Bonding

    Chemical bonds form to lower the energy of the system, the components of the system become more stable through the formation of bonds.   If we examine the periodic table, we find that the elements in Group VIII (or 18), helium, neon, argon and so on, are particularly stable, so much so that they were once labeled the "inert gases". We now know that these elements are not inert, indeed xenon forms a range of compounds, but, nevertheless, they are very stable (although now we refer to these elements as the noble gases). This stability is the result of their electronic configuration, they have a full valence shell of electrons (s², p⁶) and this imparts stability. G. N. Lewis (1916) suggested that bonds (covalent) formed to enable elements to attain this "noble gas configuration". While some of Lewis' predictions have since been proven incorrect (he suggested that electrons occupy cube-shaped orbitals), his work established the basis of what is known today about chemical bonding.  We now know that there are two main types of chemical bonding, ionic bonding and covalent bonding. 

    In ionic bonding, electrons are completely transferred from one atom to another.  In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions.  The oppositely charged ions are attracted to each other by electrostatic forces which are the basis of the ionic bond.  For example, during the reaction of sodium with chlorine, sodium loses its one valence electron to chlorine 

resulting in a positively charged sodium ion and a negatively charged chlorine ion.

    Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron.  This is typical of the relative sizes of ions to atoms, positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent.  After the reaction takes place, the charged Na⁺ and Cl^-  ions are held together by electrostatic forces, thus forming an ionic bond.  Notice also that sodium now has 10 electrons similar to neon while chorine now has 36 electrons similar to Argon.  Hence both atoms have attained the electronic configuration of a noble gas when they became ions.  As a result a higher level of stability has been achieved. 

    Ionic compounds share many features in common:

• Ionic bonds form between metals and non-metals, 
• In naming simple ionic compounds, the metal is always first, the non-metal second (i.e.. sodium chloride), 
• Ionic compounds dissolve easily in water and other polar solvents, 
• In solution, ionic compounds easily conduct electricity, 
• Ionic compounds tend to form crystalline solids with high melting temperatures. 

    This last feature, the fact that ionic compounds are solids, results from the intermolecular forces (forces between molecules) in ionic solids.  If we consider a solid crystal of sodium chloride, the solid is made up of many positively charged sodium ions (pictured at right as small gray spheres) and an equal number of negatively charged chlorine ions (green spheres).  Due to the interaction of the charged ions, the sodium and chlorine ions are arranged in an alternating fashion as demonstrated in the schematic at right.  Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction.  The concept of a single molecule becomes blurred in ionic crystals as the solid exists as one, continuous system.  Forces between molecules are comparable to the forces within the molecule, and ionic compounds tend to form crystal solids with high melting points as a result.

    The second major type of atomic bonding occurs when atoms share electrons.  As opposed to ionic bonding in which a complete transfer of electrons occurs, covalent bonding occurs when two (or more) elements share electrons.  Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons).  This most commonly occurs when two non-metals bond together.  Because both of the non-metals will tend to gain electrons, the elements involved will share electrons in an effort to fill their valence shells.  A good example of a covalent bond is that which occurs between two hydrogen atoms.  Atoms of hydrogen (H) have one valence electron in their first electron shell.  Since the capacity of this shell is two electrons, each hydrogen atom will 'want' to pick up a second electron.  In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H₂.  Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each others single electron, forming one covalent bond.  In this way, both atoms share the stability of a full valence shell. 

    Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed.  Thus covalent molecules are not  strongly attracted to one another.  As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature. 

    For every pair of electrons shared between two atoms, a single covalent bond is formed.  Some atoms can share multiple pairs of electrons, forming multiple covalent bonds.  For example, oxygen (which has six valence electrons) needs two electrons to complete its valence shell.  When two oxygen atoms form the compound O₂, they share two pairs of electrons, forming two covalent bonds. 

Resonance

    Sometimes, a single Lewis structure does not adequately represent the true structure of a molecule.  Consider the carbonate ion, CO₃^2- .  Carbon (C) has four valence electrons and oxygen (O) has six valence electrons.  Carbon is the central atom, the three oxygens are bound to it and electrons are added to fulfill the octets of the outer atoms. 

    All the available electrons have been used but carbon is electron deficient - it only has six electrons around it. So, we share a non-bonding electron pair on an oxygen with the carbon to create a double bond and thereby fulfill carbon's octet. 

becomes 

    But two other possible configurations are possible:

becomes 

becomes 

    We can sum this up as:

    From a quantum perspective, it is wrong to think that these forms are oscillating back and forth when in reality they all exist equally at the same time.  This may be be challenging to think about since you are used to thinking of electrons as discrete points which due to a rather Newtonian way of thinking of the world. 

    These structures are created as a result of attempting to form the most stable arrangement of electrons.  Clearly the task will become harder if more than eight electrons are involved in bonding.

Orbitals

    Since electrons are constantly in motion and do not represent discrete points, how can we properly visualize the bonding of electrons.  We need to appreciate that electron position can only be predicted based on probability theory.  S orbitals only contain two electrons, each traveling in different directions from each other.  As a result there is little electron-electron interaction. As a result the s orbital forms a perfect sphere as there is an equal probability of finding an electron in every direction around the nucleus of the atom.

    However, it is much easier to draw this as:

    The p orbital is more complex since p can hold six electrons, but each orbital can only hold 2 electrons.  Thus there are three equal orbitals making up p.  They are equidistant from each other, but there interaction creates a unique shape around the nucleus.  They are bilobed and can be visualized spread out along three axes. 

    Orbitals aid in our understanding of why some atoms bond the way they do.  For example carbon only has two free bonding electrons (one in each p orbital), yet carbon forms four bonds to achieve a higher stability.  How is this actually accomplished.  Carbon promotes a 2s electron to a vacant p orbital and merges it s orbital with its p orbitals creating a sp³ orbital that will accept four electrons for bonding.  The sp³ orbitals have slightly higher energy than the 2s orbital and slightly lower energy than the p orbital.  This loss of energy though is more than made up by the higher stability that can now be achieved via bonding. 

    Note that each sp3 orbital is not bilobed.  Other atoms are capable of hybridizing their orbitals to create a variety of unique orbitals that will allow increased stability through bonding.  This must be considered when deciding how atoms will bond.  If higher stability can be achieved, hybridization will likely occur.

    Bonding occurs when orbitals from two atoms overlap to the point that a stable electron arrangement occurs.  If the bond is a single bond between two of the same or different orbitals, it is called a sigma bond.  It occurs from the complete linear overlap of these orbitals.  Consider two s orbitals overlapping to create a sigma bond.

    Consider two p orbitals forming a sigma bond. 

    If we consider CH₄, we see that the sp³ orbitals sigma bond to hydrogen atoms.  In this diagram the overalpping shared electrons are shown.  Notice that their spins are opposite to each other.

    If molecules contain double or triple bonds, the first bond formed will be a sigma bond, but this then restricts the orientation of further bonding between orbitals.  This results in a sideways overlapping of orbitals creating what is called the pi bond.  Due to this stretched overlapping, the individual pi bond is always weaker than the sigma bond.

    Consider the triple bond in the following molecule.

Molecular Shape

  We know, then, what makes atoms stick together, but what limits the proportions in which they can combine and what factors control molecular shape?  The simplest and most basic factor arises from the relative sizes of the different atoms. The concept of atomic size is, however, somewhat fuzzy: the size of an atom is controlled by the extent of its electron charge cloud, and the density of electrons around an atom does not suddenly reduce to zero; it gradually decays.  Nevertheless it has proved possible to assign approximate radii to atoms.  Furthermore, purely geometrical factors will limit the ways in which we can pack atoms together. These factors become most obvious when we consider packing in crystals; thus in crystal structures containing, for example, the large cesium ion, there are often eight other atoms surrounding each cesium; whereas the smaller lithium ion has room for only four. 

    The next point is that atoms have well defined combining powers (the chemical concept of valence) which arises from specific features of their electronic structure; in particular the number of electrons in the outermost shell relative to the total number of electrons that can be present in that shell.  Indeed, the concept of valence springs from one of the oldest and most powerful ideas of theoretical chemistry —the 'electron pair' bond. Chemical bonds (like those between the two hydrogen atoms in H₂) often have a pair of associated electrons.  A key feature of chemical bonding to which we now turn concerns the fact that it can have specific directional requirements in order to minimize electron-electron repulsion within the molecules.  This requires that atoms rearrange themselves three dimensionally.   These specific geometrical requirements follow naturally from the criteria that the resulting assembly of atoms should have the lowest possible energy. More specifically, they can be understood in terms of the different shapes of the atomic electron density charge clouds and by the ways in which these can interact.  Chemical bonding occurs where the atomic charge clouds interact and overlap with each other.   Therefore to determine the shape of a covalent molecule, consider the number of electron pairs situated around the central atom of the molecules and resolve the shape ot minimize the electron-electron repulsion.  However, if a lone electron pair is present, it will influence the orientation of the other electron pairs but will not be realized in the final identity of the molecule.

    The process repeats with five electron pairs around the central atom (due to hybridization).  Here the main shape will be the trigonal bipyramidal, but as lone pairs of electrons are included the shape will be named differently as is the case for the tetrahedral shape induced by four electron pairs in the chart above.  As each lone pair is added, the name changes as the shape is not fully visualized.

Molecular Polarity

    There are, in fact, two sub-types of covalent bonds.  The H₂ molecule is a good example of the first type of covalent bond, the non-polar bond.  Because both atoms in the H₂ molecule have an equal attraction for electrons, the bonding electrons are equally shared by the two atoms, and a non-polar covalent bond is formed.  Whenever two atoms of the same element bond together, a non-polar bond is formed.  A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion).  In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.

    The primary difference between the H-O bond in water and the H-H bond is the degree of electron sharing. The large oxygen atom has a stronger affinity for electrons than the small hydrogen atoms. Because oxygen has a stronger pull on the bonding electrons, it preoccupies their time, and this leads to unequal sharing and the formation of a polar covalent bond. 

    The ability of atoms to attract shared electrons is termed electronegativity.  Electronegativity values can be determined by comparing atoms and this can be used to determine the type of bond that would result.  Trends do exist and for the most part they can be used to predict bond types.  Bond types actually exist as a continuum from non-polar to ionic depending upon the relative strength of atoms involved.  An electronegativity difference of 0 indicates perfectly equal sharing, but a value of 0-0.99 means unequal sharing.  However, this unequal sharing is not unequal enough to cause any significant polarity effects.  These bonds are still classified as nonpolar.  If the electronegativity difference between the two atoms exceeds 1.0, the unequal sharing is significant enough to be detected and the bond formed is polar.  If the unequal sharing results in an electronegativity difference of greater than 1.7, the sharing is so unequal that one atom rips an electron from the other atom creating an ionic bonding.  In fact, one would hardly call this sharing at all.  The ability to rip electrons away to create such a clear unequal sharing of electrons can be determined by looking at % ionic character of a bond.  An electronegativity of 0 has a 0% ionic character and the % increases from there until it is high enough to allow ionic bonds to form.

    Water molecules contain two hydrogen atoms bonded to one oxygen atom (blue).  Oxygen, with 6 valence electrons, needs two additional electrons to complete its valence shell.  Each hydrogen contains one electron.  Thus oxygen shares the electrons from two hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing the H valence shells. This means that water arranges its electron pairs to create a tetrahedron around the oxygen atoms, but since there are only two bonds present, the shape is bent.  Since the OH bond is polar, water exists with a discrete partial negative region and a partial positive region within the neutral molecule.  This is due to the unequal distribution of electrons.

    Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen end of the molecule develops a partial negative charge (because of the negative charge on the electrons).  For the same reason, the hydrogen end of the molecule develops a partial positive charge.  Ions are not formed, however the molecule develops a partial electrical charge across it called a dipole.  This allows water to orientate itself in different manners depending up whether a positive or a negative charge is nearby.  A nonpolar molecule may have polar bonds but it will always orient itself the same way in response to nearby charges. 

Solids

    There are two types of solids. They are crystalline and amorphous. They have differences that are very important in the study of Physical Pharmacy. They differ in the way they are arranged, their melting points, and the way they should be treated when used in making drugs.

    Crystalline solids are arranged in fixed geometric patterns or lattices. Examples of crystalline solids are ice, methanol, and sodium chloride. They have an orderly arranged units and are practically incompressible. Crystalline solids also show a definite melting point and so they pass rather sharply from solid to liquid state. There are various crystalline forms which are divided into six crystal systems or shapes. They are cubic, tetragonal, hexagonal, rhombic, monoclinic, and triclinic. The units that constitute these systems can be atoms, molecules, or ions. Ionic and atomic crystals are hard and brittle with high melting points. Molecular crystals are soft and have low melting points. Metallic crystals are composed of positively charged ions in a field of electron gas or freely moving electrons. Metals are good conductors of electricity because of the free movement of electrons in the lattice.

    Amorphous solids are solids with random unoriented molecules. Examples of amorphous solids are glass and plastic. They are considered supercooled liquids in which the molecules are arranged in a random manner some what as in the liquid state. Amorphous solids also unlike crystalline solids do not have definite melting points.

    The difference between an amorphous and crystalline solid is very important in drug making. When making a drug in solution, the drug is added to the other chemicals to prolong the shelf life. When the drug is crystallizing, if it forms a crystalline solid, there is space in the crystal for the ice to come out leaving the drug and the components. This process only takes about two or three days. If the drug forms an amorphous solid during the crystallizing phase then it takes about seven days. This is because amorphous solids do not have space for the ice to come out during the freezing therefore the ice must diffuse out. Therefore it is preferable to have crystalline solids in drug making.

Ionic Solids

    A solid that forms due to the Coulomb attraction between oppositely charged ions. For example Sodium choride consist of a close 3-dimensional array of alternating positive and negative ions having a lower energy than the separated ions.  Sodium will gives up its 3s valence electron to from Na⁺.   Chlorine is more stable by closing the shell making a Cl^-. Spherically symmetric closed shell ions are involved so the ions are arranged like closed packed spheres.  The structure is stable because the binding energy due to the net electrostatic attraction exceeds the energy spent in transferring electron to create the isolated ions from neutral ions.  The crystal structure is the one that mininimizes the energy which depends on the ions involved.  Because no free electrons can carry energy or charge they are poor conductors of heat or electricity in the solid sate.  Because of the strong electrostatic forces between ions they are hard and have high melting points.  Lattice vibrations can be excited by energies corresponding to the far infrared so they show optical absorption in that region as well as in ultraviolet; but they are transparent to visible radiation.

Molecular Solids

    They consists of molecules which are so stable that they retain much of their individuality when brought into close proximity.  Electrons in the molecules are paired so that atoms in different molecules cannot form covalent bonds with one another.  The intermolecular binding force is the weak Vand der Waals attraction; an interaction between electric dipoles.  There is also a momentary dipole created by shifting electrons within atoms.  This induced dipole is called the London Dispersion Force.  The greater the probability of shifting electrons within atoms, the stronger and more frequent the LDF.  Many organic compounds, inert gases and ordinary gases such as oxygen, nitrogen, and hydrogen form molecular solids in the solid state.  Because the binding is weak solidification takes place only at very low temperatures where the disruptive effects of thermal agitation are very small.  The absence of free electrons makes them very poor conductors of heat and electricity.

Network Solids

    In this case, electrons are shared by atoms.  Atoms are bound by shared valence electrons.  Bonds are directional and determine the geometric arrangements of atoms in the crystal structure.  The rigidity of their electron structure makes covalent solids hard and difficult to deform so high melting point.  There are no free electrons so there is no good heat or electrical conduction.  Sometimes, as for silicon and germanium they are semiconductors.  At room temperature some covalent solids like diamond are transparent as the energy to excite their electronic states exceeds that of photons in the visible region so the photons are not absorbed.  But most covalent solids absorb in the visible and are therefore opaque.  In diamond each carbon atom is bonded covalently to 4 other carbon atoms.  In diamond the basic structure is tetrahedral.  Graphite, an allotrope (different arrangement of the same element) of diamond, is softer, conducts electricity and is flexible becasue graphite has one delocalized electron per every 3 covalent intermolecular bond.  Delocalized electrons in network solids causes their properties to vary.

Metallic Solids

    These solids exhibit a binding that can be thought of as a limiting case of covalent binding in which electrons are shared by all ions in the crystal.  This is sometimes called pseudocovalent binding.  When a crystal is formed of atoms having a few weakly bound electrons in the outermost subshells, electrons can be released for individual atoms by energy released in binding.  The atoms have vacancies in their outermost subshells and there are not enough valence electrons per atom to form tight covalent bonds.  These electrons move in the combined potential of all the positive ions and are shared by all atoms in the crystal and thus are free to wander through the crystal from atom to atom there being many unoccupied electron states.   In this sense they behave like a gas in that an electron gas is interspersed between the positive ions and excerting forces on each ion that exceeds the repulsive force of other ions; hence resulting in binding.  A metallic solid is a regular lattice of spherically symmetric positive ions arranged like close-packed spheres through which electrons move.  Metallic solids are excellent conductors of electricity or heat. The electrons easily absorb energy from incident radiation or lattice vibrations and move under the influence of an applied electric field or thermal gradient.  Because radiation in visible wavelengths easily absorbed they are opaque. All the alkalies form metallic solids.  The type of binding is determined experimentally by studies of x-ray diffraction, dielectric properties and optical emissions.