Chemical BondingMetallic bonding

3 main kinds :Covalent bond results from sharing electrons between the atoms,usually found between non-metals.Ionic bond results from the transfer of electrons from a metal to a non-metal.Metallic bond: attractive force holding pure metals together.Chemical Bonds : the attractive force holding two or more atoms together

Ionic BondsBetween metals with low electronegativity with non-metals with high electronegativityInvolves electron transferHeld by strong electrostatic attraction between the oppositely charged ionshighest tendency will be between elements on the bottom left and those on the top right of the periodic table.

Eg NaClNa: 1s2 2s2 2p6 3s1Na+ (cation) : 1s2 2s2 2p6

Cl: 1s2 2s2 2p6 3s2 3p5Cl- (anion) : 1s2 2s2 2p6 3s2 3p6Electrons gain or lost in order to achieve the octet/noble gas configuration

Using Lewis Structure

CaCl2When calcium reacts with chlorine, both electrons from Ca must be lost, so we need 2 chlorine atoms to gain 1 each.

The oxidation states of various elementsGrp 1,2,3 : +1, +2, +3 respectivelyGrp 5, 6, 7 : -3, -2, -1 respectivelyThe transition element can have various oxidation no: Fe2+, Fe3+, Cu1+, Cu2+ etc

Properties of ionic Compounds solids at room temperature. form a giant crystalline lattice/network structure ie an ordered, continuous structure .very strong and high melting points and boiling pointshard, brittle do not conduct electricity when solid but do when molten or in aqueous solution are more soluble in water than in other solvents.

Lattice structure of NaClSimple cubic packing

Electronegativity

The ability of one atoms in a molecule/in a bond to attract electrons to itself.Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F).Electronegativity increases across a period and down a group.

Electronegativities of ElementsElectronegativity

Bonding ContinuumIt is more accurate to consider bonding as a gradual change from solely ionic through polar covalent to non-polar covalent.The degree of polarity determines whether a bond is classed as ionic or covalent but it is a gradual spectrum.If the difference in electronegativity is > 1.7 then it is usually . 50% ionic.A difference of 0.4 or less is considered non-polar covalent.

Covalent BondingIs the attraction of the nuclei in each atom to the valence electrons they are sharing between them. They formed discrete separate molecules When bonds are formed, the atoms become lower in energy, so more stable

The optimum distance between the two nuclei is the bond length

Electron sharing in covalent bond formation=

Carbon dioxide formed from 2 oxygen atomeach sharing 2 pairs of electrons with a carbon atom=

An oxygen molecule formed by two oxygen atoms sharing two pair of electrons=

Nitrogen molecule formed when two N atoms sharing three pair of electrons=

Pure covalent bondsEqual sharing of electrons Eg. H2(g) forms a single bond (shared pair)

Polar covalent bondUnequal sharing of electrons. One atom will have a higher electronegativity than the other, so it will pull the shared electrons closer to itself making that atom slightly more negative than the other.The Cl (3.00) is more negative than the H (2.20)

Chemical Bonds

Bond Type Single Double Triple# of es 2 4 6Notation = Bond order 1 2 3Bond strengthIncreases from Single to TripleBond lengthDecreases from Single to Triple

Relationship between bond length and bond strength

As bond length decreases, bond strength increases because as the atom gets larger, the forces of attraction gets weaker

Relationship between bond order and bond length

As the bond order increases, bond length increases. This is because as the number of electron pairs increases, the forces of attraction gets greater. This would also result in higher energy

Bond Length Values

Chemical Bonds, Lewis Symbols, and the Octet Rule

The carboxylate group=

Coordinate/Dative covalent bonds Both the shared electrons in the bond come from the same atom. Eg 1 in NH4+, the nitrogen from ammonia donate a pair of electrons when forming bonds with a H+ (a proton) Ammonium has 3 polar covalent bonds and 1 coordinate (dative) covalent bond.

Example 2 Example 3Hydronium (H3O+)Carbon monoxide (CO)

All noble gases except He has an ns2 np6 configuration. Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs).However, there are exceptions to the octet rule.The Octet Rule

Exceptions to the Octet Rule 1.Central Atoms Having Less than an Octet (8) Most typical example is BeH and BF3.Its already stable with 6 electrons surrounding the central atom

2.Central atom having more than an octetObserved only for elements in period three (n=3) and beyondEg PCl5 (cannot be seen in Nitrogen)

Beginning with the n=3 principle quantum number, the d orbitals become available (l=2)

Can use empty d orbitals to accommodate moreelectrons

3. Odd number of electrons There are currently 5 valence electrons around the nitrogen. A double bond would place 7 around the nitrogen, and a triple bond would place 9 around the nitrogen. So can never achieve the octet configuration

How to draw a Lewis Structure1. Count up the total number of valence e- in all your atoms in the formula. Eg. H2O has 1 x 2 for the Hs and 6 for O. Subtract or add electrons for any charge.2. Draw the basic structure of the molecule where (usually) the least electronegative atom is the central atom. Pay attention to the number of bonds elements usually form. Eg. H-O-H3. Add more electron pairs to complete the octets around the central atom followed by the side atoms (only 2 around H).4. If there are not enough electrons to give octet to the central atom, add double or triple bonds.5. Check your final structure has the same number of electrons as you started with.

Lewis Structure of HCN

H has 1 e-, C has 4 e- and N has 5 e- = 10 e-. Carbon forms 4 bonds, nitrogen forms 3 and hydrogen 1.The structural formula accounts for all of the carbon and hydrogen valence electrons but not for nitrogen so it must have a non-bonding pair:

Carbon tetrachlorideCarbon is the central atom. It has 4 bonding pairs.Chlorine wants to share one bonding site each.Need 4 chlorines for every one carbon(Cl has 3 lone pairs and 1 bonding pair)

More examples

Practice drawing and naming Lewis Structures

H2OCH2OO3

Lewis structure of H2O

Lewis structure of O3

Lewis structure of CO3-

VSEPR :Valence Shell Electron Pair Repulsion theory

Accounts for the geometric arrangement of electron pairs around a central atom

Bonding pairs and lone pairs around an atom in a molecule adopt positions where their mutual interactions/repulsion are minimized. Electron pairs are negatively charged and will get as far apart from each other as possible. (Same charge = repulsion)

There are five fundamental geometries for molecular shape:Molecular Shapes

Figure 9.3HyperChem

VSEPR Model

Summary of VSEPR Molecular ShapesSee Ng Web-siteHyperChemCyberChm Gems

e-pairsNotationName of VSEPR shapeExamples2AX2LinearHgCl2 , ZnI2 , CS2 , CO23AX3Trigonal planarBF3 , GaI3AX2ENon-linear (Bent)SO2 , SnCl24AX4TetrahedralCCl4 , CH4 , BF4-AX3E(Trigonal) PyramidalNH3 , OH3-AX2E2Non-Linear (Bent)H2O , SeCl25AX5Trigonal bipyramidalPCl5 , PF5AX4EDistorted tetrahedral (see-sawed)TeCl4 , SF4AX3E2T-ShapedClF3 , BrF3AX2E3LinearI3- , ICl2- 6AX6OctahedralSF6 , PF6-AX5ESquare PyramidalIF5 , BrF5AX4E2Square PlanarICl4- , BrF4-

By experiment, the H-X-H bond angle decreases 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.(LP-LP.LP-BP.BP-BP)Therefore, the bond angle decreases as the number of lone pairs increasesThe Effect of Nonbonding ElectronsHyperChem

Figure 8.10: Drawing Lewis StructuresResonance Structures

Figure 9.12HyperChem

Figure 9.11: Molecular Shape and Molecular PolarityHyperChem

Figure 9.13: Molecular Shape and Molecular PolarityHyperChem

Lewis structures and VSEPR do not explain why a bond forms.How do we account for shape in terms of quantum mechanics? What are the orbitals that are involved in bonding?We use Valence Bond Theory:Bonds form when orbitals on atoms overlap.There are two electrons of opposite spin in the orbital overlap.Covalent Bonding and Orbital OverlapGems - Movie Clip

Figure 9.14: Covalent Bonding and Orbital Overlap

To determine the electron pair geometry:draw the Lewis structure,count the total number of electron pairs around the central atom,arrange the electron pairs in one of the above geometries to minimize e--e- repulsion, and count multiple bonds as one bonding pair.

VSEPR Model (Figure 9.6)

VSEPR Model

Drawing Lewis StructuresFormal ChargeConsider:

For C: There are 4 valence electrons (from periodic table).In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.Formal charge: 4 - 5 = -1.

Drawing Lewis StructuresFormal ChargeConsider:

For N:There are 5 valence electrons.In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.Formal charge = 5 - 5 = 0.We write:CyberChm Gems

Chemical BondingLewisAXE notationVSEPR shapesPolarity

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