Covalent CompoundsChapter 8

Section 1, Covalent Bonds– Remember, ionic compounds are formed by gaining

and losing electrons– Atoms can also share electrons to become stable

•This is what occurrs in covalent bonding.

• we call compounds made by covalent bonds molecules

• The simplest form of a covalent compound is called a diatomic molecule

– They are made of atoms of the same element that are covalently bonded

– Remember, they always come in twos when they are the only element present and we learned the seven diatomic molecules in an earlier chapter• Br, I, N, Cl, H, O, F

– since the compound is formed from atoms of the same element, they have the same amount of attraction for the electrons and therefore neither will remove the electrons from the other and they must share the electrons

• covalent bonds also form between atoms of different elements

• polyatomic ions are actually formed between atoms of different elements in a covalent bond, then they bond to another ion by ionic bonding

– this means that compounds with polyatomic ions can actually be formed by two types of bonds

Covalent Bonds

Energy and Stability• most atoms are not stable by themselves

(noble gases are the exception to this rule)

• atoms become stable by getting a full outer level of electrons (8), and they do this by forming compounds with other atoms

• unbonded atoms have a high potential energy and they release some of this energy when they bond

• we call the distance between two bonded atoms at their minimum potential energy the bond length

• covalent bonds are flexible– the nuclei vibrate back and forth and the

distance between them constantly changes (the bond length is the average distance between the nuclei)

• the energy required to break the bond is called the bond energy

• the stronger the bond, the higher the bond energy and the shorter the bond length will be

Electronegativity and Covalent Bonds• Many covalent bonds form between atoms

of different elements– Every element has a different attraction for

electrons called electronegativity– this difference in attraction can determine the

type of bonds that will form– if the difference is great enough one atom will

take the electrons from the other, if not the electrons will be shared

– so we use electronegativty values to predict what type of bond will form between different elements (we subtract them)

2 Types of Covalent Bonds

• Even if the electronegativity difference is close enough to allow the sharing of electrons, the sharing is not always equal

1. nonpolar covalent – when the electrons are shared equally

2. polar covalent– when the electrons are shared unequally

• the atom with the higher electronegativity value holds the electrons more often than the other atoms

• since atoms in polar molecules do not share electrons equally the molecule ends up with partial charges on different sides

– the atom that has the electrons more often has a partial negative charge and the atom that does not have the electrons very often has a partial positive charge

Polar Molecules

Water is a Polar Molecule.

• a molecule with positive end and a negative end is called a dipole

• we use the symbom δ (delta)which means partial

– δ+ is partial positive, δ- is partial negative

• even though there are partial positives and negatives, the bond is still covalent because the electrons are still being shared (even if it is shared unequally)

• the greater the difference in the electronegativity values the greater the polarity and the greater the bond strenghth

– remember ionic bonds are the strongest

• Keep in mind that the boundaries between bond type are arbitrary, you also need to look at the atoms involved in the bond and their characteristics to know for sure

– it is a good indicator

Electronegativity and Bond Types• the general rule to predict the type of bond

that forms– if the electronegativity difference is between 0

and 0.5, the bond is probably nonpolar covalent– if the electronegativity difference is between 0.5

and 2.1, the bond is probably polar covalent– if the electronegativity difference is larger than

2.1, the bond is probably ionic

• also, remember, covalent compounds form between nonmetals and ionic bonds form between a metal and a nonmetal

Properties of Substances Depend on Bond Type• Metallic bonds are the result of the attraction

between the valence electrons in the atoms of a solid metal (metal and metal)

– all of the atoms are close together in a solid metal and all of the valence electrons are attracted to all of the atoms

– the electrons delocalize and move around all of the atoms in the metal• this results in the conducting of electrical

current

• Ionic bonds are formed by the gain and loss of electrons which results in the formation of cations and anions that are then attracted to each other

– each ion is held in place by many oppositely charged neighbors

– ionic bonds are very strong and very hard to break

– this gives them high melting and boiling points

– since they are made of ions they can also conduct electricity when in a liquid form which allows the ions free to move around

• Covalent bonds are held together by sharing electrons between 2 or more nonmetals

– the attractive forces between molecules is very small compared to the attractive forces between ions

Section 2, Drawing and Naming Molecules

• Lewis Electron-Dot Molecule– General

• all bond types involve the valence electrons in atoms

• we can use the Lewis Dot Diagrams to represent these electrons and atoms

• we can also use these diagrams to represent and visualize covalent bonds

Lewis Structures• remember, in Lewis Structures we use

the elements symbol to represent the nucleus and dots around the symbol to represent the valence electrons

– we draw the dots around each side

– we can predict the number of valence electrons from the pattern on the periodic table or we can write the electron configuration and count the outer electrons

– when we are drawing covalent bonds with Lewis structures, the shared electrons are written as dots between the 2 atoms that are sharing them• these shared electrons can also be shown as a

dash and they represent a single bond (a single pair of shared electrons)

– pairs that are not part of the bond are called unshared pairs or lone pairs

• The steps for drawing Lewis Structures for bonded atoms are as follows:

– draw the Lewis structures for each individual atom

– arrange the atoms so that the pairs that will be shared will line up

– if one atom will bond more than once it may be placed in the center of the other atoms

– distribute the dots so that each of the participating atoms has a full outer level of electrons (usually 8 unless it is hydrogen then it is 2) so that the octet rule is satisfied

– draw the bonds between the atoms by changing the shared pairs to a dash

– verify the structure by counting the electrons to see if all atoms have a full outer level

• The Lewis structure for polyatomic atoms is written the same way but some electrons will either be gained or lost to form an ion and we write the structure inside brackets with the charge on the outside.

Multiple Bonds• atoms can share more than one pair of electrons

in a covalent bond• a covalent bond that is formed by the sharing of 2

pairs of electrons is called a double bond• we represent this double bond with 2

pairs of dots between that atoms or 2 dashes

•a covalent bond that is formed by the sharing of 3 pairs of electrons is called a triple bond

•we represent this triple bond with 3 pairs of dots between that atoms or 3 dashes

•Molecules with double and triple bonds are referred to as unsaturated compounds.

• resonance structures are molecules that have 2 or more possible Lewis structures, neither structure is correct by itself, you must show all resonance structures

Naming Covalent Compounds• the method for naming covalent compounds is

similar to that for ionic compounds, except that we use prefixes to show the number of atoms for each element that is present in the molecule

• We write the less electronegative element first and we include a prefix in the name of this element only if there is more than one of them in the molecule

• We then write the more electronegative element and we include a prefix in its name that tells how many atoms of it are present in the molecule

• The prefixes are as follow:– 1 – mono, 2 – di, 3 – tri, 4- tetra, 5 – penta

• Ex. SO3 is called sulfur trioxide and N2O4 is called dinitrogen tetroxide

• Formulas for compounds are only used to tell us which atoms are present and how many of each

• We can use models to help us predict the three dimensional shape of molecules

• The model we will consider is called the VSEPR theory

• VSEPR stands for valence shell electron pair repulsion

• It is based on the fact that electrons repel each other, so this repulsion will affect the structure of the molecules

• Molecular ShapesShape # of Atoms # of Unshared

PairsExample

Linear 3 None CO2

Trigonal Planar

4 None BF3

Bent 3 1 or 2 H2O

Tetrahedral 5 None CH4

Trigonal Pyramidal

4 1 NH3

•Shape affects polarity

• The polarity of a molecule with 2 or more atoms depends on the shape of the molecule

• The atoms may have a partial charge, but they may be arranged so the charges cancel each other. In this case the molecule is considered nonpolar.

• The same charges need to be on the same side of a molecule to keep the molecules polarity.– For example, water. Because of its bent

shape, both of the δ+ hydrogen atoms are one side of the molecule while the δ- oxygen is on the opposite side of the molecule. Therefore the entire molecule is polar.

•Some molecules actually contain polar bonds within them, but because of the shape of the molecule those charges may be canceled out and the entire molecule is considered nonpolar. This molecule still contains polar bonds, but overall the molecule is nonpolar.•For instance, in carbon dioxide, the bonds between carbon and oxygen are polar, but the atom is arranged so that the charges cancel out and the entire molecule is considered nonpolar even though it contains polar bonds.

•The 2 δ- oxygens are on opposite sides of the molecule with the δ+ carbon in the middle. Since the atoms with similar charges are on opposite sides of the molecule with the other charge in between, the charges negate each other and the molecule is nonpolar.