© 2009, Prentice-Hall, Inc. Chapter 6 Concepts of Chemical Bonding Chemistry, The Central Science,...

Chapter 6Concepts of Chemical

Bonding

Chemistry, The Central Science, 11th editionTheodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten

John D. BookstaverSt. Charles Community College

Cottleville, MO

Chemical Bonding• Chemical Bond means the forces that hold

groups of atoms together and make them function as a unit.

• Bonds form because the energy of the system is lowered, making the molecule more stable.

• Bond Energy is the energy required to break a bond.

Types of chemical bonds• There are two types of chemical bonds:• Intramolecular Bonds (Interatomic Bonds) : holds atoms WITHIN a

molecule, much stronger than intermolecular bonds. There are four types of intramolecular bonds:– Ionic Bonds– Covalent Bonds (Polar Covalent, Non Polar Covalent Bonds and

Coordinate Covalent Bond)– Covalent Network Bonding– Metallic Bond

• Intermolecular Bonds: attractions BETWEEN two or more molecules already bonded by intramolecular bonds. They are of four types:-Hydrogen Bonds

-Dipole –Dipole Interactions-London Dispersion Forces-Ion- Molecule Interactions

Intramolecular Bonds

• Three basic types of bonds– Ionic

• Electrostatic attraction between ions

– Covalent• Sharing of electrons

– Metallic• Metal atoms bonded to

several other atoms

Ionic Bonding

It is due to the electrostatic attraction between a cation and an anion. Occurs between a metal and a nonmetalEN difference of atoms >1.7Solids

Energetics of Ionic Bonding

But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

How are Ionic Bonds Formed

• Ex. NaCl • Ex. CaCl2

Properties of Ionic Bonds• Crystal lattices (orderly

arrangement)• Hard, but brittle• High melting point• High boiling point• Do not conduct electricity

in solid form• Formula unit of an ionic

compound= molecule of a molecular (covalently bonded compound)

• Formula unit shows the ratio of ions present in a sample of ionic compound. Ex. Sod. Chloride’s formula unit is NaCl

Ionic Solids• What are the empirical formulas for these

compounds?– (a) Green: chlorine; Gray: cesium– (b) Yellow: sulfur; Gray: zinc– (c) Gray: calcium; Blue: fluorine

CsCl ZnS CaF2

(a) (b) (c)

Types of Ions: 1)Monatomic ions: One atom ions, Ex. Cations: Ca 2+, Mg2+, Na+, K+ Anions: Cl-, Br-, F-2)Polyatomic ions: Ions consisting of more than one atom. Ex: SO4

2-, NO3-

Criss Cross Method of Writing the Formula for Ionic Compounds

Covalent Bonding

• In covalent bonds atoms share electrons.

• There are several electrostatic interactions in these bonds:– Attractions between electrons

and nuclei– Repulsions between electrons– Repulsions between nuclei

Covalent bonding

A covalent bond is formed by a shared pair of electrons between two atoms. Between two non metals- usually liquids and gases•A group of atoms that are united by covalent bond is called a molecule and the substance made of molecules is called a molecular compound.•Molecular formula shows types and numbers of atoms tied together to form a molecule.•Formation of a covalent bond happens when there is net lowering of potential energy.•Bond length: It is the average distance between two bonded atoms.

• When forming, atoms must find the lowest energy bond length to minimize potential energy

Covalent Compund Formation

Types of Covalent BondsPolar , Non Polar covalent bonds and Coordinate Covalent BondsNon Polar Covalent Bond: e are shared equally by both nuclei equally,. Ex. H2, N2

Polar Covalent Bond: In polar covalent bond, the e tend to be more towards the atom having a higher electro negativity. Ex. H2O

Electro negativity difference and bonding: The difference in the electro negativity value of the two atoms will decide as to what kind of bond will be formed.•Electro negativity difference of 0.3 or less non polar covalent•Electro negativity difference of 0.3 – 1.6 polar covalent•Electro negativity difference of 1.7 or more Ionic bond

Polar Covalent Bonds

• Though atoms often form compounds by sharing electrons, the electrons are not always shared equally.

• Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.

• Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

Electronegativity

• Electronegativity is the ability of atoms in a molecule to attract electrons to themselves.

• On the periodic chart, electronegativity increases as you go…– …from left to right across a

row.– …from the bottom to the

top of a column.

• When two atoms share electrons unequally, a bond dipole results.

• The dipole moment is represented by an arrow pointing towards electronegative atom in the molecule.It is measured in debyes (D).

The greater the difference in electronegativity, the more polar is the bond.

Properties of covalent bonds: •Lower melting points than the ionic bond•Polar covalent bonds tend to have dipole moments and bond polarity.•Dipole Moment: It means that molecule has two poles, a positive and a negative pole.•The arrow showing dipole moment points to the center of negative charge and the tail is at the center of positive charge.•Example: HF CHCl3•It becomes more difficult to calculate dipole moment with three or more atoms in a molecule because individual dipole moments can cancel each other.Ex. CO2

Covalent Network Bond:In covalent network solid each atom is covalently bonded to its nearest neighboring atoms. The covalent bonding extends throughout a network that includes a very large number of atoms. Ex: diamond (Cx),

quartz(SiO2)x, silicon carbide(SiC)x. Such solids are

essentially giant molecules. The subscript x in these formulas indicates that the component with in the parentheses extends indefinitely. The network solids are hard and brittle (Why?). They have rather high mps.

Covalent-Network Solids

• Diamonds are an example of a covalent-network solid, in which atoms are covalently bonded to each other.– They tend to be hard and have high melting points.

Metallic Bond:The chemical bonding that results from the attraction between metal atoms an the surrounding sea of electrons is called as metallic bonding

Electron Sea Model for MetalsWith in a metal, the vacant orbitals in the atom’s outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. The freedom of motion in the network of metal atoms accounts for the high electrical and thermal conductivity characteristic of all the metals. This model also explains the malleability and ductility of the metals. (Why?)

Metallic Solids• Metals are not covalently

bonded, but the attractions between atoms are too strong to be van der Waals forces.

• In metals valence electrons are delocalized throughout the solid.

Ionic Solid Nonpolar Covalent Solids

Polar Covalent Solids

Metallic Solids

Covalent Network Solids

How is it formed?

Examples

Melting points and boiling points

Conductivity

Hardness

Brittle

Structural Particles

Forces Within Particles

Forces Between Particles

Comparison of Ionic Solids, Metallic Solids, Covalent Molecular Solids and Covalent Network Solids

Lewis Structures

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

Writing Lewis Structures

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.– If it is an anion, add one

electron for each negative charge.

– If it is a cation, subtract one electron for each positive charge.

5 + 3(7) = 26

2. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

Keep track of the electrons:

26 - 6 = 20

Writing Lewis Structures• Step 3

Step 3• Calculate number of bonds in the system:

Covalent bonds are made by sharing of e. You need 32 and you have 24. You are 8 e deficient. If you make 4 bonds ( with 2 e per bond) , you will make up the deficiency. Therefore,

• # of bonds= ( e in step 2- e in step 1)/2 =(32-

24)/2= 4 bondsStep 4• Draw the structure: The central atom is C

( usually the atom with least electro negativity will be in the center). The oxygens surround it . Because there are four bonds and only three atoms, there will be one double bond.

Step 5 • Double check your answer by counting total

number of electrons.

• Lewis structures are used to depict bonding pairs and lone pairs of electron in the molecule.

Step 1• Total number of valence electrons

in the system: Sum the number of valence electrons on all the atoms . Add the total negative charge if you have an anion. Subtract the charge if you have a cation.Example: CO3

2- Step 2• Number of electrons if each atom

is to be happy: Atoms in our example will need 8 e (octet rule) or 2 e ( hydrogen). So, for the ex.

Draw the Lewis structures for the following using above steps. Show work!A.Cl2B. CH2Cl23. NH3

Exceptions to octet rule: Expanded Octet ICl3

Exceptions to Octet: less than 8 e: H, He, Be, B

Formal Charges: Writing Lewis Structures

• Then assign formal charges.– For each atom, count the electrons in lone pairs and half

the electrons it shares with other atoms.– Subtract that from the number of valence electrons for that

atom: the difference is its formal charge.

• The best Lewis structure…– …is the one with the fewest charges.– …puts a negative charge on the most

electronegative atom.

Resonance

This is the Lewis structure we would draw for ozone, O3.

Resonance

• But this is at odds with the true, observed structure of ozone, in which…– …both O-O bonds are

the same length.– …both outer oxygens

have a charge of -1/2.

Resonance

• One Lewis structure cannot accurately depict a molecule like ozone.

• We use multiple structures, resonance structures, to describe the molecule.

Resonance

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

Resonance

• In truth, the electrons that form the second C-O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

• They are not localized; they are delocalized.

Resonance

• The organic compound benzene, C6H6, has two resonance structures.

• It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

Odd Number of Electrons

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

More Than Eight Electrons

• The only way PCl5 can exist is if phosphorus has 10 electrons around it.

• It is allowed to expand the octet of atoms on the 3rd row or below.– Presumably d orbitals in

these atoms participate in bonding.

Molecular Shapes

• The shape of a molecule plays an important role in its reactivity.

• By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

What Determines the Shape of a Molecule?

• Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

• By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

Electron Domains

• We can refer to the electron pairs as electron domains.

• In a double or triple bond, all electrons shared between those two atoms are on the same side of the central atom; therefore, they count as one electron domain.

• The central atom in this molecule, A, has four electron domains.

Valence Shell Electron Pair Repulsion Theory (VSEPR)

“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

Electron-Domain Geometries

These are the electron-domain geometries for two through six electron domains around a central atom.

Electron-Domain Geometries

• All one must do is count the number of electron domains in the Lewis structure.

• The geometry will be that which corresponds to the number of electron domains.

Molecular Geometries

• The electron-domain geometry is often not the shape of the molecule, however.

• The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

Molecular Geometries

Within each electron domain, then, there might be more than one molecular geometry.

Linear Electron Domain

• In the linear domain, there is only one molecular geometry: linear.

• NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

Trigonal Planar Electron Domain

• There are two molecular geometries:– Trigonal planar, if all the electron domains are bonding,– Bent, if one of the domains is a nonbonding pair.

Nonbonding Pairs and Bond Angle

• Nonbonding pairs are physically larger than bonding pairs.

• Therefore, their repulsions are greater; this tends to decrease bond angles in a molecule.

Multiple Bonds and Bond Angles

• Double and triple bonds place greater electron density on one side of the central atom than do single bonds.

• Therefore, they also affect bond angles.

Tetrahedral Electron Domain

• There are three molecular geometries:– Tetrahedral, if all are bonding pairs,– Trigonal pyramidal if one is a nonbonding pair,– Bent if there are two nonbonding pairs.

Trigonal Bipyramidal Electron Domain

• There are two distinct positions in this geometry:– Axial– Equatorial

Lower-energy conformations result from having nonbonding electron pairs in equatorial, rather than axial, positions in this geometry.

• There are four distinct molecular geometries in this domain:– Trigonal bipyramidal– Seesaw– T-shaped– Linear

Octahedral Electron Domain

• All positions are equivalent in the octahedral domain.

• There are three molecular geometries:– Octahedral– Square pyramidal– Square planar

Hybridization

• Refers to mixing of orbitals.

• Atomic orbitals of central atom undergo change to accommodate incoming atoms.

• Hybridization could be sp, sp2, sp3, sp3d and sp3d2.

• How do you tell the hybridization on the central atom?

Intermolecular Forces

The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

These intermolecular forces as a group are referred to as van der Waals forces.

van der Waals Forces

• Dipole-dipole interactions• Hydrogen bonding• London dispersion forces

Ion-Dipole Interactions

• Ion-dipole interactions (a fourth type of force), are important in solutions of ions.

• The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

Dipole-Dipole Interactions

• Molecules that have permanent dipoles are attracted to each other.– The positive end of one is

attracted to the negative end of the other and vice-versa.

– These forces are only important when the molecules are close to each other.

Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.

London Dispersion Forces

While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

• These forces are present in all molecules, whether they are polar or nonpolar.

• The tendency of an electron cloud to distort in this way is called polarizability.

Factors Affecting London Forces

• The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).

• This is due to the increased surface area in n-pentane.

Factors Affecting London Forces

• The strength of dispersion forces tends to increase with increased molecular weight.

• Larger atoms have larger electron clouds which are easier to polarize.

Which Have a Greater Effect?Dipole-Dipole Interactions or Dispersion Forces

• If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.

• If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

How Do We Explain This?

• The nonpolar series (SnH4 to CH4) follow the expected trend.

• The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

Hydrogen Bonding

• The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• We call these interactions hydrogen bonds.

Hydrogen Bonding

• Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.

Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

Summarizing Intermolecular Forces

Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.

Viscosity• Resistance of a liquid to

flow is called viscosity.• It is related to the ease

with which molecules can move past each other.

• Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Surface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

Phase Changes

© 2009, Prentice-Hall, Inc. Chapter 6 Concepts of Chemical Bonding Chemistry, The Central Science,...

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