Chemical BondingChapter 8 Sections 1-8

• A chemical bond is: a strong electrostatic force of attraction between atoms in a molecule or compound.

• Bonding between atoms occurs because it creates a more stable arrangement for the atoms.

Lewis Symbols – Dot Diagrams

• Convenient way to show the valence electrons

Three types of bonding• Metallic bonding – results from the attraction

between metal atoms and the surrounding sea of electrons

• Ionic bonding – results from the electrical attraction between positive and negative ions.

• Covalent bonding – results from the sharing of electron pairs between two atoms

Ionic Bonding• Many atoms transfer electrons and

other atoms accept electrons, creating cations (positive metal ions) and anions (negative nonmetal ions).

• The resulting ions are attracted to each other by electrostatic force.

Ionic Bonding• The ions closely pack together in a crystal

lattice.• This arrangement maximizes the attractive

forces among cations and anions while minimizing repulsive forces.

• Because force is proportional to the charge on each ion, larger charges lead to stronger interactions.

• Because force is inversely proportional to the square of the distance between the centers of the ions, smaller ions lead to stronger interactions.

Ionic bonding between Na and Cl

Sodium atom Chlorine atom 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p5 Sodium ion Na1+ Chlorine ion Cl1- 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p6

Covalent bonding

• In many cases electrons do not completely transfer from one atom to another.

• The electrons between atoms are shared.

Covalent bonding between H2

• Hydrogen’s electron configuration is 1s1

• Because both H atoms need 1 more electron to become isoelectronic with He, it is unlikely that either will give up an electron.

Covalent bonding between H2

↑ 1s

↓ 1s

They share the two electrons.H· + H · H : H

Types of Covalent Bonds• When electrons are shared equally the

bond is called a NONPOLAR covalent bond. (i.e. H2)

• Sometimes the electrons between two atoms are NOT shared equally. The bond created is called a POLAR covalent bond.

. . . .

H· + ·Cl: H:Cl:

. . . .

Polar Covalent Bonding

• An example of this would be HCl.

HCl moleculeHydrogen atom

[Ne] ↑↓ ↑↓ ↑↓ ↓ 3s 3p

↑ 1s

Chlorine atom

How to classify bond types• Electronegativity – ability of an atom in a

molecule to attract shared electrons to it• Each element on the periodic table is assigned

an electronegativity value (see page 353) that ranges from 0.7 to 4.0.

• The difference in the electronegativity determines the bonding type (ionic, polar covalent, or nonpolar covalent).

Electronegativity Values

If the electronegativity difference is:

1.7 and higher ionic

0.3 to 1.7 polar covalent

0.0 to 0.3 nonpolar covalent

What if I get an electronegativity difference that is 0.3 or 1.7?

• These cut-off numbers are guidelines. • It is a gradual change not stair-step.

Ionic Character

• As the electronegativity difference increases, the ionic character increases as well!

Practice ProblemsWhat type of bond will occur between

iodine and the following elements: cesium, iron, and sulfur?

Bonding between I and:

Electronegativity difference

Bond Type

Cesium

Iron

Sulfur

Determine the type of bond between the following pairs.

Bonding between

Electronegativity difference

Bond type

Li & Cl

S & O

Ca & Br

P & H

Si & Cl

S & Br

Other ways to determine bonding

• Electronegativity is not the only factor in determining bonding.

• Generally, bonds between a metal and nonmetal are ionic, and between two nonmetals the bonds are covalent.

• Examination of the properties of a compound is the best way to determine the type of bonding.

Ionic Bonding

• Ionic compounds are formed to maximize stability.

• Nonmetal - will gain electrons to become isoelectronic with nearest noble gas; called an anion

• Metal – will lose electrons to become isoelectronic with noble gas; called a cation

Transition Metals

ZincElectron configuration is 1s22s22p63s23p64s23d10

When it forms the +2 ion, it loses the 2 valence electrons in the 4s sublevel.

Zn2+ configuration is 1s22s22p63s23p63d10

Practice Problems

• Write the electron configurations for the following ions.– Fe2+

– S2-

– Mg2+

• Use electron configurations to explain why the most probable charge on the strontium ion is +2.

Li,152 pm3e and 3p

Li+, 60 pm2e and 3 p

+Does the size goDoes the size goup or down up or down when losing an when losing an electron to form electron to form a cation?a cation?

Size of Ions

• CATIONSCATIONS are are SMALLERSMALLER than the atoms than the atoms from which they come.from which they come.

• The electron/proton attraction has The electron/proton attraction has gone UP and so size gone UP and so size DECREASESDECREASES..

Li,152 pm3e and 3p

Li +, 78 pm2e and 3 p

+Forming Forming a cation.a cation.Forming Forming a cation.a cation.

Size of Ions

F,64 pm9e and 9p

F- , 136 pm10 e and 9 p

-Does the size go up or down Does the size go up or down when gaining an electron when gaining an electron to form an anion?to form an anion?

Does the size go up or down Does the size go up or down when gaining an electron when gaining an electron to form an anion?to form an anion?

Size of Ions

• ANIONSANIONS are are LARGERLARGER than the atoms from which than the atoms from which they come.they come.

• The electron/proton attraction has gone DOWN and The electron/proton attraction has gone DOWN and so size so size INCREASESINCREASES..

• Trends in ion sizes are the same as atom sizes. Trends in ion sizes are the same as atom sizes.

Forming Forming an anion.an anion.Forming Forming an anion.an anion.F, 71 pm

9e and 9pF-, 133 pm10 e and 9 p

-

Size of Ions

Trends in Ion SizesTrends in Ion SizesTrends in Ion SizesTrends in Ion Sizes

Figure 8.13Figure 8.13

Which is Bigger?

• Cl or Cl- ?• K+ or K ?• Ca or Ca+2 ?• I- or Br- ?

Which is Bigger?

• Cl or Cl- ? Cl- • K+ or K ? K• Ca or Ca+2 ? Ca• I- or Br- ? I-

Lattice Energy Effects

• 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.

Lattice Energy

• To determine which compound will have the highest lattice energy, take into consideration the following:

– The size of the ions in the compound• The smaller the size, the greater the

lattice energy– The charge of the ions in the compound

• The greater the charge, the greater the lattice energy

Calculating ∆Hf• We can take advantage of the fact the

energy is a state function and break the reaction into steps, the sum of which is the overall reaction.

• Let’s do #41 Na(s) + ½ Cl2 (g) NaCl(s)

Given the following:Lattice energy -786 kJ/molIonization energy for Na 495 kJ/molElectron affinity for Cl -349 kJ/molBond energy of Cl2 239 kJ/mol

Enthalpy sublimation for Na 109 kJ/mol

ProcessStep 1: Sublimation of Na

Na(s) Na(g) 109 kJ/molStep 2: Ionization of Na

Na (g) Na+ (g) + e- 495 kJ/molStep 3: Dissociation of Cl2

½ Cl2 (g) Cl(g) 119.5 kJ/mol

Step 4: Formation of Cl- (Electron Affinity)Cl (g) + e- Cl-(g) -349 kJ/mol

Step 5: Formation of NaCl Na+(g) + Cl-(g) NaCl(s) -786 kJ/mol

Na(s) + ½ Cl2 (g) NaCl(s) -411.5 kJ/mol

Drawing Lewis Structures Valence Electron Review

• Valence electrons are in outermost level• You can use periodic table or electron

configuration to determine valence electronsExample: Phosphorus

–Located in Group 15 or 5A–Electron configuration is 1s22s22p63s23p3

–Contains 5 valence electrons

Complete Exercise 1 on worksheet

Drawing Lewis StructuresOctet Rule

• Most useful rule for creating Lewis structures• Every atom usually has 8 valence electrons• Exception: hydrogen is good with 2 (like He)• Lines are used to link atoms together (same as

using 2 dots)

Same as

Steps to Drawing Lewis Structures1. Count valence electrons.2. Connect atoms together with bonds. In

molecules with a single atom of one element and several atoms of another element, the single atom is generally in the center with the other atoms attached to it.

3. Add electrons around outside of atoms to give each atom 8 electrons (or 2 in the case of hydrogen).

4. Count electrons used. This number must be the same as valence electrons.

PCl3

Bonding Pairs

Lone Pairs(a.k.a. nonbonding electrons)

5+(3*7)=26 e-

Complete Exercise 2.

Helpful Hints• Carbon atoms form 4 bonds.• Nitrogen atoms form 3 bonds.• Oxygen atoms form 2 bonds.• Hydrogen atoms form 1 bond.• Fluorine atoms form 1 bond.• Other halogens (Cl, Br, and I) frequently form 1

bond (but not always).

Determining the Central Atom• In a molecule, the atom that typically forms the

greatest number of bonds is in the center, with other atoms attached to it.

• Example: CH3Cl– Carbon forms 4 bonds– Hydrogen forms 1 bond– Chlorine forms 1 bondSO CARBON IS IN THE MIDDLE WITH HYDROGEN AND

CHLORINE AROUND IT! Don’t forget electrons on chlorine to make 8!

• Complete exercise 3.

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).

• Generally, bond distances shorten with multiple bonding.

H H O O N N

Covalent BondingCovalent Bonding

Octet in each case

Resonance

• Occurs when more than one valid Lewis structure can be written for a particular molecule.

Odd Number of Electrons…

NO Number of valence electrons = 11

N O N O

NO2Number of valence electrons = 17

Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons)

Molecules and atoms which are neutral (contain no formal charge) and with an

unpaired electron are called Radicals

N OO N OO N OO

Beyond the Octet

• Elements in the 3rd period or higher can have more than an octet if needed.

• Atoms of these elements have valence d orbitals, which allow them to accommodate more than eight electrons.

More than an Octet…

PCl5

Elements from the 3rd period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding

P : (Ne) 3s2 3p3 3d0

Number of valence electrons = 5 + (5 x 7) = 40 P

Cl

ClCl

ClCl

SF4S : (Ne) 3s2 3p4 3d0

Number of valence electrons = 6 + (4 x 7) = 34

SF

F

F

FThe Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.

Formal ChargeDifference between the # of valence electrons in the free atom and the # of electrons assigned to that atom in the Lewis structure.

FC = formal charge; G.N. = Group Number#BE = bonding electrons; #LPE = lone pair electrons

If Step 4 leads to a positive formal charge on an inner atom beyond the second row, shift electrons to make double or triple bonds to minimize formal charge, even if

this gives an inner atom with more than an octet of electrons.

LPEBE ##

2

1- G.N. FC

Formal Charge

Not as good Better

CO O(-1) (0) (+1)

CO O(0) (0) (0)

Molecular ShapesMolecular Shapes• Lewis structures give atomic connectivity:

they tell us which atoms are physically connected together. They do not tell us the shape.

• The shape of a molecule is determined by its bond angles.

• Consider CCl4: experimentally we find all Cl-C-Cl bond angles are 109.5.Therefore, the molecule cannot be planar.All Cl atoms are located at the vertices of a

tetrahedron with the C at its center.

Molecular Shape of CClMolecular Shape of CCl44

VSEPR TheoryVSEPR Theory

In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimized this repulsion.

We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

Why is VSEPR Theory Important?• Gives a specific shape due to the number

of bonded and non-bonded electron pairs in a molecule

• Tells us the actual 3-D structure of a molecule

• In bonding, electron pairs want to be as far away from each other as possible.

VSEPR and Resulting GeometriesVSEPR and Resulting Geometries

How does VSEPR THEORY work?We can use VSEPR theory using 4 steps.1. Draw the Lewis Structure for the

molecule.Example: SiF4

F

F-Si-F

F

How does VSEPR THEORY work?

We can use VSEPR theory using 4 steps1. Draw the Lewis Structure for the

molecule.2. Tally the number of bonding pairs and

lone (non-bonding) pairs on the center atom.

F

F-Si-F

F

Bonding pairs: 4

Lone pairs on central atom: 0

How does VSEPR THEORY work?We can use VSEPR theory using 4

steps1. Draw the Lewis Structure for the

molecule2. Tally the number of bonding pairs

and lone pairs on the center atom.3. Arrange the rest of the atoms so

that they are as far away from each other as possible.

Si

F

FF F

How does VSEPR THEORY work?We can use VSEPR theory using 4 steps1. Draw the Lewis Structure for the molecule2. Tally the number of bonding pairs and

lone pairs on the center atom.3. Arrange the rest of the atoms so that they

are as far away from each other as possible

4. Give the type of geometry the molecule has:

Tetrahedral

Another Example:To determine the electron pair geometry:

1) draw the Lewis structure;2) count the total number of electron pairs around the central atom.3) arrange the electron pairs in one of the geometries to minimize e-e repulsion.4) multiple bonds count as one bonding pair for VSEPR

The VSEPR ModelThe VSEPR Model

Predicting Molecular GeometriesPredicting Molecular Geometries

The VSEPR ModelThe VSEPR Model

Predicting Molecular GeometriesPredicting Molecular Geometries

The VSEPR ModelThe VSEPR Model

Difference between geometry and shapeDifference between geometry and shapeGeometry:We determine the geometry only looking at electrons.All the atoms that obey the octet rule have the same tetrahedral-like geometry.

Shape:We name the shape by the positions of atoms.We ignore lone pairs in the shape.

The VSEPR ModelThe VSEPR Model

Predicting ShapePredicting Shape

Shape

The VSEPR ModelThe VSEPR Model

Predicting ShapePredicting Shape

Shape

The VSEPR ModelThe VSEPR ModelThe Effect of Nonbonding Electrons and The Effect of Nonbonding Electrons and Multiple Bonds on Bond AnglesMultiple Bonds on Bond AnglesBy experiment, the H-X-H bond angle decreases on moving from C to N to O:

Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs.Therefore, the bond angle decreases as the number of lone pairs increase.

OHH

104.5O107O

NHH

HC

H

HHH109.5O

The VSEPR ModelThe VSEPR Model

The Effect of Nonbonding Electrons and The Effect of Nonbonding Electrons and Multiple Bonds on Bond AnglesMultiple Bonds on Bond AnglesSimilarly, electrons in multiple bonds repel more than electrons in single bonds.

C OCl

Cl111.4o

124.3o

The VSEPR ModelThe VSEPR Model

Molecules with Expanded Valence ShellsMolecules with Expanded Valence ShellsAtoms that have expanded octets have AB5 (trigonal bipyramidal) or AB6 (octahedral) electron pair geometries.

Examples:PF5 trigonal bipyramidal SCl6 octahedral

The VSEPR ModelThe VSEPR Model

Molecules with Expanded Valence ShellsMolecules with Expanded Valence Shells

The VSEPR ModelThe VSEPR Model

Molecules with Expanded Valence ShellsMolecules with Expanded Valence Shells

The VSEPR ModelThe VSEPR ModelMolecules with More than One Central AtomMolecules with More than One Central AtomIn acetic acid, CH3COOH, there are three central atoms.We assign the geometry about each central atom separately.

Hybrid Orbitals• In bonding, s and p orbitals are used in

bonding. It is easy to tell which ones are used by looking at our molecule.

• For example, CH4. Looking again at the Lewis structure, we see that there are 4 bonds. We call this sp3 hybridized.

Hybrid Orbitals

• Regions of electron density-EACH BOND AND LONE PAIR OF ELECTRONS ON THE CENTRAL ATOM IS KNOWN AS A REGION OF ELECTRON DENSITY.

• 2 regions of electron density-sp hybridized• 3 regions of electron density-sp2 hybridized• 4 regions of electron density-sp3 hybridized

Hybridizationspsp Hybrid Orbitals Hybrid OrbitalsThe two lobes of an sp hybrid orbital are 180 apart.

Hybrid Orbitals

spsp22 Hybrid Orbitals Hybrid OrbitalsImportant: when we mix n atomic orbitals we must get n hybrid orbitals.sp2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.)The large lobes of sp2 hybrids lie in a trigonal plane.All molecules with trigonal planar electron pair geometries have sp2 orbitals on the central atom.

Hybridization

Hybridization

spsp33 Hybrid Orbitals Hybrid Orbitalssp3 Hybrid orbitals are formed from one s and three p orbitals. Therefore, there are four large lobes.Each lobe points towards the vertex of a tetrahedron.The angle between the large lobes is 109.5All molecules with tetrahedral electron pair geometries are sp3 hybridized.

Hybridization

Hybrid OrbitalsHybrid Orbitals

Hybrid OrbitalsHybrid OrbitalsSummarySummaryTo assign hybridization:

1. Draw a Lewis structure.2. Assign the geometry using VSEPR theory.3. Use the geometry to determine the hybridization.4. Name the shape by the positions of the atoms.

Hybridization and Multiple Bonds•Multiple bonds overlap differently and are called -bonds and -bonds

•All single bonds are •Double bonds contain 1 and 1 bond

•Triple bonds contain 1 and 2 bonds

Bond Energy

Covalent Bonding & Orbital Overlap

•As two nuclei approach each other their atomic orbitals overlap.•As the amount of overlap increases, the energy of the interaction decreases.•At some distance the minimum energy is reached.•The minimum energy corresponds to the bonding distance (or bond length).

Covalent Bonding & Orbital Overlap

•As the two atoms get closer, their nuclei begin to repel and the energy increases.

•At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).

Bond Energies• Bond breaking requires energy (endothermic).• Bond formation releases energy (exothermic).H = D(bonds broken) D(bonds formed)

energy required energy released

Bond Energies