Chemical Bonding and Lewis StructuresChemical

Bonding and Lewis Structures

Chemical Bonding Chemical Bonds are the forces that hold

atoms together. Atoms form bonds in order to attain a minimal

energy state. Bond formation is an exothermic process (just

as bond breaking is endothermic) The type and strength of bond that forms

between reacting particles dictates the physical and chemical characteristics of the molecule or ion in question.

Bond energy is the energy needed to break a bond and is an indication of bond strength.

The Relationship Between Electronegativity and Bond Type

Ionic Bonds

Ionic bonds occur between ions due to electrostatic attraction between positive cations and negative anions.

Form between atoms with large differences in electronegativity (>1.7), usually between a metal and a nonmetal.

Relative strength can be determined using Coulomb’s Law E = k(q1)(q2)

r2

K= coulomb’s constant; q1=charge on one ion; q2= charge on other ion; r=distance between ions

Ionic Bond, continued

The strength of an ionic bond is directly proportional to the magnitude of the charges involved and inversely proportional to the square of the distance between them.

Ionic solids tend to have high melting points

Often soluble in polar solvents, such as water.

Ionic Solid

UNIT CELL Crystal lattice – at the lattice sites in the unit cell

are positive and negative ions held together by Coulombic attraction force between positive and negative ions.

CONDUCTIVITY Solid has no conductivity due to zero empty

valence orbitals. The ions in the crystal are isoelectronic with Group 18.

Molten and aqueous solutions will conduct due to mobile ions.

Covalent Bonds

Nonpolar covalent bonds form between atoms of nonmetals with nearly identical electronegativities while polar covalent bonds for between nonmetals with dissimilar electronegativities (0.4-1.7).

Covalent bonds within molecules are strong, but the binding forces between molecules are relatively weak.

Molecular solids usually have low melting points.

Usually soluble in nonpolar solvents (carbon tetrachloride)

Which of the following bonds would be the least polar yet still be considered polar covalent?

Mg-O C-O O-O Si-O N-O

React 5

Metallic Bonding

Occurs in metallic solids Metal atoms usually have large, positively

charged nuclei and few valence electrons. Nuclei are positioned in a regular geometric

array (lattice) by electrostatic repulsion. Valence electrons are attracted equally by all

nuclei. Leads to the “sea of electrons” model with

nuclei bobbing in a “sea of electrons” Useful for explaining the physical characteristics of

metals (ie, conductivity) Metals have a wide range of melting points.

Metallic Solid

CRYSTAL LATTICEAt lattice sites in unit cell are positive ions

held together by mobile valence electrons traveling through empty valence orbitals.

POSITIVE CHARGE DENSITYSmaller ionic radius = higher melting point

CONDUCTIVITYSolid: excellent due to empty valence

orbitalsLiquid: good due to mobile ions.

Intermolecular Forces

The group of weaker attractive forces between atoms or molecules. NOT BONDS!!!!!

Van der Waals forces London dispersion forces-all atoms and molecules; caused by

temporary dipoles created as electrons move about the nucleus. Strength depends on the number of electrons moving, so molecules

with larger masses (more electrons) have greater London forces. Dipole-dipole force- an attraction between opposite polar ends

of adjacent molecules. Hydrogen bonding- occurs when hydrogen bonds with a very

electronegative anoin (F, O, N) resulting in a very polar molecule. Strongly positive and negative ends of the molecule have stronger

interactions than either london or dipole-dipole forces.

The Effect of an Electric Field on Hydrogen Fluoride Molecules

Lewis Structure

Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule

There are rules for Lewis Structures that are based on the formation of a stable compound

Atoms want to achieve a noble gas configuration

Octet & Duet Rules

Octet Rule – atoms want to have 8 valence electrons

Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

Steps for drawing Lewis Structures

1. Sketch a simple structure with a central atom and all attached atoms

2. Add up all of the valence electrons for each individual atom

If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge

If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

1. Subtract 2 electrons for each bond drawn2. Complete the octet on the central atom &

subtract those electrons3. Complete the octet on the surrounding atoms

& subtract those electrons4. Get your final number

If 0 you are done! If + add that many electrons to the central atom If - need to form multiple bonds to take away

that many electrons

Steps for drawing Lewis Structures

Examples

CCl4 Sketch a simple structure with a central

atom and all attached atoms Cl │Cl – C – Cl │ Cl

Examples

Add up all of the valence electrons for each individual atom

4 + 4(7) = 32 Subtract 2 electrons for each bond drawn 32-8 = 24 Complete the octet on the central atom &

subtract those electrons Done

Examples

Complete the octet on the surrounding atoms & subtract those electrons

24 – 24 = 0 Final number = 0…DONE! Final structure is… __ │Cl │ __ │ __│Cl – C – Cl │ │ │Cl │

Examples

HF

Examples

NH3

Examples

NO+

Exceptions to the octet rule

Sometimes the central atom violates the octet rule and has more or less than 8 valence electrons

Keep using the same rules to draw Lewis Structures

Examples

SF4

ICl3

XeF4

ICl4-

Resonance

When more than one Lewis Structure can be written for a particular molecule

Resonance structure – all possible Lewis structures that could be formed

The actual structure is the average of all of the structures

You MUST show all structures!

Examples

SO3

NO2-

NO3-

Covalent Network Crystal

CRYSTAL LATTICEAt the lattice site in the unit cell there are

atoms held together by strong covalent bonds.

Si, SiO2, C (diamond), and C (graphite)This crystal has the highest melting point.

SUMMARY

Covalent network solids have the highest melting points because of the strongest forces holding the crystal together (covalent bonds).

Metallic solids generally have the next highest melting points, but the larger the ionic radius, the lower the melting point.

Ionic crystals next. Molecular crystals last.

Consider IMF to rank these. The stronger the IMF, the higher the melting point.