chemical bonding

The Microscopic Structure of Matter Types of Chemical Forces Formation of Ionic Compounds Properties of Ionic and Covalent Compounds Formation of Covalent Compounds Lewis Structures - introduction Molecular Geometry - introduction Molecular Polarity – introduction Making Lewis structures & determining their

characteristics

CHEMICAL (ELECTROMAGNETIC)

FORCES

Electrostatic Attractions Electron Orbital Bonding

Attractions between oppositely charged chemical species or

oppositely charged regions of chemical species

Attractions due to overlap of electron

orbitals

Integer charges in ionic lattices = IONIC

BONDING

Partial charges

Hydrogen Bonding

Occupied with occupied orbitals

COVALENT BONDING

“Pooling” of orbitals METALLIC &

SEMIMETALLIC BONDING

Occupied with unoccupied orbitals

CO-ORDINATE COVALENT BONDING

Unbounded chem. species NETWORK COVALENT

Discrete chem. species MOLECULAR COVALENT

Homogeneous

Dipole – dipole (non H-bond)

London (induced dipole)

Ion – ion Ion – dipole

Mixed dipole Mixed H-bond

Mixed London

Heterogeneous Between species from same parent species Between species from

different parent species

“Bonding”: General Ideas and TypesGeneral Principle: Atoms interact with atoms in such a way as to

achieve the best SET of electronic configurations for all the atoms involved

Chemical “Bond” - The attractive force that holds the atoms of a substance together

Ionic “bonding” = Electrostatic attraction = Exchanging electrons as a strategy to get more stable configurations

Covalent (molecular) bonding = Sharing orbitals (with or without electrons) as a strategy to get more stable configurations

Metallic bonding: Pooling electrons in orbitals as a strategy to get more stable configurationscommon metals: Fe, Cu, Sn, Pb (elements);bronze (an alloy of two metals)Will not cover

Principles of Bond FormationOnly the valence electrons of atoms are available for

chemical bondingValence electrons = electrons in outermost (valence) shell of

an atomCertain arrangements of electrons are more stable than

othersHe, Ne, Ar, Kr, Xe do NOT form compounds

Chemical bonds are formed because the energy is lowered upon bond formationThe number of favorable factors outweighs the number of

unfavorable ones!

Rationale for Ionic “Bonding” Atoms of elements exchange electrons such as to produce

ions: 1. that have electronic configurations of the nearest noble gas

if at all possible or that have other stable electronic configurations

1. “duet” for He, “octet” for other noble gases

2. that produce salts with strong ion-ion attractions (lattice forces)

There is a limit to the number of electrons that can be donated or accepted (+3 or - 3)

Electron Dot Structures

Electron dot structure help to determine the number of electrons that can be lost or gained to reach a stable octet (duet)

CB N O FLi Be

• What would Li, Be do?• Would B lose 3 of its 5 electrons; would Al be more

or less reluctant to lose 3 electrons?• What would F do readily, O less readily and N even

less readily?• What is C’s dilemma?

Na

+ +

Na+

Cl Cl-

Cl- Na+ClNa

Formation of IonsRepeat but with electronic configurations

e-

(1s22s22p63s23p5)

(1s22s22p6)(1s22s22p63s1)

(1s22s22p63s23p6)

Ca

+ +

Formation of Ions – another exampleFormation of Ions – another exampleFormation of Ions – another exampleFormation of Ions – another example

Ca2+

O O2-

O2- Ca2+ OCa

2 e-(1s22s22p63s23p64s2)

(1s22s22p6)(1s22s22p4)

(1s22s22p63s23p6)

Structures of Ionic CompoundsCrystalline state

Network of cations and anions with overall zero chargeArrangement of cations and anions depends on size and

charge of ionDifferent types of macroscopic crystals as a consequence

Size of network can vary (small and large crystals)Dispersed state

Requires “dispersing agent”Solvent (usually water); heat

SolutionsRequires that water can pull apart crystal LATER

Molten saltsHave to be above melting point of salt

Crystals of Cesium Halides

CsF

CsCl

CsI

CsBr

Crystals of copper salts

Formulas of Binary Ionic CompoundsBinary Ionic Compounds

Two different ions, one metal, one nonmetal

Metal ion + (cation) Nonmetal ion - (anion)Total + = Total –

Formulas of Binary IonicsMetal ion first, nonmetal followsSubscripts indicate number of each

ion used to get balance of chargeTypes of Binary Ionic

CompoundsFixed charged metals

Metals from columns 1, 2, 3 & 13Variable charged metals

All other metalsCharges on monatomic nonmetals

ions are CONSTANT

*Method for Making Compounds

O2-

e-K K+2

O 2 e-

22= K2O

Properties of Ionic CompoundsIonic compounds have very high melting points. They are always

solids at room temperature.Attractions between cations and anions are very strongThe larger the charges, the greater the attraction, the higher the melting point

(usually)Usually hard and brittle

Ionic compounds dissociate into cations and anions when melted.They are electrical conductors when melted.

Ionic compounds dissociate into cations and anions IF they can dissolve in water.Aqueous solutions with dissolved ions (electrolytes) are electrical conductors If an ionic compound cannot dissolve in water,, the crystal lattice forces must

be strong Ionic compounds with larger charges tend not to dissolve in water LATER

Rationale for Covalent Bonding

NOTE: treatment covalent bonding, geometry, polarity (4.5, 4.7, 4.8) in your book is lame please go with this treatment instead

Atoms of electron rich elements share orbitals (with or without electrons) such as to produce: electronic configurations of the nearest noble gas (in most cases)molecular compounds with strong bonds (in all cases)

An electron rich element is one that has half or more of its outer valence shell filledall nonmetals & hydrogen

“Sharing” of orbitals can be:Equal each atom contributes the same number of electronsUnequal atoms contribute different numbers of electrons“Parasitic” (co-ordinate covalent) one atom contributes all the electrons in

a bond

VSEPRnot in your textbook

Atomic vs molecular orbitalsAtoms of each element are born with atomic orbitals

Just fine if you are born perfect and stay single !!!Orbitals need to adjust when orbitals of one atom interact

with orbitals of another atomHybridization (making friends, marriage) Formation of molecular orbitals

Still only 2 electrons per orbital

Valence Shell Electron Pair RepulsionOrbitals used in bonding as well as “leftover” orbitals stay as

far away from each other as possible in order to avoid electron repulsions

Need to know geometry and angles in order to know how to spread out orbitals

The Octet and Duet RulesFor H the nearest noble gas is helium

H in covalent compounds will have a single bond sharing 2 electrons = DUET RULE

For other nonmetals, the nearest noble gas has a completed outer valence shell of 8 electronsthese elements will form enough bonds such that

afterwards the total number of electrons is 8 = OCTET RULE

total number of electrons can be any combination of shared and unshared electrons

Sharing does not have to be equal or symmetricBreaking the octet rule occurs we won’t do

Types of Covalent Bonds

• Single Bonds

• A:B (A-B) - one pair of electrons shared by 2 atoms

• Overlap of two orbitals (one from each atom)

• Double Bonds

• A::B (A=B) - two pairs of electrons shared by 2 atoms

• Overlap of 4 orbitals (two from each atom)

• Triple Bonds

• A:::B ( A B) - three pairs of electrons shared by 2 atoms

• Overlap of 6 orbitals (three from each atom)

– IMPORTANT: NOTICE THAT WHERE THE ELECTRONS COME FROM IS NOT SPECIFIED

Lewis Structures

• Lewis Structure - A picture indicating the way in which the valence electrons are distributed in a molecule

• Electrons in bonds are represented by lines

• no dots are placed between atoms!

• Nonbonding electrons are represented by pairs of dots

• Each atom is given its preferred number of bonds, if possible

• Preferred number of bonds = Number of unpaired dots in the dot structure of the free atom = valence

• If the preferred number of bonds is less or greater than the preferred number, this difference has to be made up by neighboring atoms

Table 4.11

Preferred Number of Bonds

Arrangements of Groups vs Molecular Geometry

• ARRANGEMENTS = Name for the number of groups around a central atom• Groups are:

• connected atoms (by bonds due to overlapping orbitals)

• unshared pair of electrons

• Groups are arranged about a central atom in such a way as to keep the groups as far apart as possible

• MOLECULAR GEOMETRIES = Name that describes how atoms are arrayed around a central atom• Molecular geometries are subclasses of arrangements

LINEAR

TETRAHEDRAL

PLANAR

ARRANGEMENT

Two double or one triple bond

One double bond

All single bonds

Example 1: Correlation between Lewis structure, arrangement and molecular geometry

Example 2: Correlation between Lewis structure, arrangement and molecular geometry

Example 3: Correlation between Lewis structure, arrangement and molecular geometry

ElectronegativityThe most important idea in chemistry

• Electronegativity - A measure of the ability of a nucleus in an atom to attract electrons in a bond toward itself.

• The most electronegative elements would be those that have large effective nuclear charge

• Elements can be given numerical values indicating relative ability

• Electrons in bonds will be “polarized” to the more electronegative atom in the bond

Figure 4.5

Bond Polarity

• Nonpolar Bond - A covalent bond in which the electrons are shared equally between the bonded atoms.• Bonded atoms have the same electronegativity

• Polar Bond - A covalent bond in which there is unequal sharing of electrons between the bonded atoms.– Bonded atoms have unequal electronegativities

– The atom having the higher electronegativity will have a slight negative charge (-)

– The atom having the lower electronegativity will have a slight positive charge. (+)

Table 4.14

Table 4.15

Molecular Polarity

• Nonpolar Molecule - A molecule that does not have a net positive end and a net negative end.

• All bonds are either nonpolar or polarities in bonds cancel out because of symmetry

• Polar Molecule - A molecule that has a slight positive charge on one end and a slight negative charge on the other end.

• Bonds are polar and do not cancel each other out = molecules are asymmetric

A “Formula” for Success for Structure, Geometry and Polarity

Formula -------->Lewis structure Tool = follow the rules for making a Lewis structure

Lewis Structure -------> Arrangement Tool = Count the number of groups around the central atom and assign

the arrangement name groups = atoms (however connected, single, double or triple) and electron

pairs

Arrangement -------> Geometry Tool = Count the number of atoms attached to the central atom (NOT the

number of bonds) staying in the same arrangement family, assign the geometry subclass name

Arrangement -------> Bond Angles Tool = assign all bond angles that connect atoms to each other

do NOT leave the arrangement family

Arrangement -----> Polarity (applies only to neutral species) Tool = inspect the molecule in 3D and determine whether one end of the

molecule is different than another if there is a difference, the molecule is polar, otherwise it is nonpolar

Systematic approach to making Lewis structures & determining their characteristics

Making the Lewis structureAll species will obey octet rule (except for H)No non octet species (such as in your text)

Determining the correct arrangementBond angles always the same for a given arrangement

Determining the geometry SUBCLASSDetermining whether the species has overall polarity

Lewis Structures – Facts to start with

Determine the number of “normal" covalent bonds by knowing the valence of elementF, Cl, Br, I = 1 O, S, Se, Te = 2N, P, As = 3 C, Si = 4

Some species are charged: this has to be accounted forReasons for charge Later

When you make a structure, do NOT worry about “breaking” the rules for bonding for certain elements, either too few or too manySome atoms will have more than the usual number, others less,

BUT these compensate each otherHave faith in the rules you will get the right answer

Determine the valence for each of the elements present in the

formula and multiply by the number of each of these element present

Divide this number by 2 since the valence is how many bonds you need for each atom and bonds are shared between atoms

Place the least electronegative atom (but never H) in the center and place the other atoms around this central atom

Put bonds in such a way as to give each atom its preferred number of bonds (valence), if possible.You may have to double and triple up bondsIf an atom has too many bonds, a neighboring atom will have

too few: this is because of unequal sharing but overall the normal bonding number will be correct

Add lone pairs to complete the octet as needed around each atom (but not H!)

Check your structure

Lewis Structures – Neutral Species

Lewis Structures – Charged species Determine the valence for each of the elements present in the

formula and multiply by the number of each element present Subtract one for each overall negative charge and add one for

each overall positive charge. This takes care of the shortage or excess of bonds due to charge

Divide number you get by 2 since the valence is how many bonds you need for each atom & bonds are shared between atoms

Place the least electronegative atom (but never H) in the center and place the other atoms around this central atom

Put bonds in such a way as to give each atom its preferred number of bonds (valence), if possible.

You may have to double and triple up bonds There will be a “missing” bond from the overall normal

number for every negative charge and an extra bond from the normal number for every positive charge

Add lone pairs to complete the octet as needed around each atom (but not H!)

Check your structure

Making Lewis Structures

F2, O2, N2

HCl, H2O, NH3, SBr2, PF3, SiCl4

CO2, SO2, HCN

NH4+,NO3

-, NO2-, CO3

2-, SO42-

ClO4-, BrO2

-

H2SO4, HCO3-, CH2O, SCN-, OCS

*

Determining Arrangements & Bond Angles around central atom(s) – Step by Step

1. Count the number of groups GROUPS = ELECTRONS PAIRS AND BONDING

GROUPS ONE ELECTRON PAIR = ONE GROUP ONE BONDING GROUP = ONE GROUP

A bonding group = a single or a double or a triple bond

2. Arrangement & Bond Angle Chart 2 groups = linear = 180 3 groups = (trigonal) planar = 120 4 groups = tetrahedral = 109.5

*

IN DETERMINING GEOMETRY DO NOT LEAVE ARRANGEMENT

1. Linear all geometries linear2. Planar

3 atoms planar2 atoms and an electron pair bent

3. Tetrahedral4 atoms tetrahedral3 atoms and an electron pair

pyramidal2 atoms and two electron pairs bent

Determining Geometries around central atom(s) – Step by Step

*

Rules for Overall Polarity of Neutral Species

Arrangement of Groups

Linear

Trigonal Planar

Tetrahedral

Polarityif atoms on either side of the

central atom are different, the species is polar

if any group around central atom is different, the species is polarbe careful with resonance

if any group is different, the species is polar

ALWAYS USE THE ARRANGEMENT TO DETERMINE POLARITY *

LINEAR

TETRAHEDRAL

PLANAR

ARRANGEMENT

Two double or one triple bond

One double bond

All single bonds

180

120

109.5

BOND ANGLE SUBCLASS

THE ULTIMATE USEFUL SUMMARY SLIDE

One of the S-O bonds is double