Chapter 6 Chemical Bonding. Types of Chemical Bonding Chemical bond – Mutual electrical attraction...

Chapter 6

Chemical Bonding

Types of Chemical Bonding

Chemical bond– Mutual electrical attraction between nuclei

and valence electrons of different atoms that binds the atoms together

– Creates more stable compoundsIonic bonds

– Electrical attraction between large numbers of cations and anions

Covalent bonds– Sharing electron pairs between two atoms

Types of Chemical Bonds

Determine if bond is ionic or covalent by difference in electronegativities– Electronegativity difference 1.7 or less is

covalentBonding between diatomic molecule is

covalentNon-polar covalent - e- shared equally (≤0.3)Polar covalent - unequal attraction for e-

(0.3 - 1.7)– Electronegativity of 1.7 to 3.3 is ionic

Polar - uneven distribution of charge

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Predicting Bonds

4

Problems

Use electronegativity differences to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?

Problems

Use electronegativity differences to classify bonding between chlorine, Cl, and the following elements: calcium, Ca; oxygen, O; and bromine, Br. Indicate the more negative atom in each pair.

Covalent Bonding

Molecule– Neutral group of atoms that are covalently

bonded together– H2O, sugar,O2

Chemical formula– Shows relative number of atoms in a

compoundDiatomic Molecule

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Formation of a Covalent Bond

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Energy and Stability

Energy released when covalent bond formed

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Covalent Bonding

Bond length– Distance between two bonded atoms at

minimum potential energyAtoms vibrate back and forthDepends on atoms that have

combinedBond energy

– Energy required to break a chemical bond

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Polarity and Bond Strength

The greater the electronegativity difference, the greater the polarity and the stronger the bond

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Problem

Arrange the following bonds in order of increasing bond length, from shortest bond to longest

Bond Bond Energy (kJ/mol)

H-F 569H-I 299

H-Cl 432H-Br 366

Octet Rule

Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

Atoms with 8 electrons in outer shell are stable– Except:

Hydrogen and Helium - stable with 2 electronsBoron - forms bonds surrounded by 6 electronsBeryllium - forms bonds surrounded by 4 e-

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Electron-Dot Strucutres

Notation in which only valence electrons of element are shown by dots

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Electron-Dot Notation

Na Cl

Kr B

Ne1

2

34

5

67

8

H N

S Ba

Lewis Structures

Combine 2 electon-dot structures to show shared electrons– Unshared or Lone pairs

– Indicates kind, number, arrangement, and bonds

– Shared electrons represented by dash– F-F

F F

Lewis Structures

Least electronegative atom is central

– Except hydrogen– Carbon is usually central

Multiple bond represented– Single bond– Double bond– Triple bond

Lewis Structures1. Determine the total number of valence

electrons in the compound.2. Arrange the atoms’ symbols to show how

they are bonded and show valence electrons as dots

3. Compare the number of valence electrons used in the structure to the number available from step 1.

4. Change to a single dash each pair of dots that represents two shared electrons.

5. Be sure that all atoms, with the exception of hydrogen, follow the octet rule.

Problems

Draw the Lewis structure of iodomethane, CH3I.

Draw the Lewis structure of ammonia, NH3

Draw the Lewis structure for hydrogen sulfide, H2S

Draw the Lewis structure for methanal, CH2O, which is also known as formaldehyde.

Problems

Draw the Lewis structure for:– Carbon dioxide, CO2

– Hydrogen cyanide, HCN– IBr– CH3Br– C2HCl– SiCl4– F2O

Polyatomic Ions

A charged group of covalently bonded atoms

Combine with other ions to form ionic compounds

Add/subtract appropriate number of electrons

Place brackets around structureShow the charge of ion outside of

brackets

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Polyatomic Ions

Draw Lewis structures for:– NH4

+

– SO4-

– PO43-

– CO32-

Ionic Bonding

Composed of cations and anions to make neutral compound– NaCl– Cannot be isolated and examined like

moleculesForm crystal lattice to stabilize

– Forces between like-charged ions and opposite-charged ions

– Na+ surrounded by 6 Cl-– Cl- surrounded by 6 Na+

Lattice Energy

Ionic vs. Molecular Compounds

Ionic– Stronger bonds– Higher melting and

boiling points– Hard but brittle– Electrical

conductors when dissolved

– May separate when dissolved

Molecular– Weaker bonds– Lower melting and

boiling points

Metallic Bonding

Chemical bonding the results from attraction between metal atoms and surrounding sea of electrons

Highest energy levels have few electrons

Many vacant orbitals

Metallic Bonding

Metallic Properties– High electrical and thermal conductivity– Absorb many light frequencies– Shiny– Malleable– Ductile– Heat of vaporization → Bond strength

VSEPR Theory

Valence-Shell Electron-Pair RepulsionRepulsion between sets of valence-

level electrons surrounding an atom causes sets to be oriented as far apart as possible– Electrons of bonded atoms want to be as

far away from each other as possible

VSEPR Theory

CO2

– Shared pairs or oriented as far away from each other as possible

– 180 apart– Linear – AB2

BF3– 120 apart– Trigonal-Planar– AB3

VSEPR Theory

CH4

– 109.5– Tetrahedral– AB4

VSEPR & Unshared Electrons

Lone pairs occupy space and influence shape of molecule– H2O → AB2E2 → bent– NH3 → AB3E → trigonal-pyramidal

Unshared electrons repel electrons more strongly than shared electrons

Double and triple bonds treated like single bonds

Problem

Use VSEPR theory to predict the molecular geometry of:– AlCl3– CBr4

– AlBr3

– SF6

– CH2Cl2

Problem

Use VSEPR theory to predict the molecular geometry of:– CO2

– ClO3-

– SF2

– PCl3

Hybridization

Mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies

Example → CH4

– 2s and 2p orbitals hybridize to form 4 identical orbitals called sp3 orbitals

Group 15 and 16– Nitrogen - sp3

Hybrid Orbitals

sp sp2

sp3

Intermolecular Forces

Forces of attraction between molecules

Boiling point is used to measure force of attraction between particles– Weaker than bonds– Boiling points of ionic compounds higher

than covalent moleculesDipole-Dipole ForcesHydrogen BondingLondon Dispersion Forces

Dipole-Dipole Forces

Strongest force is between polar molecules

Dipole– Equal but opposite charges that are

separated by short distance– Direction is from positive pole to negative– Represented by arrow pointing toward

negative pole with a crossed tail

Dipole-Dipole Forces

Forces of attraction between polar molecules are short-range forces

Dipole-Dipole forces cause higher boiling points

Hydrogen Bonding

Many bonds with hydrogen are highly polar because of large electronegativity difference

Hydrogen therefore has positive charge in many compounds

Hydrogen atoms bonded to highly electronegative atom is attracted to an unshared pair of electrons– Example → H2O

Represented by dotted lines

Hydrogen Bonding

Hydrogen Bonding

IceIce Liquid WaterLiquid Water

London Dispersion Force

Electrons are in continuous motionSlight uneven electron distributionTemporary dipoleIntermolecular attractions resulting

from constant motion of electrons and creation of instantaneous dipole

Only intermolecular force acting on noble gases

Strength increases with number of e-

Chapter Review

Pg. 209– 3, 6acfg, 9, 16, 19abcfg, 21bde, 22, 24c,

29ab, 33, 45bc, 46bdf, 47cd, 48acdf, 49a