Chemical Bonding

Posted on February 25, 2009

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Ionic compounds form when metals and nonmetals react
The attraction between positive and negative ions is called an ionic bond
The ionic compounds form because the potential energy of the system decreases
Consider the example of sodium chloride
The energy change when NaCl forms can be calculated using the ionization energy (IE) of sodium, the electron affinity (EA) of chlorine, and the lattice energy of NaCl
Starting from 1 mole of gas phase atoms:
Na(g)  Na+(g) + e- +495.4 kJ (IE of sodium)
Cl(g) + e-  Cl-(g) -348.8 kJ (EA of chlorine)
Na+(g)+Cl-(g)NaCl(s) -787.0 kJ (-lattice energy)
Net: -640.4 kJ
It turns out that for any ionic compound, the chief stabilizing influence is the lattice energy
The size of the lattice energy depends on ion size and charge
The lattice energy increases with charge because the ions attract each other more strongly
Example: KCl (709 kJ) vs CaO (3401 kJ)
Smaller ions have larger lattice energies because they get closer together
Example: NaCl (778 kJ) vs KCl (709 kJ) and LiF (1033 kJ) vs LiCl (845 kJ)
The lattice energy can be calculated using a Born-Haber cycle
Noble gas configurations are very stable and can be useful in predicting ion charges
Consider the case of sodium:
Na(g)  Na+(g)+e- IE= 466 kJ/mol
Na+(g) Na2+(g)+e- IE=4563 kJ/mol
Na 1s22s22p63s1
Na+ 1s22s22p6 (noble gas core)
Na2+ 1s22s22p5
Formation of Na+ is relatively inexpensive
Na2+ doesn’t (ordinarily) form because breaking into the noble gas core costs a to much energy
All noble gases (except He) have 8 valence electrons
This is called an octet of electrons
Most of the representative elements tend to gain or lose electrons until they have achieved the configuration of the nearest noble gas
For example: Na and K lose electrons to achieve an octet of electrons while Cl and O gain electron to achieve an octet of electrons
The octet rule works best for ionic compounds of Group IA and IIA metals from Period 3 down and for the anions of the nonmetals
It fails for Li and Be because they achieve the He (1s2) electron configuration
It also doesn’t work for hydrogen which can form H- (electron configuration: 1s2) when it reacts with very reactive metals
The octet rule doesn’t work well for transition metals and post transition metals
For these cations:
The first electrons lost by an atom or ion are always those from the outer shell (with the largest value of n)
Within a given shell: the f (subshell) is emptied before the d, which is emptied before the p, which is emptied before the s
Consider the case of iron:
Neutral: Fe [Ar]3d64s2
Loss of 4s electrons: Fe2+ [Ar]3d6
Loss of a 3d electron: Fe3+ [Ar]3d5
Many transition elements form multiple cations (like iron)
Often, one of the cations has a charge of +2
The relative stability of the ions formed is difficult to predict
Lewis symbols provide a convenient way to keep track of valence electrons
In this notation the symbol of the element is surrounded by dots (or similar symbols) that represent the atom’s valence electrons
All the elements in a group have a similar Lewis symbol because they have the same number of valence electrons
Ions are treated in a similar fashion
Many substances comprised only of nonmetals occur as molecules
Molecules involve electron sharing
Covalent bonds are characterized by their bond distance, or average distance between the nuclei, and the bond energy, or amount of energy released when the bond forms
Lewis symbols can be used to represent the covalent or electron pair bond
Both hydrogens are considered to have two electrons
For simplicity, electron pair bonds are usually represented by a dash
Example: hydrogen molecule is represented as H-H
Formulas drawn with Lewis symbols are called Lewis formulas or Lewis structures
The term structural formula is also used because it shows how the atoms in the molecule are attached to each other
Many molecules obey the octet rule:
When atoms form covalent bonds, they tend to share sufficient electrons so as to achieve an outer shell having eight electrons
In most of their covalently bonded compounds, the number of covalent bonds formed by carbon, nitrogen, and oxygen are four, three, and two, respectively
One shared pair of electrons is called a single bond
Double and triple bonds are also common:
Organic compounds will frequently be used as examples later in the text
In general, organic compounds are held together with covalent bonds
The simplest hydrocarbons are the alkanes with the general formula CnH2n+2
The first three alkanes are methane, ethane, and propane
In condensed form they are written:
methane: CH4
ethane: CH3CH3
propane: CH3CH2CH3
Things get more complicated starting with alkanes containing four carbons
Hydrocarbons that contain one double bond have the general formula CnH2n and are called alkenes
Hydrocarbons that contain one triple bond have the general formula CnH2n-2 and are called alkynes
Most organic compounds contain elements in addition to carbon and hydrogen
These are considered to be hydrocarbon derivatives
Using the symbol “R” to represent any hydrocarbon fragment (such as CH3-, or CH3CH2-) important families include:
Ball and stick models are common
When two identical atoms form a covalent bond each atom has an equal share of the bond’s electron pair
When different kinds of atoms combine, one nuclei usually attracts the electrons in the bond more strongly
The magnitude of the polarity is expressed in terms of the dipole moment
Dipole moments are frequently reported in units of Debye (D)
The dipole moments and bond lengths for some diatomic molecules are:
Electronegativity is the term used to describe the relative attraction of an atom for the electrons in a bond
The element with the larger electronegativity will carry the partial negative charg
The difference in electronegativity provides an estimate for the degree of polarity of the bond
There is no sharp dividing line between ionic and covalent bonding: ionic bonding and nonpolar covalent bonding represent the extremes
A bond is mostly ionic when the electronegativity difference between the two atoms is large
The degree of polarity, or ionic character, varies continuously with the electronegativity difference
In general, electronegativity increases bottom to top in a group and left to right in a period
Metal reactivity refers to the tendency of the metal to undergo oxidation
The lower the electronegativity the easier a metal is to oxidize
For nonmetals, reactivity is usually gauged by the ability to act as an oxidizing agent
In general, the oxidizing ability of nonmetals increases from left to right in a period and bottom to top in a group
This makes fluorine, found in the upper right of the periodic table, the strongest oxidizing agent
Single displacement reactions may be predicted from nonmetal reactivity
Consider the halogens (Group VIIA): a halogen as an element will oxidize the anion of any halogen below it
F2 will oxidize Cl-, Br-, and I-
Example: F2 + 2Cl-  2F- + Cl2
Cl2 will oxidize Br- and I-
Example: Cl2 + 2Br-  2Cl- + Br2
Br2 will oxidize I-
Example: Br2 + 2I-  2Br- + I2
Lewis structures are useful because they give a simple way to describe the structure of molecules
Not all structures obey the octet rule
Most nonmetals beyond Period 2 form structures with more than eight electrons
Examples: PCl5 and SF6
In some compounds the central atom has less than eight electrons
Common examples include compounds of beryllium and boron
Examples: BeCl2 and BCl3
Lewis structures describe how atoms share electrons in chemical bonds
The bond length and bond energy are related to the number of electron pairs shared between to atoms
For bonds between the same elements the bond length and bond energy depend on the bond order
The bond order is the number of pairs of electrons shared between two atoms
A single bond has bond order of 1; a double bond a bond order of 2; and a triple bond a bond order of 3
The bond order is a measure of the amount of electron density in a bond
More electron density gives a stronger bond
Consider the average bond lengths and bond energies for carbon-carbon bonds:
Bond lengths and bond energies are obtained from experiment
The preferred Lewis structure is the one that best fits the experimental data
The preferred Lewis structure for sulfuric acid violates the octet rule:
Structure I obeys the octet rule, but is not consistent with experiment
Structure II violates the octet rule, but is consistent with experiment
Structure II is the preferred Lewis structure
Formal charge is the apparent charge on an atom
The formal charge on a atom is calculated by subtracting the number of valence electrons assigned to it in a Lewis structure from the number of valence electrons in an isolated atom
Consider the sulfur atoms in the two structures for sulfuric acid:
Structure I: formal charge on S = 6 – (4 + 0) = +2
Structure II: formal charge on S = 6 – (6 + 0) = 0
When several Lewis structures are possible, those with the smallest formal charges are the most stable and preferred
Note that the formal charges for all atoms in a Lewis structure sum to the charge on the species
Some molecules and ions are not well represented by a single Lewis structure
Consider the case of the formate ion
Experiment gives a single carbon-oxygen bond length
A combination of structures is needed to describe this ion
These are called resonance structures and the ion is said to be a resonance hybrid of the contributing structures
Two resonance structures are required for the formate ion because two equivalent carbon-oxygen double bonds can be formed
Note that three resonance structures would be required to represent SO3
The total energy of a resonance hybrid is lower in energy than any one of its resonance structures
This energy lowering is called the resonance energy
Consider the formation of the ammonium ion from ammonia and a hydrogen ion in solution
The nitrogen donates both of the electrons when forming the bond to H+
This is called a coordinate covalent bond
The concept of a coordinate bond can be useful when trying to understand what happens to atoms in reactions
For example, addition compounds involve coordinate covalent bonds and can result when two small molecules “join”

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