Chapter 8 : Covalent Bonding

Section 8.1: Molecular Compounds

What is a molecule? A molecular compound? ● A molecule is a neutral group of atoms joined together by covalent bonds● A molecular compound is a compound that composed of molecules ● A covalent bond is a bond formed by the sharing of electrons between atoms ● This is like a ‘tug of war’ between atoms for electrons, joining the atoms

together

What are common examples of molecules? ● Oxygen gas consists of oxygen

molecules● Each oxygen molecule is made of

two covalently bonded oxygen molecules

● This is referred to as a diatomic molecule

○ A molecule that contains two atoms

● Other common examples are Hydrogen, nitrogen and the halogens

How are molecules represented? ● Molecular formulas are used to show how many atoms of each element a molecule of a substance

contains ○ Example - H2O

● Subscript refers to the number of atoms of that element in the molecule. In water there are two hydrogens bonded to one oxygen atom

● These formulas are not always the lowest whole number ration - why? ● C4H5 represents Butane. C2H5 would be a different molecule ● Molecular formulas do not tell you about a molecule's structure. For this you need a structural formula

What else does the formula not tell you? ● The chemical formula

does not tell you the overall shape of the molecule

● Carbon dioxide - linear● Water - bent

So what are the key differences between ionic and molecular compounds? ● The formulas of molecular compounds describe a molecule● The formulas of ionic compounds describe a formula unit

Characteristics of molecular and ionic compounds continued

The Nature of Covalent Bonding

Section 8.2

How do Covalent bonds relate to the Octet rule? ● Like Ionic bonds, noble gas

electron configurations are key● Atoms will share electrons so each

atom has a noble gas electron configuration

○ This is more common for nonmetals in the periodic table

● Atoms are able to share electrons until they complete their octet

○ The octet rule still applies!

What is meant be a single covalent bond? ● A single covalent bond forms when two atoms share a pair of electrons

○ The attraction between the shared electrons and the positive nuclei of the atoms holds the atoms in the molecule together

● On a structural formula this single bond is represented as a dashed line ● Halogens are an example of atoms that form diatomic molecules with a single

covalent bond. ○ When this occurs, each halogen will have three unshared pairs; a pair of valence electrons

that is not shared between atoms

What about more complex molecules?

Water: Top Right, two unshared pairs

Ammonia : Top Left, one unshared pair

Methane: Right: No unshared pairs

What is a double covalent bond? A triple bond? ● Sometimes atoms must share two or three pairs of electrons to reach a noble

gas electron configuration● Double bonds involves two shared pairs of electrons - example carbon

dioxide (CO2) ● A triple covalent bond is the sharing of three pairs of electrons - example

Nitrogen (N2)

What is a coordinate covalent bond? ● A coordinate covalent bond is an atom in which one atom provides both

bonding electrons● Example - Carbon monoxide. How does bonding in carbon monoxide work? ● Once a coordinate covalent bond forms, it is like any other covalent bond. It is

often drawn as an arrow in a structural formula

What are some of the exceptions to the octet rule?

● Try and draw an electron dot diagram for Nitrogen Dioxide (NO2) ● You can’t draw one! The same applies for all molecules with odd number of

valence electrons● Some molecules that do have an even number of valence electrons still fail to

follow the octet rule - example BF3● Some molecules expand the octet to more than 8 - example PCl5

How does energy relate to bonds? ● When atoms combine to form a bond energy is released . The product is

more stable than the reactants ○ Breaking bonds would require energy

● The energy needed to break a bond is the bond dissociation energy, expressed in kj/mol

● The larger the bond dissociation energy, the stronger the bond, and the shorter the bond length

○ Double and tripple bonds have higher bond dissociation energies than their single bond counterparts

● Bond dissociation energies also contribute to reactivity of molecules - high energy means unreactive molecule.

Bond dissociation energy examples

What does an electron dot structure for Ozone (O3) look like?● Ozone has two different types of bonds - one single coordinate covalent bond

and one double covalent bond. ● This would mean that the bonds in Ozone would have different lengths,

however measurements suggest that the two bonds are the same length, being the average of the two structures.

● When you can draw two or more valid electron dot structures for a molecule or polyatomic ion it is referred to as resonance structures

Section 8.3Bonding Theories

What is a molecular orbitals ● A molecular orbital is an orbital that applies to the entire molecule● Individual atomic orbitals can combine to produce a different shaped molecular

orbital ● There are a number of similarities between atomic and molecular orbitals

○ Both contain two electrons when filled○ Those that can be filled by two electrons of a covalent bond are referred to as bonding orbitals

● Covalent bonds are caused by an imbalance between attraction and repulsion of the nuclei and electrons

● Bonds can form because the attraction between atoms is stronger than repulsions

What is VSEPR Theory? ● VSEPR theory states that the repulsion between electron pair causes

molecular shapes that allow for the greatest distance between valence electron pairs

○ Stands for Valence Shell Electron Pair Repulsion Theory

● It is used to estimate the 3D shape of molecules ● Essentially the valence electron pairs want to stay as far apart as possible,

which can give rise to a number of different molecular shapes

What is going on in Linear Molecules? ● Linear is the simplest shape● The valence electron pairs are furthest apart when arranged in a straight line

○ Example - Carbon Dioxide

● In molecules such as methane it is not so straightforward

What is happening in these diagrams?

The relationship between molecular shapes and unshared pairs ● Unshared pairs of electrons are important when predicting molecular shapes ● Unshared pairs of electrons will be held closer to the nucleus than bonded

pairs - why?● Because, there is no atom competing for them● Unshared pairs will always repel bonded pairs - dictating the overall shape of

a molecule● This explains why the angle between bonds is water is less than in ammonia

Can VSEPR theory be used to predict types of bonds in a molecule?

● No, instead this can be achieved through orbital hybridization● Hybridization is the mixing of several atomic orbitals to form the same total

number of equivalent hybrid orbitals ● Orbital hybridization provides information about both molecular bonding and

molecular shape ● In this example, orbital hybridization is the only way that carbon can form 4

bonds

Section 8.4Polar bonds and molecules

What is the difference between a polar and a nonpolar covalent bond? ● Nonpolar: Covalent bond in which two atoms share the electrons equally● Polar: A covalent bond in which two atoms share the electrons unequally

○ The more electronegative atom will attract the electrons more strongly, giving slight changes to the atoms in the bond

When do polar covalent bonds form? ● Whenever there is a difference in electronegativity values for the two atoms

involved ● Electronegativity differences are used to assess the strength of the polar bond

Relationship between polar bonds and polar molecules● The presence of a polar bond in a molecule often makes the entire molecule

polar○ This means that one end of the molecule is slightly positive, and the other is slightly negative

● When the molecule has two poles of opposite charge it is called a dipolar molecule, or dipole.

● Sometimes molecules can have polar bonds but not be polar themselves - how?

● Example - carbon dioxide and water. Which is Polar?

What effects can bond polarity cause? ● Bond polarity can cause intermolecular attractions; weak attraction between

molecules● These can be important, and can determine the state of matter of a molecular

compound● Van der Waals forces: the two weakest intermolecular attractions

○ Dipole interactions: Polar molecules are attracted to one another (oppositely charged parts). Weaker version of ionic bonds

○ Dispersion forces: The weakest of all intermolecular interactions, caused by motion of electrons.

Dispersion forces explained…...

● Dispersion forces are caused by the motion of electrons● All molecules experience dispersion forces - not just

polar molecules● They occur when moving electrons of one molecule

happen to move to one side of the molecule ● Their force then repels electrons in a nearby molecules,

causing a slight attraction● Halogen diatomic molecules attract each other by

dispersion forces● Fluorine and chlorine are gases due to weak dispersion

forces. Bromine is a liquid due to more electrons and more dispersion forces

What is a hydrogen bond? ● Hydrogen bonds are a special case of dipole

interaction, formed when hydrogen is covalently bonded to a very electronegative atom, and is also weakly bonded to an unshared electron pair

● Example: water● They are the strongest of the intermolecular

forces, but only have about 5% the strength of an average covalent bond

● They always involve hydrogen, but can be found in other molecules besides water

● Form as a result of a highly polar bond and lack of shielding in the hydrogen

How do intermolecular attractions relate to Molecular properties?

● The physical properties of compounds relate to the type of bond it displays (largely ionic vs covalent), but the diverse range is largely due to varying intermolecular forces

● Melting and Boiling points of most molecular compounds are low compared to ionic compounds, why?

● Some covalent compounds exist where all the atoms are covalently bonded to each other - a network solid

○ Melting would require breaking of all the covalent bonds throughout the solid

● Example - diamond. Each carbon atom is bonded to four other carbons - all atoms are connected. Explains why diamonds is so hard, does not melt, and vaporizes to gas at 3500 C. Essentially carbo behaves as a single molecule