Relative acidities of functional groups predict their reactivity. The acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid–base reactions. By examining the Ka or the pKa it is possible to determine the relative acidity of a compound. The larger the value of pKa, the smaller the extent of dissociation at any given pH. Another way of saying this is the smaller the pKa, the stronger the acid or the more likely the compound is to dissociate a hydrogen. Acids with a pKa value of less than about −2 are said to be strong acids; the dissociation of a strong acid is such that the concentration of the undissociated acid is too small to be measured. A selection of functional groups and their respective pKa are shown below.

In aqueous solution, the equilibrium of acid dissociation can be represented as shown below where HA represents the acid A- the conjugate base.

Polyprotic acids are acids that can lose more than one proton. The constant for dissociation of the first proton may be denoted as Ka1 and the constants for dissociation of successive protons as Ka2, Ka3. Phosphoric acid, H₃PO₄, is an example of a polyprotic acid capable of losing three protons.

 

Groups of atoms (also known as moities) or bonds within molecules that are responsible for the characteristic chemical reactions of those molecules are known as functional groups. Functional groups are atoms connected by covalent bonds. It is common to see molecules comprised mainly of a carbon backbone with functional groups attached to the chain. The functional group gives the molecule its properties; functional groups are centers of chemical reactivity.

Relative reactivity can be modified by nearby functional groups. For example, a carboxylic acid can become more acidic due to the electronegativity of nearby atoms. For example, acetic acid’s pka is not as low as the pka of trichloroacetic acid due the presence of hydrogens as opposed to chlorine atoms. Inductive effects refer to those electronic effects of an atom or functional group can contribute through single bonds such as saturated (sp3) carbon atoms. Atoms or functional groups that are electronegative relative to hydrogen such as the halogens, oxygen, nitrogen withdraw electron density through the single bond structure of a compound and can assist in the stabilization of negative charge that may form in reactions. One such reaction where -Cl groups can have a stabilizing (enhancing) effect is the ionization of acids.

The terms resonance and induction refer to the electronic effects that atoms or functional groups may have within a compound. These effects are dependent on the valence, bonding order and electronegativity of atoms, as well as the molecular geometry. Resonance may be defined as bonding or sharing of electrons between more than two atoms. Typical covalent and ionic bonding involves sharing (covalent) or transferring (ionic) electron pairs between two atoms. An example of resonance is shown below. Note the direction of the arrow (from partial negative to partial positive or from lone pair to positive charge)

 

 

The structure of molecules determine their function. The reactivity of a given atom is specified by the outermost shell of electrons, called the valence electrons. The valence electrons determine the structure of a molecule.

In order to understand the structure of molecules, scientists consider the position and movement of electrons, specifically the valence electrons-the electrons in the outermost shell available to react and form bonds. A Lewis structure reveals the valence electrons arranged to satisfy the octet rule. Explore the steps to generating a Lewis structure for nitrate.

Electronegativity measures the force of an atom’s attraction for electrons. Increasing from left to right within a period of the Periodic Table due to increasing positive charge on the nucleus leading to a stronger attraction for electrons as well as increasing from bottom to top as the decreasing distance of the valence electrons from the nucleus, electronegativities of atoms help predict the type of bonds formed.  If the difference in electronegativity between two atoms is less than 0.5, the atoms form a nonpolar covalent bond. If the difference in electronegativity is from 0.5 to 1.9, the atoms form a polar covalent bond. Finally, if the electronegativity difference is greater than 1.9, the atoms form an ionic bond.

Once atoms form bonds to satisfy the octet rule, which specifies that atoms of main-group elements tend to combine in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas, the next step in understanding reactivity is to calculate formal charge.

Valence electrons do not belong to any one atom in a molecule or ion. Quantum mechanics tells us that electrons are shared by a few neighboring atoms, or even by the whole molecule. Atoms differ in electronegativity and hybridization, so it is inaccurate to assume that these electrons are shared equally. A formal charge is a comparison of electrons owned by an atom in a Lewis structure compared to the number of electrons possessed by the same atom in its unbound, free atomic state. Formal charge provides some indication of electron distribution within a molecule or ion, providing a starting point to predict chemical and physical properties.

The procedure to determine formal charges on the atoms of an ion or molecule has three steps. The process is illustrated using hydronium ion (H3O+ ); an ion very frequently encountered in organic and biochemical reaction mechanisms.

Step 1: Draw a Lewis structure for the molecule, including all unpaired electrons. Be sure to show all nonbonded electrons, as these influence formal charges.

Step 2: Assign the formal charge to each atom. Formal charge is calculated using this formula: FC = GN – UE – 1/2 BE

Where:

FC = formal charge

GN = periodic table group number (number of valence electrons in free, nonbonded atom)

UE = number of unshared electrons

BE = number of electrons shared in covalent bonds.

Step 3: The sum of the formal charges of all atoms must equal the overall charge on the structure. The best Lewis structure or resonance contributing structure has the least number of atoms with formal charge. Equivalent atoms have the same formal charge.