Bonding: General Concepts

Bonding: General Concepts
Chapter 8

Overview Types of chemical bonds, Electronegativity, Bond polarity and Dipole Moments. The Ions: electron configurations, size, formula, lattice energy calculations. Covalent bonds: model, bond energies, chemical reactions. Lewis structure, exceptions to the octet rule, resonance. Molecular structure models from Valence Shell Electron Pair Model “VSEPR” for single and multiple bonds.

Bonds Forces that hold groups of atoms together and make them function as a unit. NaCl – attraction is electrostatic since Na+ and Cl- are the Stable forms for these elements. This is an example of “Ionic Bonding”

Bond Energy It is the energy required to break a bond.
It gives us information about the strength of a bonding interaction, as well as radius.

Bond Length The distance where the system energy is a minimum.

Figure 8. 1: (a) The interaction of two hydrogen atoms
Figure 8.1: (a) The interaction of two hydrogen atoms. (b) Energy profile as a function of the distance between the nuclei of the hydrogen atoms.

Change in electron density as two hydrogen atoms approach each other.

Covalent Bond No electron transfer
Electrons are shared between two atoms, positioned between the two nuclei Example: H2, O2, H2O, CO2, etc.

Ionic Bonds Formed from electrostatic attractions of closely packed, oppositely charged ions. Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6
[Ne] Li Li+ + e- e- + F - F - Li+ + Li+

Ionic Compound Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed.

Ionic Bonds and Coulomb’s Law
Q1 and Q2 = numerical ion charges r = distance between ion centers (in nm) Attractive forces are (-), repulsive are (+).

Covalent or Ionic? Covalent and ionic are simply extreme cases.
Most molecules share electrons but “Unequally” due to the difference in electronegativity and electron affinity. This will give rise to a dipole moment and the molecule becomes polar.

Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling simple model Δ= (H  X)actual  (H  X)expected (H-H) + (X-X) 2 (H-X)experimental # (H-X)expected = If  = 0 => no polarity

Classification of bonds by difference in electronegativity
Bond Type 0 to 0.1 Covalent  2 Ionic 0 < and <2 Polar Covalent Increasing difference in electronegativity Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e-

Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent 9.5

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms. The molecule is called “Dipolar”. H F electron rich region electron poor region e- poor e- rich F H d+ d- Dipole Moment

Figure 8.2: The effect of an electric field on hydrogen fluoride molecules.
When no electric field is present, the molecules are randomly oriented. (b) When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their positive ends toward the negative pole.

Comparison of Ionic and Covalent Compounds

Polyatomic Molecules May exhibit dipole moment depending on their structure i.e. arrangement in space

Figure 8. 4: (a) The charge distribution in the water molecule
Figure 8.4: (a) The charge distribution in the water molecule. (b) The water molecule in an electric field. V-Shape

Figure 8.5: (a) The structure and charge distribution of the ammonia molecule. The polarity of the N—H bonds occurs because nitrogen has a greater electronegativity than hydrogen. (b) The dipole moment of the ammonia molecule oriented in an electric field. Look for a “NET DIPOLE”

Look for a “NET DIPOLE” N H H H Trigonal Pyramidal Structure

The carbon dioxide molecule CO2: The opposed bond polarities cancel out, and the carbon dioxide has no dipole moment: Non-polar molecule Note: The C-O bond is polar, but the net dipole is Zero Example: SO3, CCl4, etc.

Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4 O H S O dipole moment polar molecule dipole moment polar molecule C H C O no dipole moment nonpolar molecule no dipole moment nonpolar molecule

Dipoles (polar molecules) and Microwaves

Compounds Two nonmetals react: They share electrons to achieve NGEC (Noble Gas Electron Configurations) and form covalent bonds. A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied (cation) to achieve NGEC. The valence electron configuration of the nonmetal (anion) achieves NGEC.

Ions Ionic compounds are always electrically neutral e.g. they have the same amount of +ve and –ve charges. Common ions have noble gas configurations.

Electron Configurations of Cations and Anions
Of Representative Elements Na [Ne]3s1 Na+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. Ca [Ar]4s2 Ca2+ [Ar] Al [Ne]3s23p1 Al3+ [Ne] H 1s1 H- 1s2 or [He] Atoms gain electrons so that anion has a noble-gas outer electron configuration. F 1s22s22p5 F- 1s22s22p6 or [Ne] O 1s22s22p4 O2- 1s22s22p6 or [Ne] N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Ground State Electron Configurations of the Elements
ns2np6 Ground State Electron Configurations of the Elements ns1 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 d10 d1 d5 4f 5f

Cations and Anions Of Representative Elements
+1 +2 +3 -3 -2 -1

Notes on Ions Hydrogen may form H+ or H- Tin forms Sn2+ and Sn4+.
Transition metals exhibit a more complicated behavior.

Example of Ionic Compounds
MgO magnesium oxide is formed of Mg2+ and O2-. CaO formed from Ca2+ and O2-. Al2O3 is formed of 2Al3+ and 3O2-.

Figure 8.7: Sizes of ions related to positions of the elements on the periodic table.

Isoelectronic Ions Contain the the same number of electrons
8O 1s22s22p4 O2- 1s22s22p6 9F 1s22s22p5 F- 1s22s22p6 11Na 1s22s22p63s1 Na+ 1s22s22p6 12Mg 1s22s22p63s2 Mg2+1s22s22p6 13Al 1s22s22p63s23p1 Al3+ 1s22s22p6 Which one you expect to have the smallest radius? And why?

Radii of Isoelectronic Ions
O2> F > Na+ > Mg2+ > Al3+ largest smallest 13 protons vs. 10 electrons

Example Choose the largest ion in each of the following groups:
Li+, Na+, K+, Rb+, Cs+ Ba2+ , Cs+ , I- , Te2-

To Reiterate Cation is smaller than parent molecule.
Anion is larger than parent molecule. Size increase down in a group (+ve or –ve). The larger the mass number the smaller the size for isoelectronic cations and anions.

Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s23d6 Mn: [Ar]4s23d5 Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5 Fe3+: [Ar]4s03d5 or [Ar]3d5

Why Compounds Exist? The driving force behind the naturally occurring compounds (such as NaCl, H2O, etc.) is to yield a stable lower energy form. A stable form is a an arrangement of atoms, held together by bonding that prevent decomposition. This bonding energy when ions condense from gas phase into ionic solid is called Lattice Energy.

Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M+(g) + X(g)  MX(s) Lattice energy is negative (exothermic) from the point of view of the system.

Formation of an Ionic Solid
1. Sublimation of the solid metal M(s)  M(g) [endothermic] 2. Ionization of the metal atoms M(g)  M+(g) + e [endothermic] 3. Dissociation of the nonmetal 1/2X2(g)  X(g) [endothermic] 4. Formation of X ions in the gas phase: X(g) + e  X(g) [exothermic] 5. Formation of the solid (LATTICE) MX M+(g) + X(g)  MX(s) [quite exothermic] Lattice Energy

The energy changes involved in the formation of solid lithium fluoride from its elements.
From Gas to Solid Lattice Energy

The structure of lithium fluoride
The structure of lithium fluoride. Called also the NaCl structure where each ion is surrounded by 6 of the other ions Applicable for all alkali-metals/halogen except the Cesium salts.

Q1, Q2 = charges on the ions r = shortest distance between centers of the cations and anions The magnitude of Q’s will determine how strong is the lattice.

Comparison of the energy changes involved in the formation of solid sodium fluoride and solid magnesium oxide. Note all ions are isoelectronic

Electrostatic (Lattice) Energy
Lattice energy (E) increases as Q increases and/or as r decreases. Cmpd lattice energy MgF2 MgO LiF LiCl 2957 3938 1036 853 Q= +2,-1 Q= +2,-2 r F < r Cl

Partial Ionic Character of the Covalent Bond
Introduce the concept of percent ionic character in a polar covalent bond % ionic character = Measured dipole of X-Y Calculated dipole of X+Y- x 100

Ion Pairing Experiments showed that none of the studied systems have
100% ionic character despite large differences in electronegativity. The reason is “Ion Pairing” during motion, ions approach closely and for a short period of time their charges will be neutralized before moving away from each other.

Molten NaCl conducts an electric current, indicating the presence of mobile Na+ and Cl- ions.
The more mobile the ions the Stronger the current and the brighter the lamp!

The relationship between the ionic character of a covalent bond and the electronegativity difference of the bonded atoms. Note: These are “Molten” salts and not “Aqueous” salts! <50% Non-Ionic >50% Ionic

A Model for Covalent Bond
Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

Fundamental Properties of Models
A model does not equal reality. Models are oversimplifications, and are therefore often wrong. Models become more complicated as they age. We must understand the underlying assumptions in a model so that we don’t misuse it.

e.g. CH4 is 1652 KJ lower in energy than 1 mole of C and 4 mole of H.
Generally, bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms e.g. CH4 is 1652 KJ lower in energy than 1 mole of C and 4 mole of H. Stability can be determined in terms if model called “Chemical Bond” C-H bond energy = = 413 KJ/mol 1652 KJ/mol 4

Single bond < Double bond < Triple bond
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Energy H2 (g) H (g) + DH0 = kJ Cl2 (g) Cl (g) + DH0 = kJ HCl (g) H (g) + Cl (g) DH0 = kJ O2 (g) O (g) + DH0 = kJ O N2 (g) N (g) + DH0 = kJ N Bond Energies Single bond < Double bond < Triple bond

Using this model, one can determine other bond energies.
CH3Cl is composed of 3 C-H bond and 1 C-Cl bond C-H bond is known 413 KJ/mol C-Cl bond energy = 1572 – 3(413) = 339 KJ/mol

Average bond energy in polyatomic molecules
H2O (g) H (g) + OH (g) DH0 = 502 kJ OH (g) H (g) + O (g) DH0 = 427 kJ Average OH bond energy = 2 = 464 kJ

Bond Energy and Bond Length
Note : The shorter the bond the higher the bond energy

Bond Energies and Enthalpy of Reaction
Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic). H = D(bonds broken)  D(bonds formed) Energy required Energy released Reactants Products

Use bond energies to calculate the enthalpy change for:
H2 (g) + F2 (g) HF (g) DH0 = SBE(reactants) – SBE(products) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) H 1 436.4 F 1 156.9 Type of bonds formed Number of bonds formed Bond energy (kJ/mol) Energy change (kJ) H F 2 568.2 1136.4 DH0 = – 2 x = kJ 9.10

Localized Electron Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Localized Electron Model
1. Description of valence electron arrangement (Lewis structure). 2. Prediction of geometry (VSEPR model). 3. Description of atomic orbital types used to share electrons or hold lone pairs.

Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration.

Writing Lewis Structures
Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge. Complete an octet for all atoms except hydrogen If structure contains too many electrons, form double and triple bonds on central atom as needed.

5 + (3 x 7) = 26 valence electrons
Write the Lewis structure of nitrogen trifluoride (NF3). Step 1 – N is less electronegative than F, put N in center 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 Actual in 3D F N

Lewis structure of F2 F F lone pairs single covalent bond

Lewis structure of water
single covalent bonds 2e- 8e- 2e- H + O + H O H or Double bond – two atoms share two pairs of electrons 8e- 8e- 8e- double bonds O C or O C double bonds Triple bond – two atoms share three pairs of electrons triple bond 8e- N 8e- or N triple bond

4 + (3 x 6) + 2 = 24 valence electrons
Write the Lewis structure of the carbonate ion (CO32-). Step 1 – C is less electronegative than O, put C in center 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) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24 O C

Carbon Dioxide CO2 or O C .. O C .. O C .. Octet No octet yet BUT
Do not Break Symmetry O C ..

Exceptions to the Octet Rule
The Incomplete Octet Be – 2e- 2H – 2x1e- 4e- BeH2 H Be B – 3e- 3F – 3x7e- 24e- 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24 F B BF3

Exceptions to the Octet Rule
Odd-Electron Molecules N – 5e- O – 6e- 11e- NO N O The Expanded Octet (central atom with principal quantum number n > 2) S F S – 6e- 6F – 42e- 48e- 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 SF6

Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive (BF3, BeH2, etc..). 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals (SF6, PCl5, etc.). When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Note: For molecules that has several Third row (or higher) elements the extra electrons should be placed on the central atom. I3- I I I ..

Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures. The value of the resonance bond is in between a single bond and double bond. Bond Energy C=C > C C > C-C ….

Resonance Structure of Ozone: O3
+ - O + - What are the resonance structures of the carbonate (CO32-) ion? O C - O C - O C -

How the charges in the previous examples were depicted?
Calculate the Formal charge on each atom.

Formal Charge 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.

Two possible skeletal structures
of formaldehyde (CH2O) H C O H C O Calculate and minimize Formal Charges

( ) - -1 +1 H C O = 1 2 = 4 - 2 - ½ x 6 = -1 = 6 - 2 - ½ x 6 = +1
C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 H C O 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 formal charge on C = 4 - 2 - ½ x 6 = -1 formal charge on O = 6 - 2 - ½ x 6 = +1

( ) - H C O = 1 2 = 4 - 0 - ½ x 8 = 0 = 6 - 4 - ½ x 4 = 0 C – 4 e-
H C O C – 4 e- O – 6 e- 2H – 2x1 e- 12 e- 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 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 formal charge on C = 4 - 0 - ½ x 8 = 0 formal charge on O = 6 - 4 - ½ x 4 = 0

Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. Lewis structures with large formal charges are less plausible than those with small formal charges. 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 -1 +1 H C O

Formal Charge Not as good Better Charges are minimized

VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions.

Predicting a VSEPR Structure
1. Draw Lewis structure. 2. Put pairs as far apart as possible. 3. Determine positions of atoms from the way electron pairs are shared. 4. Determine the name of molecular structure from positions of the atoms.

Valence shell electron pair repulsion (VSEPR) model:
Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear B

0 lone pairs on central atom 2 atoms bonded to central atom
Cl Be 0 lone pairs on central atom 2 atoms bonded to central atom

Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear trigonal planar trigonal planar AB3 3

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral trigonal bipyramidal trigonal bipyramidal AB5 5

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral AB5 5 trigonal bipyramidal AB6 6 octahedral octahedral

Molecular structure of PCl6-
Octahedral Geometry

Octahedral electron arrangement for Xe

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal planar trigonal planar AB3 3 trigonal planar AB2E 2 1 bent

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB4 4 tetrahedral tetrahedral trigonal pyramidal AB3E 3 1 tetrahedral

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB4 4 tetrahedral tetrahedral AB3E 3 1 tetrahedral trigonal pyramidal AB2E2 2 2 tetrahedral bent H O V-Shape

bonding-pair vs. bonding
pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding <

See-Saw VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 trigonal
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 trigonal bipyramidal distorted tetrahedron AB4E 4 1 See-Saw

VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1 trigonal bipyramidal distorted tetrahedron trigonal bipyramidal AB3E2 3 2 T-shaped Cl F

VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1 trigonal bipyramidal distorted tetrahedron AB3E2 3 2 trigonal bipyramidal T-shaped trigonal bipyramidal AB2E3 2 3 linear I

Three possible arrangements of the electron pairs in the I3- ion.
Least Lone pair repulsions

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB6 6 octahedral square pyramidal Br F AB5E 5 1 octahedral

VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB6 6 octahedral AB5E 5 1 octahedral square pyramidal square planar Xe F AB4E2 4 2 octahedral

Possible electron-pair arrangements for XeF4.

Predicting Molecular Geometry
Draw Lewis structure for molecule. Count number of lone pairs on the central atom and number of atoms bonded to the central atom. Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO2 and SF4? S F S O AB4E AB2E distorted tetrahedron bent

The molecular structure of methanol CH3OH
The molecular structure of methanol CH3OH. (a) The arrangement of electron pairs and atoms around the carbon atom. (b) The arrangement of bonding and lone pairs around the oxygen atom. (c) The molecular structure.

Links http://www.molecules.org/VSEPR_table_c.html

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ANSWER HMCLASS PREP: Table 8.1

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ANSWER HMCLASS PREP: Figure 8.3

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ANSWER HMCLASS PRESENT: Animation: Electron Pair Repulsion, Four Pairs
HMCLASS PREP: Figure 8.16

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ANSWER HMCLASS PREP: Table 8.6

Bonding: General Concepts

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