
d
2
: Octahedr al (e.g., SF
6
).
• Examples :
– C
2
H
6
: sp
3
-sp
3
?? bonds (C–C, C–H).
– C
2
H
4
: sp
2
-sp
2
?? bond (C–C), ?? bond (C=C), sp
2
-s ?? bonds (C–H).
– C
2
H
2
: sp-sp ?? bond (C–C), two ?? bonds (C= C), sp-s ?? bonds (C–H).
7 Molecular Orbital (MO) Theory
• Principle : Atomic orbitals combine to form molecular orbitals (bonding: lower energy;
antibonding: higher energy).
• Linear Combination of Atomic Orbitals (LCA O) :
– Bonding MO (?? , ?? ): Constructive interference (e.g., ?? 1s = ??
??
+??
??
).
– Antibonding MO (??
*
, ??
*
): Destructive interference (e.g., ??
*
1s = ??
??
-??
??
).
• Conditions for Orbital Combination :
– Similar energy .
– Same symmetry .
– Maximum overlap.
• Energy Order (O
2
, F
2
) : ?? 1s <??
*
1s <?? 2s <??
*
2s <?? 2p
??
<?? 2p
??
=?? 2p
??
<??
*
2p
??
=??
*
2p
??
<??
*
2p
??
.
• Bond Order :
1
2
(N
??
– N
??
), where N
??
= electrons in bonding MOs, N
??
= electrons in antibonding
MOs.
• Examples :
– H
2
: (?? 1s)
2
, bond order = 1, diamagnetic.
– He
2
: (?? 1s)
2
(??
*
1s)
2
, bond order = 0, unstable.
– O
2
: (?? 1s)
2
(??
*
1s)
2
(?? 2s)
2
(??
*
2s)
2
(?? 2p
??
)
2
(?? 2p
??
)
2
(?? 2p
??
)
2
(??
*
2p
??
)
1
(??
*
2p
??
)
1
, bond order = 2, par a-
magnetic.
• Magnetic Properties : Diamagnetic (all electrons paired), Par amagnetic (unpaired elec-
trons).
3
Page 4


Chemical Bonding and Molecular Structure Cheat
Sheet (Class 11 CBSE)
1 Introduction to Chemical Bonding
• Chemical Bond : Attr active force holding atoms, ions, etc., together in a molecule or com-
pound.
• Wh y Atoms Combine : T o achieve stability b y attaining noble gas electron configur ation
(lower energy state).
• K ey Theories :
– K össel-Lewis Approach
– V alence Shell Electron Pair Repulsion (VSEPR) Theory
– V alence Bond (VB) Theory
– Molecular Orbital (MO) Theory
2 K ö ssel-Lewis Approach
• Lewis S ymbols : Dots around element symbol represent valence electrons (e.g., Na· , Cl:· ).
• Octet Rule : At oms combine to achieve 8 electrons in their valence shell (noble gas config-
ur ation).
– Ionic Bond : Electron tr ansfer (e.g., Na? Na
+
+ e
-
; Cl + e
-
? Cl
-
; forms NaCl).
– Covalent Bond : Electron sharing (e.g., Cl
2
: Cl:Cl, sharing one electron pair).
• Limitations of Octet Rule :
– Incomplete Octet : e.g., LiCl, BeH
2
, BCl
3
(centr al atom < 8 electrons).
– Odd-Electron Molecules : e.g., NO , NO
2
(odd number of electrons).
– Expanded Octet : e.g., PF
5
, SF
6
(centr al atom > 8 electrons due to d-orbitals).
3 Types of Bonds
• Ionic (Electrovalent) Bond :
– F ormed b y electron tr ansfer between electropositive (low ionization enthalp y) and
electronegative (high electron gain enthalp y) elements.
– Stabilized b y lattice enthalp y (energy to separ ate 1 mole of ionic solid into gaseous ions,
e.g., NaCl: 788 kJ/mol).
• Covalent Bond :
– F ormed b y sharing electron pairs (single, double, or triple bonds).
– Single Bond : One shared pair (e.g., H
2
).
– Double Bond : Two shared pairs (e.g., O
2
, CO
2
).
1
– Triple Bond : Three shared pairs (e.g., N
2
, C
2
H
2
).
• Hydrogen Bond :
– W eak bond between H (bonded to F , O , N) and another electronegative atom (F , O , N).
– Intermolecular : Between different molecules (e.g., H
2
O , HF).
– Intr amolecular : Within the same molecule (e.g., o-nitrophenol).
– Stronger than van der W aals forces but weak er than covalent/ionic bonds.
4 Bond Par ameters
• Bond Length : Equilibrium distance between nuclei of bonded atoms (e.g., H
2
: 74 pm, C=C:
133 pm).
– Covalent r adius = half the distance between two similar atoms in a covalent bond.
• Bond Angle : Angle between orbitals containing bonding electron pairs (e.g., H
2
O: 104.5
?
).
• Bond Enthalp y : Energy to break 1 mole of a bond (e.g., H
2
: 435.8 kJ/mol, N
2
: 946 kJ/mol).
– Higher bond order? higher bond enthalp y , shorter bond length.
• Bond Order : Number of bonds between atoms (e.g., H
2
: 1, O
2
: 2, N
2
: 3).
• Dipole Moment (?? ) : Product of charge and distance (?? = ??× ?? , in Deb ye, 1D = 3.33564 ×
10
-30
C· m).
– Polar Covalent Bond : Unequal electron sharing (e.g., HF: H
+??
–F
-??
).
– Nonpolar Covalent Bond : Equal sharing (e.g., H
2
, O
2
).
– Net dipole depends on molecular geometry (e.g., CO
2
: linear , ?? = 0 ; H
2
O: bent, ?? = 1.85
D).
5 VSE PR Theory
• Principle : Electron pairs (bond pairs, lone pairs) around centr al atom repel each other ,
minimizing repulsion to determine molecular geometry .
• Repulsion Order : Lone pair-lone pair (lp-lp) > Lone pair-bond pair (lp-bp) > Bond pair-
bond pair (bp-bp).
• Geometries (No Lone Pairs) :
– 2 electron pairs: Linear (e.g., BeCl
2
, 180
?
).
– 3 electron pairs: Trigonal planar (e.g., BCl
3
, 120
?
).
– 4 electron pairs: T etr ahedr al (e.g., CH
4
, 109.5
?
).
– 5 electron pairs: Trigonal bip yr amidal (e.g., PCl
5
, 120
?
/90
?
).
– 6 electron pairs: Octahedr al (e.g., SF
6
, 90
?
).
• Geometries (With Lone Pairs) :
– AB
2
E: Bent (e.g., SO
2
,~ 119.5
?
).
– AB
3
E: Trigonal p yr amidal (e.g., NH
3
, 107
?
).
– AB
2
E
2
: Bent (e.g., H
2
O , 104.5
?
).
– AB
4
E: See-saw (e.g., SF
4
).
– AB
3
E
2
: T-shaped (e.g., ClF
3
).
2
6 V a lence Bond (VB) Theory
• Principle : Coval ent bond forms b y overlap of half-filled atomic orbitals with opposite spins.
• Types of Overlap :
– Sigma (?? ) Bond : Head-on overlap (s-s, s-p, p-p).
– Pi (?? ) Bond : Sidewa ys overlap (p-p).
– Sigma bonds are stronger due to greater overlap.
• Hybridisation : Mixing of atomic orbitals to form equivalent h ybrid orbitals.
– sp : Linear , 180
?
(e.g., BeCl
2
).
– sp
2
: Trigonal planar , 120
?
(e.g., BCl
3
).
– sp
3
: T etr ahedr al, 109.5
?
(e.g., CH
4
, NH
3
, H
2
O).
– sp
3
d : Trigonal bip yr amidal (e.g., PCl
5
, axial bonds longe r due to more repulsion).
– sp
3
d
2
: Octahedr al (e.g., SF
6
).
• Examples :
– C
2
H
6
: sp
3
-sp
3
?? bonds (C–C, C–H).
– C
2
H
4
: sp
2
-sp
2
?? bond (C–C), ?? bond (C=C), sp
2
-s ?? bonds (C–H).
– C
2
H
2
: sp-sp ?? bond (C–C), two ?? bonds (C= C), sp-s ?? bonds (C–H).
7 Molecular Orbital (MO) Theory
• Principle : Atomic orbitals combine to form molecular orbitals (bonding: lower energy;
antibonding: higher energy).
• Linear Combination of Atomic Orbitals (LCA O) :
– Bonding MO (?? , ?? ): Constructive interference (e.g., ?? 1s = ??
??
+??
??
).
– Antibonding MO (??
*
, ??
*
): Destructive interference (e.g., ??
*
1s = ??
??
-??
??
).
• Conditions for Orbital Combination :
– Similar energy .
– Same symmetry .
– Maximum overlap.
• Energy Order (O
2
, F
2
) : ?? 1s <??
*
1s <?? 2s <??
*
2s <?? 2p
??
<?? 2p
??
=?? 2p
??
<??
*
2p
??
=??
*
2p
??
<??
*
2p
??
.
• Bond Order :
1
2
(N
??
– N
??
), where N
??
= electrons in bonding MOs, N
??
= electrons in antibonding
MOs.
• Examples :
– H
2
: (?? 1s)
2
, bond order = 1, diamagnetic.
– He
2
: (?? 1s)
2
(??
*
1s)
2
, bond order = 0, unstable.
– O
2
: (?? 1s)
2
(??
*
1s)
2
(?? 2s)
2
(??
*
2s)
2
(?? 2p
??
)
2
(?? 2p
??
)
2
(?? 2p
??
)
2
(??
*
2p
??
)
1
(??
*
2p
??
)
1
, bond order = 2, par a-
magnetic.
• Magnetic Properties : Diamagnetic (all electrons paired), Par amagnetic (unpaired elec-
trons).
3
8 Resonance
• Concept : When a single Lewis structure cannot describe a molecule accur ately , multiple
structures (canonical forms) are used, and the a ctual structure is a resonance h ybrid.
• Examples :
– O
3
: Two structures with O–O single and O=O double bonds; h ybrid has equal bond
lengths (128 pm).
– CO
2-
3
: Three structures with one C=O double bond and two C–O single bonds; all C–O
bonds equivalent.
• K ey Points :
– Resonance lowers energy , stabilizing the molecule.
– Canonical forms are h ypothetical; the h ybrid is the actual structure.
9 F ormal Charge
• F ormula : FC = V alence electrons – Non-bonding electrons –
1
2
(Bonding electrons).
• Use : Helps select the most stable Lewis structure (lowest formal charges).
• Example (O
3
) :
– Centr al O: FC = 6 – 2 –
1
2
(6) = +1.
– End O (double bond): FC = 6 – 4 –
1
2
(4) = 0.
– End O (single bond): FC = 6 – 6 –
1
2
(2) = - 1.
10 K ey Examples
• Lewis Structures :
– H
2
S: H–S–H (bent, 2 lone pairs).
– CO
2-
3
: Resonance h ybrid of three structures.
– HCOOH: O=C(OH)–H (C has sp
2
, double bo nd with O).
• VSEPR Shapes :
– BeCl
2
: Linear (sp).
– NH
3
: Trigonal p yr amidal (sp
3
, 1 lone pair).
– H
2
O: Bent (sp
3
, 2 lone pairs).
• MO Configur ations :
– N
2
: Bond order = 3, diamagnetic.
– O
+
2
: Bond order = 2.5, par amagnetic.
– O
-
2
: Bond order = 1.5, par amagnetic.
4