CHAPTER 6CHEMICAL BONDING

COVALENT MOLECULES

PART 1:CHEMICAL BONDS

CHEMICAL BONDS Mutual electrical attraction between the

nuclei and valence electrons of different atoms that binds the atoms together

Atoms are usually less stable by themselves than combined with other atoms

TYPES OF BONDING Ionic Bonding

Chemical bonding resulting from the attractions of cations and anions In pure ionic bonding, atoms give up

electrons completely to other atoms, which accept them completely to form ions which attract one another

TYPES OF BONDING Covalent Bonding

Chemical bond that shares the electrons between two atoms

A pure covalent bond results in the electrons being “owned” equally by two atoms Both nuclei have these electrons around

them an equal amount of timeExample of purely covalent bond : F2

Both atoms attract each others’ electrons equally (both have an electronegativity of 4.0, so the pull on the electrons is exactly the same from both atoms)

CLASSIFYING BONDS:IONIC OR COVALENT Most compounds do not have purely

covalent or ionic bondsThe bonds fall somewhere on a continuum

Look at the bond’s percentage of ionic characterBonds with 0 to 50% ionic character are classified as covalentBonds with > 50% ionic character are classifiedas ionic

IONIC VS. POLAR COVALENT VS. NONPOLAR COVALENT BONDS Ionic Bonds

Electrons get transferred between atoms forming ions that attract one another

Polar Covalent BondElectrons shared between atoms unequally

Nonpolar Covalent BondElectrons shared between atoms equally

PART 2:CLASSIFYING BONDS USING

ELECTRONEGATIVITY

CAN USE ELECTRONEGATIVITY DIFFERENCE AS A GUIDE We can predict the bond type between

two atoms by calculating the absolute value of the electronegativity difference between them.

Electronegativity Difference

Bond Type

0- 0.3 Nonpolar Covalent

>0.3 – 1.7 Polar Covalent

>1.7 Ionic

IONIC BONDING If the absolute value of the

electronegativity difference between two atoms is > 1.7, then the bond they form is probably ionic

Example #1: NaClNa – electronegativity of 0.9Cl – electronegativity of 3.0Electronegativity difference = | 0.9 – 3.0 | = 2.1

An electron moves from the Na atom to the Cl atom causing Na to have a +1 charge and Cl to have a -1 charge

The Na + ion is then attracted to the Cl- ion forming a very strong bond

POLAR COVALENT BONDS If the absolute value of the electronegativity

difference between two atoms is >0.3 - 1.7, then the bond they form is probably polar covalent

Covalent because the electrons are shared between the atoms

Polar because the electrons are shared unequally between the atoms (spend more time around one nucleus)

Example: BF3 Boron – electronegativity of 2.5 Fluorine – electronegativity of 4.0 Electronegativity difference of | 4.0-2.5|= 1.5

Polar covalent - electrons spend more time around the F nucleus due to greater electronegativity

Electron cloud is distorted towards the Fluorine atoms (more electron density) resulting in a partial negative charge on the fluorine atoms and a partial positive charge on the boron atom.

NONPOLAR COVALENT BONDS If the absolute value of the

electronegativity difference between two atoms is 0.0- 0.3, then the bond they form is probably nonpolar covalent

Electrons are shared equally between the nuclei of the bonded atoms, resulting in a balanced distribution of electrical charge

Ionic character of 0 to 5% ionic character

Example: F2

Electronegativity difference of |4.0 – 4.0| = 0.0

PART 3:COVALENT

COMPOUNDS

COVALENT COMPOUNDS Molecules

Neutral group of atoms held together by covalent bonds

Can exist on its own as an individual unit Molecular formula

Shows the relative numbers of atoms of each kind in a chemical compound by using element symbols and subscripts to show the number of atoms Examples:

H2O - water

O2 - oxygen

C12H22O11 - sucrose

BOND ENERGY Energy required to break a chemical

bond and form neutral isolated atomsUnits kj/moleEnergy required to break one mole of bonds

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

EXCEPTIONS TO THE OCTET RULE Hydrogen – needs total of 2 electrons Beryllium – needs total of 4 electrons Boron – needs 6 electrons Aluminum – needs 6 electrons Atoms of elements in groups 15, 16, 17,

and 18 can sometimes form expanded octetsCan have more than 8 electrons

surrounding the nucleusUsually not N, O, or F

LEWIS STRUCTURES AKA ELECTRON DOT NOTATION Symbols represent the nucleus and

inner shell electrons Dot pairs between the symbol represent

pairs of shared electrons Dots adjacent to symbol represent

unshared pairs of electrons

STEPS FOR DRAWING LEWIS DOT STRUCTURES1. Count up the number of valence

electrons.  E.g. in CCl4 there is 1 C which has 4 valence

electrons and 4 Cl atoms which each have 7 valence electrons  Therefore for C 1 x 4 = 4 and for Cl 4 x 7 = 28 Adding 4 and 28 together gives you 32 electrons

2. Identify the central atom and put the other atoms around the central atom.

If C is present, put C in the center. Otherwise, the least electronegative atom goes in

the center. H never goes in the center. Halogens don’t usually go in the center unless they

are bonded to other halogens

STEPS FOR DRAWING LEWIS DOT STRUCTURES (CONT’D)3. Put a pair of electrons (dots) between

the central atom and the outside atoms.

4. Start distributing the remaining electrons

Subtract those electrons placed in step 3 from the total in step 1 to find electrons remaining

Put the remaining electrons around the outside atoms first until all atoms either have an octet or are satisfied (Exceptions to the Octet Rule for the exceptions).

5. If there are electrons left over after satisfying the octet rule, put them as unshared pairs on the central atom.

STEPS FOR DRAWING LEWIS DOT STRUCTURES (CONT’D)6.  If there are too few electrons

for all atoms (including the central atom) to have an octet (or be satisfied if it is an exception), then double and or triple bonds should be considered.

C, N, and O are most atoms likely to form double or triple bonds. 

DRAWING LEWIS DOT STRUCTURES FOR POLYATOMIC IONS To draw the Lewis dot structure for a

polyatomic ion (a covalently bonded molecule that has a charge), you follow the same steps except the following:When counting up the valence electrons in

step 1, add the absolute value of the charge if the charge is negative to the total of the valence electrons. If the charge is positive, subtract the absolute value of the charge from the total valence electrons.

After completing the Lewis dot structure, put the entire structure in brackets with the charge on the right upper corner on the outside of the right bracket.

MULTIPLE COVALENT BONDS Some elements, especially C, N, and O

can share more than one electron pair.Double bond – 2 pairs of electrons are

sharedTriple bond – 3 pairs of electrons are shared

MULTIPLE BONDS ENERGIES AND BOND LENGTHS Double bonds have a greater bond

energy and shorter bond lengths than single bonds

Triple bonds have a greater bond energy and shorter bond lengths than single bonds

Triple bonds stronger than double Double bonds stronger than single

Bond Avg. Bond Length (pm)

Avg bond energy (kj/mole)

C-C 154 346

C=C 134 612

C= C 120 835

C-N 147 305

C=N 132 615

C=N 116 887

RESONANCE STRUCTURES Bonding in molecules or ions that

cannot be correctly represented by a single Lewis structureDraw possible Lewis structures with double-

headed arrow in between

Bond between oxygens is a hybrid of a single and double bond

PART 4: CHEMICAL FORMULAS, NAMING, AND

SHAPES

CHEMICAL FORMULA Chemical formula indicates the relative

number of atoms of each kind in a chemical compound. Molecular formula indicates the number of

atoms of each element contained in a single molecule

E.g. C8H18

Ionic compounds have formula units, not molecules (Remember, these ions are located in a lattice of positively and negatively charged ions – are not found as individual units)

One formula unit is used to designate the simplest ratio of the cations to anions

E.g. in a crystal of NaCl, the formula NaCl is used to represent one unit

NAMING BINARY MOLECULAR COMPOUNDS (CHAPTER 7 P. 227-229)

Prefix system of naming – must learn prefixesNumber Prefix

1 Mono-

2 Di-

3 Tri-

4 Tetra-

5 Penta-

6 Hexa-

7 Hepta-

8 Octa-

9 Nona-

10 Deca-

NAMING BINARY MOLECULAR COMPOUNDS Element with the smaller group number is given first

in the name If both elements are in same group, element whose

period number is greater is given first. If first element subscript is one, no prefix is given. If subscript is greater than 1, then the prefix

corresponding to the subscript is added to the beginning of the element name.

The second element is always given a prefix indicating the number of atoms and the element root + ide is used.

The –o or –a at the end of a prefix is usually dropped when the word following the prefix begins with another vowel i.e. monoxide instead of monooxide or pentoxide instead of pentaoxide

BINARY MOLECULE NAME EXAMPLESFormula Prefix-system name

N2O Dinitrogen monoxide

NO Nitrogen monoxide

NO2 Nitrogen dioxide

N2O3 Dinitrogen trioxide

N2O4 Dinitrogen tetroxide

N2O5 Dinitrogen pentoxide

VSEPR THEORY Theory used to determine the shape of

moleculesBased on the premise that repulsion

between the sets of valence electrons surrounding an atom causes the electron pairs to be oriented as far apart as possible.

Looks at the bonded pairs and lone pairs on the central atom only Bonded pairs are set as far apart from one

another as possible

MOLECULAR GEOMETRIES Some

representative molecular geometries – need to learn molecular geometries located on VSEPR Summary sheet in Chapter 6 folder

HOW TO READ VSEPR SUMMARY SHEET Must determine molecule type from Lewis Dot

structure – must draw this first Let A = central atom (in example below, N) Let B = outside atoms regardless of whether they are the

same element or not (3 outside atoms ) Let E = represent the # of lone pairs on the central atom

(only) (one lone pair on N) Let subscripts for B and E represent the number of outside

atoms and number of lone pairs on the central atom Example: NF3 is AB3E N is central atom or A, the 3 Fs are represented as B3 and there is one lone pair on the

central atom or E Molecular geometry - from VSEPR Summary sheet is trigonal pyramidal

BOND ANGLES Angles between the bonded pairs of

atoms around the central atom Bond angles between bonded pairs tend

to be larger for those atoms without lone pairs on the central atomLone pairs cause increased repulsion and

force bonded pairs closer together

POLAR MOLECULES Electrons are unevenly distributed

throughout the moleculeBecause of this there are areas of partial

positive charge and areas of partial negative charge

Forms dipoles Equal but oppositecharges separatedby short distances

More positive more negative

Partial +

Partial -

FACTORS DETERMINING POLARITY OF MOLECULES Polar Bonds present in molecule Symmetry of molecule

Do dipoles when taken together, cancel out or form a dipole in one direction? If cancel out - then nonpolar If overall dipole is present – polar

GUIDELINES FOR DETERMINING POLARITY If molecule type is AB2,AB3,AB4, AB5, AB6,

AB2E3, or AB4E2, and the outside atoms are the same element, then the molecule is nonpolar

If molecule type is AB and A and B are the same element, then the molecule is nonpolar

All others are polarE.g. CH4 is AB4 all of the B’s are the same

(H), therefore the molecule is nonpolarCHCl3 is AB4 but not all of the B’s are H, (1 H

and 3 Cls), therefore the molecule is polar

HYBRIDIZATION When bonding occurs, orbitals from the

bonding atoms become mixed and form new orbitals that are equivalent energies (i.e. hybrid orbitals)Hybrid orbitals

Orbitals of equal energy produced by the combination of two or more orbitals on the same atom# of orbitals produced = # of orbitals that

have combined

HYBRID ORBITALS AND MULTIPLE BONDS All single bonds are sigma bonds – σ

Atomic orbitals overlap and form hybrid orbitals

Double and Triple bonds have one sigma bond In addition, double bonds also have one π

bond Pi bonds (π) are overlapping atomic p orbitals on

adjacent atomsTriple bonds also have two π bonds

HYBRID ORBITALS# of bonding pairs and lone pairs around central atom

Hybrid Orbital Formed

# of hybrid orbitals

Geometry of electron Pairs (not geometry of atoms)

2 sp 2 Linear

3 sp2 3 Trigonal planar

4 sp3 4 Tetrahedral

5 dsp3 or sp3d

5 Trigonal bipyramidal

6 d2sp3 or sp3d2

6 octahedral

Count double and triple bonds as one bonding pair since there is only one sigma bond

PART 5: INTERMOLECULAR FORCES

INTERMOLECULAR FORCES VS. INTRAMOLECULAR FORCES

Weaker Stronger

Intermolecular forces Between molecules 3 types

Dipole-dipole interaction

Hydrogen Bonding London Dispersion

Forces

Intramolecular Forces Within molecules True bonding

Covalent bonding

3 TYPES OF INTERMOLECULAR FORCES1. Dipole-dipole interactions

Attraction between molecules that occurs between oppositely charged ends of partially charged molecules

Occurs between polar molecules

TYPES OF INTERMOLECULAR FORCES2. Hydrogen bonding – special case of

dipole-dipole interactionNOT Bonding!!!!! – intermolecular forceOccurs when highly polarized H is attracted

to the lone pair of a highly electronegative N, O, and F

Occurs with molecules having N, O, or F directly bonded to a H

Strongest intermolecularforce

TYPES OF INTERMOLECULAR FORCE3. London Dispersion Forces

Weakest intermolecular forceOccurs due to random motion of electrons

in molecules forming temporary dipoles (areas of partial charge) which attract each other (oppositely charged ends attract)

Due to temporary nature of the dipoles, attractions are fairly weak but when added together can be significant

Occurs in all molecules (polar and nonpolar)

STRENGTH OF INTERMOLECULAR FORCES Hydrogen Bonding > Dipole-Dipole

Interactions> London Dispersion Forces Strength of intermolecular forces

determines whether molecules haveHigh surface tensionHigh boiling pointCapillary action

DETERMINING THE INTERMOLECULAR FORCES ASSOCIATED WITH A MOLECULE

Is the molecule polar or nonpolar

If polar, is thereAn N,O, or F directlyBonded to an H in the structure?

If yes, then the molecule has hydrogen bonding and London Dispersion Forces

If nonpolar, London Dispersion Forces only

If no, then the molecule has a dipole-dipole interaction and London DispersionForces

PROPERTIES OF SUBSTANCES ASSOCIATED WITH INTERMOLECULAR FORCES Surface Tension

Force that pulls adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size

Molecules at the surface do not have other like molecules on all sides of them and consequently they cohere more strongly to those directly associated with them on the surface forming a surface "film” or “skin”.

SURFACE TENSION IN MOLECULES The greater the strength of the intermolecular

forces between molecules in a substance the greater the surface tension.

Hexane and Water Which will have the greatest surface tension?

Hexane is nonpolar and only has London dispersion forces whereas water is very polar and has hydrogen bonding as well as London Dispersion Forces Water has higher surface tension

SURFACE TENSION AND TEMPERATURE The greater the temperature of the

substance, the lower the surface tension.The greater the temperature, the greater

the average kinetic energy of the molecules, the further apart they are from one another and the lower the force of attraction between them.

BOILING POINTS AND INTERMOLECULAR FORCES Boiling

Change of a liquid to bubbles of vapor that appear throughout the liquid (not just from surface like evaporation)

Boiling points are determined by the strength of the intermolecular forces between moleculesThe greater the strength of the

intermolecular forces between molecules, the higher the boiling point.

BOILING POINT TRENDS In Groups 15, 16, and

17 there is a more elevated boiling point than expected for compounds with N, O, and F directly bonded to a H due to the very strong hydrogen bonding that occurs. Much larger energy than

normal is required to overcome the attractive force between molecules to allow the liquid to escape into the gas phase.

Higher than expected B.P.

BOILING POINTS ARE DETERMINED BY STRENGTH OF INTERMOLECULAR FORCES Example 1:

Hexane or Water

Which substance has the higher Boiling PointHexane has only London Dispersion Forces

whereas water has hydrogen bonding and London Dispersion forces

Due to strength of hydrogen bonding, water will have higher boiling point

BOILING POINTS ARE DETERMINED BY STRENGTH OF INTERMOLECULAR FORCES

Example 2: Isopropyl alcohol vs. waterBoth polar with hydrogen bonding and

London dispersion forcesWhich has the higher boiling point?

Water has stronger intermolecular forces due to compact shape and very polar bonds

Isopropyl alcohol is still polar but due to carbons and hydrogen areas of the molecule that are more nonpolar in nature

Due to stronger intermolecular forces, water has the higher boiling point

Nonpolar areas

Polararea

Highly polar molecule

Example 3: C8H18 (octane) vs. C4H10 (butane)Which will have a higher boiling point?

Both have only London Dispersion Forces but octane has stronger intermolecular forcesMore randomly moving electrons that

can form temporary dipoles in more places in the molecule

Therefore octane has a higher boiling point than butane

Octane Butane

BOILING POINTS ARE DETERMINED BY STRENGTH OF INTERMOLECULAR FORCES

CAPILLARY ACTION Attraction of the surface of a liquid to the

surface of a solid Responsible for meniscus seen with water in a

graduated cylinder

Water molecules are attracted to the side of the graduated cylinder and these water molecules attract more water molecules so that water “travels” up the side

If tube is small enough in diameter (capillary tube), then water will rise against gravity due to attractive forces of water molecules

CAPILLARY ACTION IS DETERMINED BY STRENGTH OF INTERMOLECULAR FORCES

The stronger the intermolecular forces the more capillary action will be seen and the larger the meniscus.Hexane – nonpolar – only weak London

Dispersion Forces will not show a meniscus in a graduated cylinder

Water - very polar – strong hydrogen bonds – there will be a pronounced meniscus in a graduated cylinder and will readily flow up a capillary tube against gravity