CHEMICAL BONDING LEWIS THEORY OF BONDING

Results from the transfer of electrons from a metal to a non- metal. A chemical bond between oppositely charged ions Held together by electrostatic attraction

IONIC BONDING

Formed when an orbital from 2 different atoms overlap Electrons must have opposite spins

COVALENT BONDING

CHEMICAL BONDS Bond Type Single Double Triple # of e’s Notation — =  Bond order Bond strength Increases from Single to Triple Bond lengthDecreases from Single to Triple

TYPES OF BONDING CONDITIONS BETWEEN ELEMENTS Low Electronegativity and low Ionization energy (Metals) High electronegativity and High Ionization energy (Non-metals) Low Electronegativity and low Ionization energy (Metals) Metallic bondingIonic bonding (transferring of electrons between atoms) High electronegativity and High Ionization energy (Non-metals) Ionic bondingCovalent bonding (sharing of electrons between atoms)

Electronegativity

LEWIS BOND THEORY Atom/ions are stable when they have a noble-gas like arrangement of electrons – full octet (valence shell) Electrons are the most stable when paired. Atoms form bonds to achieve a full octet.

LEWIS DIAGRAMS OR STRUCTURES A convention developed to “show” the relationship between atoms when they form bonds. Why is it necessary? predict where the electrons are in a molecule needed to predict the shape of a molecule

LEWIS DIAGRAMS FOR IONIC COMPOUNDS  Identify the number of valence shell electrons and determine the charge on the ion using the “stable octet rule”.  Write the elemental symbol, place dots to represent the electrons in the valence shell, enclose in square brackets and write the ionic charge as a superscript. [Na] + or [ Cl ] -

DRAWING LEWIS STRUCTURES 1.Identify the central atom (lowest electronegativity or atom with the highest bonding capacity) 2.Add up the valence electrons for all the atoms (if there is a charge on the compound, add this in too) 3.Place the other atom around the central atom 4.Draw the Lewis dot diagram for all the atoms – make sure to keep track of the #of electrons. 5.Using arrows, start pairing up electrons to complete the atoms octet 6.Rearrange to show electron sharing, keeping in mind the shape 7.Add bonds for any shared electrons. 8.Count the number of electrons around each atom to make sure you are not disobeying the octet rule.

STRUCTURAL DIAGRAMS FOR COVALENT COMPOUNDS Draw the Lewis Diagram for nitrogen trifluoride (NF 3 ). Step 1. Count the valence electrons N = 5 F = ( 7) = 26 valence electrons

STRUCTURAL FORMULA FOR COVALENT COMPOUNDS Step 2. Write a skeletal structure. Use the least electronegative atom in the centre Electronegativity: N = 3.0 & F = 4.0 FF F N = a pair of e - (a single bond)

Step 3. Complete the octets for each terminal atom (except H) FF F N    

Step 4. Assign any additional electrons as lone pairs on the central atom  FF F N   

Example 2. COCl 2 (24 electrons) CCl O    

Step 5. Make multiple bonds where necessary to complete the octets.  CCl O      CCl O    

Example 3. Chlorate ion, ClO 3 - ((1 x 7) + (3 x 6) + 1) = 26 ClOO O  :   : 

In some covalent compounds, the bonds between atoms occur because one atom has donated both electrons to the covalent bond. This is called a coordinate covalent bond. N : H H H H+H+ + N H H H H + Nitrogen supplies the two lone pair electrons to this N-H bond. The H + ion has no electrons.

To determine the number of coordinate covalent bonds – subtract the bonding capacity (lone valence electrons) from the number of bonds the atom has. N H H H H + Nitrogen Bonds  4 Bonding capacity  3 Coordinate bonds  4-3=1

In some compounds the SCH3U guidelines may not “work”. On occasion, both elements have the same electronegativity or there may be two or more possible Lewis Structures. e.g. CS 2 (both electronegativities = 2.5) is it S=C=S or C=S=S ? Exceptions to the Octet Rule

In such situations, one determines the Formal Charge. The option with the lowest formal charge has the most stable and viable structure. The Formal Charge for an atom is the number of valence electrons in the free neutral atom minus the number of valence electrons assigned to the atom in the Lewis structure.

Formal Charge = (# valence electrons)-(# of bonds)-(# of unshared e - )

CSS Valence electrons466 Electrons assigned646 Formal Charge-220 C=S=S  

SCS Valence electrons646 Electrons assigned646 Formal Charge000 S=C=S  

In some structures the Lewis structure does not represent the true structure of the compound. Bond order is the number of shared pairs of electrons between two atoms. (i.e. – the number of bonds between two atoms) As the bond order increases... The length of the bond decreases. The energy associated with breaking the bond increases.

The number of shared electrons in a bond affects its length and energy. Bond TypeBond Order Bond Length (pm) ( m) Bond Energy (kJ/mol) C-O C=O C-C C=C C C-N C=N C N

CHO 2 - is a polyatomic ion with the following Lewis structure.  The C-O bond lengths are experimentally determined to be between C-O and C=O. The bond order is neither 1 or 2, but considered to be somewhere in between (i.e.- 1.5).  The Lewis structure does not support the experimental data. C H O O

The actual structure is a resonance hybrid of the two resonance structures. C H O O C H O O The resonance structure does not “flip-flop” back and forth between the two. It is a hybrid form of the two.

The bond order for NO 3 - is, (1+1+2)/3=1.33. The resonance structure is... N O O O N O O O N O O O

CC C CC C H H H HH H CC C CC C H H H HH H Benzene has a bond order of 1.5. ( )/6=1.5

The work on quantum theory in conjunction with the success of Lewis structures resulted in the inevitable connections between the two areas of study. Linus Pauling, a friend of Gilbert Lewis, connected the two with the valence bond theory.

PRACTICE COMPLETE THE LEWIS DIAGRAM AND CHEMICAL BONDING WORKSHEETS

Linus Carl Pauling (February 28, 1901 – August 19, 1994)

The half filled orbital of one atom overlaps a half filled orbital of a second atom to form a full orbital with two electrons spinning in opposite directions. The bonding atoms arrange themselves in order to maximize the overlap of the half- filled orbitals. Maximum overlapping of the orbitals creates a bonding orbital with a lower energy and increased stability.

The combination of electrostatic repulsion and opposing magnetic fields (due to the electron’s spin) creates the stability associated with a bonding orbital. Negative charge Electrostatic repulsion Negative charge Electron Spin “North” magnetic field “South” magnetic field Magnetic attraction

During this process, two atoms approach each other and allow their half filled orbitals to overlap and form the stability of a filled bonding orbital.

In some situations a more advantageous bonding scenario can be established by promoting electrons from a full orbital to a similar empty orbital to create two half filled orbitals that are available for bonding. The resulting orbital is a mixture of the two original orbitals and is called a hybrid orbital.

Electrons are promoted from the full “s- orbital” into an empty “p-orbital” which results in hybrid orbitals that have one electron per orbital and characteristics unique to the newly formed orbitals. The new hybrid orbitals are free to become involved in bonds by overlapping with other half filled valence orbitals.

1s 2s2p C (z = 6) 1s 2s2p C (z = 6) 1s sp 3 C (z = 6) sp 3 hybridization 1s 2s2p C (z = 6) 1s 2s2p C (z = 6) 1s sp 2 C (z = 6) sp 2 hybridization p 1s sp C (z = 6) sp hybridization pp

Sigma (  ) bonds – The “end-to-end” overlapping of half filled orbitals to make a full bonding orbital of lower energy level (i.e. – more stable) They occur between “s”, “p” and hybrid orbitals (“sp”,“sp 2 ” & “sp 3 ”) to make single covalent bonds.

Pi (  ) bonds – The “side-to-side” overlapping of half filled “p” orbitals to make more stable filled bonding orbitals.

A combination of “  ” and “  ” bonds makes double and triple bonds. Single bonds Sigma only (  ) Double bonds 1 Sigma (  ) and 1 Pi (  )

A combination of “  ” and “  ” bonds makes double and triple bonds. Triple bonds 1 Sigma (  ) and 2 Pi (  ) Pi (  bond Sigma (  bond

Valence shell electron pair repulsion (VSEPR) theory Only valence shell electrons of the central atom are important in the molecular shape. Valence shell electrons will repel to maximize the distance between the pairs Bond pairs and lone pairs behave similarly

e-pairsNotationName of VSEPR shapeExamples 2AX 2 LinearHgCl 2, ZnI 2, CS 2, CO 2 3AX 3 Trigonal planarBF 3, GaI 3 AX 2 ENon-linear (Bent)SO 2, SnCl 2 4AX 4 TetrahedralCCl 4, CH 4, BF 4 - AX 3 E(Trigonal) PyramidalNH 3, OH 3 - AX 2 E 2 Non-Linear (Bent)H 2 O, SeCl 2 5AX 5 Trigonal bipyramidalPCl 5, PF 5 AX 4 EDistorted tetrahedral (see-sawed) TeCl 4, SF 4 AX 3 E 2 T-ShapedClF 3, BrF 3 AX 2 E 3 LinearI 3 -, ICl 2 - 6AX 6 OctahedralSF 6, PF 6 - AX 5 ESquare PyramidalIF 5, BrF 5 AX 4 E 2 Square PlanarICl 4 -, BrF 4 -