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Chem.Unit 4 and Study guide
Objectives written  below refer to notes given out in class for studying
Chapter 6: Chemical Bonding
Section 6-1: Introduction to Chemical Bonding
1. Define chemical bond.
2. Explain why most atoms form chemical bonds.
3. Describe ionic and covalent bonding.
4. Explain why most chemical bonding is neither purely ionic nor purely covalent.
5. Classify bonding type according to electronegativity differences.
Section 6-2: Covalent Bonding and Molecular Compounds
1. Define molecule and molecular formula.
2. Explain the relationships between potential energy, distance between approaching atoms, bond length, and bond energy.
3. State the octet rule.
4. List the six basic steps used in writing Lewis structures.
5. Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.
6. Explain why scientists use resonance structures to represent some molecules.
Section 6-3: Ionic Bonding and Ionic Compounds
1. Compare and contrast a chemical formula for a molecular compound with one for an ionic compound.
2. Discuss the arrangements of ions in crystals.
3. Define lattice energy and explain its significance.
4. List and compare the distinctive properties of ionic and molecular compounds.
5. Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information.
Section 6-4: Metallic Bonding
1. Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors.
2. Explain why metal surfaces are shiny.
3. Explain why metals are malleable and ductile but ionic-crystalline compounds are not.
Section 6-5: Molecular Geometry
1. Explain VSEPR theory.
2. Predict the shapes of molecules or polyatomic ions using VSEPR theory.
3. Explain how the shapes of molecules are accounted for by hybridization theory.
4. Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces.
5. Explain what determines molecular polarity.

Basic Concepts of Chemical Bonding
Lewis Symbols and the Octet Rule
Why are some substances chemically bonded molecules and others are an association of ions?
   depends upon the electronic structures of the atoms
   nature of the chemical forces within the compounds
A broad classification of chemical forces:
Ionic bonds
Covalent bonds
Metallic bonds
Ionic bonds - electrostatic forces that exist between ions of opposite charge
   typically involves a metal with a nonmetal
Covalent bonds - results from the sharing of electrons between two atoms
   typically involves one nonmetallic element with another
Metallic bonds
   found in solid metals (copper, iron, aluminum)
   each metal bonded to several neighboring groups
   bonding electrons free to move throughout the 3-dimensional structure
Lets look at the preferred arrangements of electrons in atoms when they form chemical compounds
Lewis Symbols and the Octet Rule
Valence electrons reside in the outer shell and are the electrons which are going to be involved in chemical interactions and bonding (valence comes from the Latin valere, "to be strong").
Electron-dot symbols (Lewis symbols):
   convenient representation of valence electrons
   allows you to keep track of valence electrons during bond formation
   consists of the chemical symbol for the element plus a dot for each valence electron
Electron configuration is [Ne]3s23p4, thus there are six valence electrons. Its Lewis symbol would therefore be:
   The dots (representing electrons) are placed on the four sides of the atomic symbol (top, bottom, left, right)
   Each side can accommodate up to 2 electrons
   The number of valence electrons in the table below is the same as the column number of the element in the periodic table (for representative elements only)
Atoms often gain, lose, or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table
Because all noble gasses (except He) have filled s and p valence orbitals (8 electrons), many atoms undergoing reactions also end up with 8 valence electrons. This observation has led to the Octet Rule:
Atoms tend to lose, gain, or share electrons until they are surrounded by 8 valence electrons
Note: there are many exceptions to the octet rule (He and H, for example), but it provides a useful model for understanding the basis of chemical bonding.
Basic Concepts of Chemical Bonding
Ionic Bonding

Ionic Bonding
Sodium metal reacts with chlorine gas in a violently exothermic reaction to produce NaCl (composed of Na+ and Cl- ions):
2Na(s) + Cl2(g) -> 2NaCl(s)
These ions are arranged in solid NaCl in a regular three-dimensional arrangement (or lattice):
   The chlorine has a high affinity for electrons, and the sodium has a low ionization potential. Thus the chlorine gains an electron from the sodium atom. This can be represented using electron-dot symbols (here we will consider one chlorine atom, rather than Cl2):

   The arrow indicates the transfer of the electron from sodium to chlorine to form the Na+ metal ion and the Cl- chloride ion. Each ion now has an octet of electrons in its valence shell:
Na+ 2s22p6
Cl- 3s23p6
Energetics of Ionic Bond Formation
The formation of ionic compounds (like the addition of sodium metal and chlorine gas to form NaCl) are usually extremely exothermic.
The loss of an electron from an element:
   Always endothermic (takes energy to strip the e' from the atom)
   Na(g) -> Na+(g) + 1e- DH = 496 kJ/mol
The gain of an electron by a nonmetal:
   Generally exothermic (energy released)
   Cl(g) + 1e- -> Cl-(g) DH = -349 kJ/mol
The formation of NaCl from Na and Cl would thus require the input of 147 kJ/mol. However, it appears to be a highly exothermic reaction.
Ionic compounds are stable due to the attraction between unlike charges:
   The ions are drawn together
   Energy is released
   Ions form solid lattice
Lattice energy:
the energy required to separate completely a mole of a solid ionic compound into its gaseous ions
It is a measure of just how much stabilization results from the arranging of oppositely charged ions in an ionic solid.
To completely break up a salt crystal:
NaCl(s) -> Na+(g) + Cl-(g) DHlattice = +788 kJ/mol
Thus, -788 kJ/mol is given off as heat energy when 1 mol of NaCl is incorporated into the salt lattice.
So, forming the ions from Na(g) and Cl(g) requires the input of +147 kJ/mol, these ions incorporate into the salt lattice liberating -788 kJ/mol, for an overall highly exothermic release of -641 kJ/mol.
The magnitude of the lattice energy depends upon the charges of the ions, their size and the particular lattice arrangement.
The potential energy of two interacting charged particles is:
Q1 = charge on first particle
Q2 = charge on second particle
d = distance between centers of particles
k = 8.99 x 109 J m/C2
Thus, the interaction increases:
   As the charges increase
   As the two charges are brought closer together
The minimum distance between oppositely charged ions is the sum of the atomic (ionic) radii. Although atomic radii do vary, it is not over a considerable range, thus, the attraction between two ions is determined primarily by the charge of the ions.
Electron configuration of ions
How does the energy released in lattice formation compare to the energy required to strip away another electron from the Na+ ion?

Polyatomic ions
In polyatomic ions, two or more atoms are bound together by covalent (chemical) bonds. They form a stable grouping which carries a charge (positive or negative). The group of atoms as a whole acts as a charged species in forming an ionic compound with an oppositely charged ion.
Basic Concepts of Chemical Bonding
Sizes of Ions

Sizes of Ions
Sizes of ions influence:
   packing of ions in ionic lattices, and therefore, the lattice energy
   biological recognition - some ions can pass through certain membrane channels, others may be too large
The size of an ion is influenced by:
   nuclear charge
   number of electrons
   valence orbitals
   formed by removing one or more valence electrons
   vacates the most spatially extended orbitals
   decreases the total electron-electron repulsion in the outer orbital
Cations are therefore smaller than the parent atom
   formed by addition of one or more valence electrons
   fills in the most spacially extended orbitals
   increases electron-electron repulsion in outer orbital
Anions are therefore larger than the parent atom
For ions of the same charge (e.g. in the same group) the size increases as we go down a group in the periodic table
As the principle quantum increases the size of both the parent atom and the ion will increase
How does the nuclear charge affect ion size?
Consider the following collection of ions:
O2- 10 8
F- 10 9
Na+ 10 11
Mg2+ 10 12
Al3+ 10 13

Each ion:
   contains the same number of electrons (10; with configuration 1s22s22p6) and are thus termed a collection of isoelectronic ions
   varies in the nuclear charge
The radius of each ion decreases with an increase in nuclear charge:

Oxygen and fluorine precede neon and are nonmetals, sodium, magnesium and aluminum come after neon and are metals.

Basic Concepts of Chemical Bonding
Covalent Bonding

Covalent Bonding
Ionic substances:
   usually brittle
   high melting point
   organized into an ordered lattice of atoms, which can be cleaved along a smooth line
the electrostatic forces organize the ions of ionic substances into a rigid, organized three-dimensional arrangement
The vast majority of chemical substances are not ionic in nature
   gases and liquids, in addition to solids
   low melting temperatures
G.N. Lewis
reasoned that an atom might attain a noble gas electron configuration by sharing electrons
A chemical bond formed by sharing a pair of electrons is called a covalent bond
The diatomic hydrogen molecule (H2) is the simplest model of a covalent bond, and is represented in Lewis structures as:

The shared pair of electrons provides each hydrogen atom with two electrons in its valence shell (the 1s) orbital.
In a sense, it has the electron configuration of the noble gas helium
When two chlorine atoms covalently bond to form Cl2, the following sharing of electrons occurs:

Each chlorine atom shared the bonding pair of electrons and achieves the electron configuration of the noble gas argon.
In Lewis structures the bonding pair of electrons is usually displayed as a line, and the unshared electrons as dots:
   The shared electrons are not located in a fixed position between the nuclei. In the case of the H2 compound, the electron density is concentrated between the two nuclei:
The two atoms are bound into the H2 molecule mainly due to the attraction of the positively charged nuclei for the negatively charged electron cloud located between them
For the nonmetals (and the 's' block metals) the number of valence electrons is equal to the group number:
Bonds needed to form valence octet

Examples of hydride compounds of the above elements (covalent bonds with hydrogen:

Thus, the Lewis bonds successfully describe the covalent interactions between various nonmetal elements
Multiple bonds
The sharing of a pair of electrons represents a single covalent bond, usually just referred to as a single bond
In many molecules atoms attain complete octets by sharing more than one pair of electrons between them.
Two electron pairs shared a double bond
Three electron pairs shared a triple bond

Because each nitrogen contains 5 valence electrons, they need to share 3 pairs to each achieve a valence octet.
   N2 is fairly inert, due to the strong triple bond between the two nitrogens
   The N - N bond distance in N2 is 1.10 Å (fairly short)
From a study of various Nitrogen containing compounds bond distance as a function of bond type can be summarized as follows:
   N-N 1.47Å
   N=N 1.24Å
   N=N 1.10Å
As a general rule, the distance between bonded atoms decreases as the number of shared electron pairs increases

Basic Concepts of Chemical Bonding
Bond Polarity and Electronegativity

Bond Polarity and Electronegativity
The electron pairs shared between two atoms are not necessarily shared equally
Extreme examples:
1. In Cl2 the shared electron pairs is shared equally
2. In NaCl the 3s electron is stripped from the Na atom and is incorporated into the electronic structure of the Cl atom - and the compound is most accurately described as consisting of individual Na+ and Cl- ions
For most covalent substances, their bond character falls between these two extremes
Bond polarity is a useful concept for describing the sharing of electrons between atoms
   A nonpolar covalent bond is one in which the electrons are shared equally between two atoms
   A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. If this relative attraction is great enough, then the bond is an ionic bond
A quantity termed 'electronegativity' is used to determine whether a given bond will be nonpolar covalent, polar covalent, or ionic.
Electronegativity is defined as the ability of an atom in a particular molecule to attract electrons to itself
(the greater the value, the greater the attractiveness for electrons)
Electronegativity is a function of:
   the atom's ionization energy (how strongly the atom holds on to its own electrons)
   the atom's electron affinity (how strongly the atom attracts other electrons)
(Note that both of these are properties of the isolated atom)
For example, an element which has:
   A large (negative) electron affinity
   A high ionization energy (always endothermic, or positive for neutral atoms)
   Attract electrons from other atoms
   Resist having its own electrons attracted away
Such an atom will be highly electronegative
Fluorine is the most electronegative element (electronegativity = 4.0), the least electronegative is Cesium (notice that are at diagonal corners of the periodic chart)

General trends:
   Electronegativity increases from left to right along a period
   For the representative elements (s and p block) the electronegativity decreases as you go down a group
   The transition metal group is not as predictable as far as electronegativity
Electronegativity and bond polarity
We can use the difference in electronegativity between two atoms to gauge the polarity of the bonding between them
Electronegativity Difference
4.0 - 4.0 = 0
4.0 - 2.1 = 1.9
4.0 - 1.0 = 3.0
Type of Bond
Nonpolar covalent
Polar covalent
Ionic (non-covalent)
   In F2 the electrons are shared equally between the atoms, the bond is nonpolar covalent
   In HF the fluorine atom has greater electronegativity than the hydrogen atom.
The sharing of electrons in HF is unequal: the fluorine atom attracts electron density away from the hydrogen (the bond is thus a polar covalent bond)
The H-F bond can thus be represented as:

   The 'd+' and 'd-' symbols indicate partial positive and negative charges.
   The arrow indicates the "pull" of electrons off the hydrogen and towards the more electronegative atom
   In lithium fluoride the much greater relative electronegativity of the fluorine atom completely strips the electron from the lithium and the result is an ionic bond (no sharing of the electron)
A general rule of thumb for predicting the type of bond based upon electronegativity differences:
   If the electronegativities are equal (i.e. if the electronegativity difference is 0), the bond is non-polar covalent
   If the difference in electronegativities between the two atoms is greater than 0, but less than 2.0, the bond is polar covalent
   If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic

Basic Concepts of Chemical Bonding
Drawing Lewis Structures
Drawing Lewis Structures
The general procedure...
1. Sum the valence electrons from all atoms
   Use the periodic table for reference
   Add an electron for each indicated negative charge, subtract an electron for each indicated positive charge
2. Write the symbols for the atoms to show which atoms are attached to which, and connect them with a single bond
   You may need some additional evidence to decide bonding interactions
   If a central atom has various groups bonded to it, it is usually listed first: CO32-, SF4
   Often atoms are written in the order of their connections: HCN
3. Complete the octets of the atoms bonded to the central atom (H only has two)
4. Place any leftover electrons on the central atom (even if it results in more than an octet)
5. If there are not enough electrons to give the central atom an octet, try multiple bonds (use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds)
Draw the Lewis structure of phosphorous trichloride (PCl3)
This is an example of a central atom, P, surrounded by chlorine atoms
1. We will have 5(P) plus 21 (3*7, for Cl), or 26 total valence electrons
2. The general symbol, starting with only single bonds, would be:

3. Completing the octets of the Cl atoms bonded to the central P atom:

4. This gives us a total of (18 electrons) plus the 6 in the three single bonds, or 24 electrons total. Thus we have 2 extra valence electrons which are not accounted for. We will place them on the central element:
5. The central atom now has an octect, and there is no need to invoke any double or triple bonds to achieve an octet for the central atom. We are finished.
Draw the Lewis structure for the NO+ ion
1. We will have 5 (N) plus 6 (O) minus 1 (1+ ion), or 10 valence electrons
2. The general structure starting only with single bonds would be:
3. Completing the octet of the O bonded to N:
4. This gives us a total of 6 plus 2 for the single bond, or 8 electrons. There are 2 unaccounted for electrons and we will place them on the N:
5. There are only 4 atoms on the N atom, not enough for an octet, so lets try a double bond between the N and O:

The oxygen still has an octet, but the N only has 6 valence electrons, so lets try a triple bond:

Each atom now has a valence octet. We are finished.
The brackets with the + symbol are used to indicate that this is an ion with a net charge of 1+
Formal Charge
In some cases we can draw several different Lewis structures which fulfill the octet rule for a compound. Which one is the most reasonable?
One method is to tabulate the valence electrons around each atom in a Lewis structure to determine the formal charge. The formal charge is the charge that an atom in a molecule would have if we considered each atom to have the same electronegativity in a compound.
To calculate formal charge, assign electrons to their respective atoms as follows:
All of the unshared electrons are assigned to the atom on which they are found
The bonding electrons are divided up equally between each atom involved in the bond
The number of valence electrons expected in the isolated atom is compared to the actual number of electrons assigned from the Lewis structure:
The formal charge equals the number of valence electrons in the isolated atom, minus the number of electrons assigned in the Lewis structure
Example: Carbon Dioxide (CO2)
Carbon has 4 valence electrons
Each oxygen has 6 valence electrons, therefore our Lewis structure of CO2 will have 16 electrons:

One way we could draw the Lewis structure is:

Another way we could draw the Lewis structure is:

Both structures fulfill the octet rule. But what are the formal charges?

Which structure is correct? In general, when several Lewis structures can be drawn the most stable structure is the one in which:
   The formal charges are the smallest
   Any negative charge is found on the most electronegative atom
In the above case, the second structure is the one with the smallest formal charges (i.e. 0 on all the atoms).
   Furthermore, in the first possible Lewis structure the carbon has a formal charge of 0 and one of the oxygens it is bonded to has a formal charge of +1.
   Oxygen is more electronegative than Carbon, so this situation would seem unlikely.
It is important to remember that formal charges do not represent the actual charges on the atoms. Actual charges are determined by the electronegativity of the atoms involved.
Basic Concepts of Chemical Bonding
Resonance Structures
Resonance Structures
The Lewis structure of ozone (O3)
1. Sum of valence electrons = (6*3) = 18
2. Drawing the bond connectivities:
3. Complete the octets of the atoms bonded to the central atom:

4. Place any leftover electrons (18-16 = 2) on the central atom:

5. Does the central atom have an octet?
   NO, it has 6 electrons
   Add a multiple bond (first try a double bond) to see if the central atom can achieve an octet:
6. Does the central atom have an octet?
   YES, we are done
   Ozone would appear to have one single bond, and one double bond
However... known facts about the structure of ozone
   The bond lengths between the central oxygen and the other two oxygens are identical:

   We would expect that if one bond was a double bond that it should be shorter than the other (single) bond
   Since all the atoms are identical (oxygens) which atom is chosen for the double bond?

These Lewis structures are equivalent except for the placement of the electrons (i.e. the location of the double bond)
Equivalent Lewis structures are called resonance structures, or resonance forms
The correct way to describe ozone as a Lewis structure would be:

This indicates that the ozone molecule is described by an average of the two Lewis structures (i.e. the resonance forms)
The important points to remember about resonance forms are:
   The molecule is not rapidly oscillating between different discrete forms
   There is only one form of the ozone molecule, and the bond lengths between the oxygens are intermediate between characteristic single and double bond lengths between a pair of oxygens
   We draw two Lewis structures (in this case) because a single structure is insufficient to describe the real structure
The Nitrate (NO3-) ion:
1. Count up the valence electrons: (1*5) + (3*6) + 1(ion) = 24 electrons
2. Draw the bond connectivities:
3. Add octet electrons to the atoms bonded to the center atom:
4. Place any leftover electrons (24-24 = 0) on the center atom:
5. Does the central atom have an octet?
   NO, it has 6 electrons
   Add a multiple bond (first try a double bond) to see if the central atom
   can achieve an octet:

6. Does the central atom have an octet?
   Are there possible resonance structures? YES

Note: We would expect that the bond lengths in the NO3- ion to be somewhat shorter than a single bond
Basic Concepts of Chemical Bonding
Exceptions to the Octet Rule

Exceptions to the Octet Rule
There are three general ways in which the octet rule breaks down:
1. Molecules with an odd number of electrons
2. Molecules in which an atom has less than an octet
3. Molecules in which an atom has more than an octet
Odd number of electrons
Draw the Lewis structure for the molecule nitrous oxide (NO):
1. Total electrons: 6+5=11
2. Bonding structure:
3. Octet on "outer" element:
4. Remainder of electrons (11-8 = 3) on "central" atom:
5. 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.
We appear unable to get an octet around each atom
Less than an octet (most often encountered with elements of Boron and Beryllium)
Draw the Lewis structure for boron trifluoride (BF3):
1. Add electrons (3*7) + 3 = 24
2. Draw connectivities:
3. Add octets to outer atoms:
4. Add extra electrons (24-24=0) to central atom:
5. Does central electron have octet?
   NO. It has 6 electrons
   Add a multiple bond (double bond) to see if central atom can achieve an octet:
6. The central Boron now has an octet (there would be three resonance Lewis structures)
   In this structure with a double bond the fluorine atom is sharing extra electrons with the boron.
   The fluorine would have a '+' partial charge, and the boron a '-' partial charge, this is inconsistent with the electronegativities of fluorine and boron.
   Thus, the structure of BF3, with single bonds, and 6 valence electrons around the central boron is the most likely structure
BF3 reacts strongly with compounds which have an unshared pair of electrons which can be used to form a bond with the boron:

More than an octet (most common example of exceptions to the octet rule)
PCl5 is a legitimate compound, whereas NCl5 is not.

Expanded valence shells are observed only for elements in period 3 (i.e. n=3) and beyond
   The 'octet' rule is based upon available ns and np orbitals for valence electrons (2 electrons in the s orbitals, and 6 in the p orbitals)
   Beginning with the n=3 principle quantum number, the d orbitals become available (l=2)
The orbital diagram for the valence shell of phosphorous is:

Third period elements occasionally exceed the octet rule by using their empty d orbitals to accommodate additional electrons
Size is also an important consideration:
   The larger the central atom, the larger the number of electrons which can surround it
   Expanded valence shells occur most often when the central atom is bonded to small electronegative atoms, such as F, Cl and O.
Draw the Lewis structure for ICl4-
1. Count up the valence electrons: 7+(4*7)+1 = 36 electrons
2. Draw the connectivities:

3. Add octet of electrons to outer atoms:

4. Add extra electrons (36-32=4) to central atom:

5. The ICl4- ion thus has 12 valence electrons around the central Iodine (in the 5d orbitals)
Basic Concepts of Chemical Bonding
Strengths of Covalent Bonds
Strengths of Covalent Bonds
The stability of a molecule is a function of the strength of the covalent bonds holding the atoms together.
How do we measure the strength of a covalent bond?
Bond-dissociation energy (i.e. "bond energy")
The more stable a molecule (i.e. the stronger the bonds) the less likely the molecule is to undergo a chemical reaction
Bond Energies and the Enthalpy of Reacti

Basic Concepts of Chemical Bonding
Oxidation Numbers
Oxidation Numbers
When a covalent bond forms between two atoms with different electronegativities the shared electrons in the bond lie closer to the more electronegative atom:

   The oxidation number of an atom is the charge that results when the electrons in a covalent bond are assigned to the more electronegative atom
   It is the charge an atom would possess if the bonding were ionic
In HCl (above) the oxidation number for the hydrogen would be +1 and that of the Cl would be -1
in oxidation numbers we write the sign first to distinguish them from ionic (electronic) charges
Oxidation numbers do not refer to real charges on the atoms, except in the case of actual ionic substances.
Oxidation numbers can be determined using the following rules:
1. The oxidation number for an element in its elemental form is 0 (holds true for isolated atoms and elemental substances which bond identical atoms: e.g. Cl2, etc)
2. The oxidation number of a monoatomic ion is the same as its charge (e.g. oxidation number of Na+ = +1, and that of S2- is -2)
3. In binary compounds (two different elements) the element with greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds of the element (e.g. in the compound PCl3 the chlorine is more electronegative than the phosphorous. In simple ionic compounds Cl has an ionic charge of 1-, thus, its oxidation state is -1)
4. The sum of the oxidation numbers is zero for an electrically neutral compound and equals the overall charge for an ionic species.
5. Alkali metals exhibit only an oxidation state of +1 in compounds
6. Alkaline earth metals exhibit only an oxidation state of +2 in compounds
The chlorine is more electronegative and so its oxidation number is set to -1. The overall molecule is neutral, so the oxidation number of P, in this case, is +3.
The oxygen is more electronegative and receives an oxidation number of -2. The overall molecule has a net charge of 2- (an overall oxidation number of ­2), therefore, the C must have an oxidation state of +4, i.e. (3*-2) + 'C' = -2.
Examples of Sulfur
Sulfur (2.5) is more electronegative than hydrogen (2.1), thus it has an oxidation number of -2. The hydrogen will have an oxidation number of +1.
This is an elemental form of sulfur, and thus would have an oxidation number of 0.
Chlorine (3.0) is more electronegative than sulfur (2.5), thus it has an oxidation number of -1. The sulfur thus has an oxidation number of +2.
Sodium (alkali metal) always has an oxidation number of +1. The oxygen (3.5) is more electronegative than sulfur (2.5), thus the oxygen would have an oxidation number of -2. The sulfur would therefore have an oxidation number of +4.
The oxygen is more electronegative and thus has an oxidation number of -2. The sulfur thus has an oxidation number of +6.
   Sulfur exhibits a variety of oxidation numbers (-2 to +6)
   In general the most negative oxidation number corresponds to the number of electrons which must be added to give an octet of valence electrons
   The most positive oxidation number corresponds to a loss of all valence electrons
Oxidation Numbers and Nomenclature
Compounds of the alkali (oxidation number +1) and alkaline earth metals (oxidation number +2) are typically ionic in nature.
Compounds of metals with higher oxidation numbers (e.g. tin +4) tend to form molecular compounds
   In ionic and covalent molecular compounds usually the less electronegative element is given first.
   In ionic compounds the names are given which refer to the oxidation (ionic) state
   In molecular compounds the names are given which refer to the number of molecules present in the compound
Ionic Molecular
MgH2 magnesium hydride H2S dihydrogen sulfide
FeF2 iron(II) fluoride OF2 oxygen difluoride
Mn2O3 manganese(III) oxide Cl2O3 dichlorine trioxide