In Atoms Part 4B, we learned that covalent bonding is favoured over ionic bonding when atoms are of similar electronegativity. Electrons are shared within electron pairs, forming a stable balance between attractive and repulsive forces. Electrons from both atoms are attracted to the two positively charged nuclei, while the electrons themselves, being negatively charged, experience a repulsive force between each other. Electrons push each other away like two south poles of a magnet. The reason they pair up is so that the most electrons possible can fit into the lowest possible energy orbitals, and by pairing up, electrons can share an orbital. This lowers the overall energy of the atom so it is favoured.
Covalent bonding doesn't necessarily mean, however, that the electron clouds involved are perfectly evenly dispersed across the molecule, a phenomenon called electron delocalization. When two or more atoms with different electronegativity bond covalently, the electrons may be shared unequally between them, creating areas with slightly different charges. These bonds are called polar covalent bonds. If the polarity is extreme, the bond created is an ionic bond. Polar covalent bonds, therefore, fall midway in the spectrum between non-polar covalent bonds and ionic bonds.
Water Has Strongly Polar Covalent Bonds
Polar bonds create tiny accumulations of charge. Atoms with high electronegativity, such as oxygen (3.44), pull on electrons strongly, so when oxygen bonds covalently with other atoms, such as hydrogen (2.20), it often forms strongly polar molecules. A water molecule is an example, shown below left as a Lewis dot diagram:
Notice the two lone pairs of electrons around the oxygen atom. The molecule appears bent because the two positively charged hydrogen atoms are flexed away from the negatively charged filled oxygen orbitals. Negative charge accumulates around the central oxygen atom (red), leaving the outer parts of the two hydrogen atoms with a net positive charge (blue) in the diagram below:
Polar molecules such as sugar, and ionic solids, such as sodium chloride, dissolve easily in (polar) water. Oils and waxes are nonpolar molecules, so they don't dissolve in water but they will dissolve in very nonpolar solvents like cyclohexane, which is a ring of six carbon atoms covalently bonded together. Charge is distributed very evenly in this molecule. "Like dissolves in like" is the common rule of thumb here. Cyclohexane, not surprisingly, won't dissolve in water. Instead it will form a separate distinct layer, just like oil does.
Water is a very fascinating polar molecule. I started off the "Our Universe" articles by exploring water. At very high temperatures and pressures, inside Jupiter for example, water transforms into a fully ionic state - a soup of positive hydrogen ions and negative oxygen ions. At even higher pressures, water will become a superionic solid. Oxygen crystalizes out of the soup and hydrogen ions float freely within the oxygen lattice.
One Molecule - Two Kinds of Bonds: Polyatomic Ions
In Atoms Part 4A, I mentioned phosphorus [P] as an example of a half-filled valence orbital. This makes phosphorus reactive - it wants to fill that orbital and find a more stable, lower-energy, argon-like electron configuration. One way it can do this is by reacting with four oxygen [O] atoms to form a complex ion, PO43-. This ion is called phosphate. It is one of many different kinds of polyatomic ions, ions made of more than one kind of atom. The electronegativity of phosphorus (2.29) and oxygen (3.44) are close enough that they form covalent bonds, but the values are far enough apart that the covalent bonds are polar in nature. Negative charge accumulates around the oxygen atoms. A phosphate ion drawn as a Lewis structure is shown below left:
Lewis structures are often simplified using lines for bonds instead of drawing electron "dots." As we saw in the last article, a single line means a single bond (two dots). A double line means a double bond (four dots). This polyatomic ion (left) acts as a single unit that has an overall charge of -3. Like all covalently bonded molecules, phosphorus can lower its energy by sharing its five valence electrons with oxygen atoms. Do you notice something odd about the phosphorus atom here? It is breaking the octet rule. It has more electrons around itself than you'd expect - ten, rather than eight. Elements in the third period and below can accommodate more bonds than the rule allows, an exception to the octet rule.
Oxygen atoms (3.44) are more strongly electrostatic than phosphorus (2.19) (see the electrostatic periodic table far below), so each oxygen atom tries to draw two electrons to itself. If you look at the periodic table (below) you'll see that oxygen is two electrons short (eight) of the noble gas atom, neon (ten).
One oxygen atom uses two of its six valence electrons to form a double bond. The other four valence electrons of this oxygen form two lone pairs. The other three oxygen atoms can't do the same thing. Phosphorus only has five valence electrons available for bonding, so these oxygen atoms each attract one phosphorus electron into a single bond. In doing so, they are one electron short of an octet.
They fix this unstable situation by attracting an extra electron (which can be from water or another atom). In doing so, they each attain a single negative charge. It is energetically favourable for any atom to pair up its electrons. The three negative charges mean that the phosphate ion has a charge of -3.
In reality, phosphate doesn't arbitrarily choose one oxygen atom to double bond with and then stick with this particular atom, creating one shorter bond and three longer bonds. These bonds display resonance instead. All three bonds are the same length and strength - a hybrid between a double bond and a single bond.
The formation of a polyatomic ion with phosphorus is energetically favourable to oxygen atoms but it leaves three oxygen atoms one electron short of a stable full valence shell, a better but still energetically unfavourable situation. The complex ion has achieved its goal of saving energy by electron sharing but it could improve things even more. Three of the four oxygen atoms are ready to react . . .
A Brief Introduction To Acids And Bases
Many atoms will react with this polyatomic ion to form an ionic compound. Hydrogen [H] found in dissociated form in water is one example. Three hydrogen atoms will bond with phosphate's three oxygen atoms to create phosphoric acid, H3PO4. Phosphoric acid is an ionic compound containing a highly polar complex ion - phosphate. It is a crystalline powder at room temperature that dissolves easily in water, and when it does, just like table salt, it dissociates into two ions:
H3PO4 ↔ H+ + H2PO4-
H2PO4- ↔ H+ + HPO42-
HPO42- ↔ H+ + PO43-
Notice that I've drawn three reactions instead of just one. Phosphoric acid doesn't just dissociate once in water, it dissociates three times, each time releasing an H+ ion. H+ ions react with water to create H3O+ ions. The concentration of H3O+ ions, called hydronium ions, determines the pH of a solution. As the concentration of H3O+ increases, the phosphate solution becomes more and more acidic, until eventually equilibrium is reached. This is basically how an acid works.
If phosphate is added to a solution that is very acidic, around pH 2, almost all phosphate ions will be in the PO43- form. If phosphate is added to pure water, around pH 7, you will get a mixture of phosphate ions, most of which will be HPO42-. If you add phosphate to a very basic solution (pH 12) you will get mostly H2PO4- ions.
All hydrogen phosphate (ionic) compounds dissolve readily in water but the phosphate complex ion, being covalently bonded, does not. It stays intact.
Bonding Determines the Chemical and Physical Properties of Molecules
The strong electrostatic bonds in ionic compounds give these compounds typical chemical and physical properties. As we've seen with sodium chloride and phosphoric acid, ionic compounds tend to dissolve easily in water. They also conduct electricity very well in solution with water, as well as when they are in a molten state, but not in their solid state. Ionic compounds have very high melting and boiling points. They are almost always found as solids at room temperature, usually as tightly packed lattices or crystals. This tight packing means charged ions or electrons can't move easily and that is why they don't conduct electricity as solids.
The physical and chemical properties of covalent compounds tend to vary much more than they do for ionic compounds. These compounds may be solids, liquid or gases at room temperature. They are not usually electrically conductive. Polar covalent compounds are soluble in water but nonpolar ones are not. Examples of covalent compounds are water, carbon dioxide gas, and diamond.
This nine-minute video by Montana science teacher, Paul Andersen, reviews the covalent and ionic chemical bonds we've explored in a very user-friendly way:
METALLIC BONDING: A SPECIAL KIND OF COVALENT BOND
Metals make up a large part of the periodic table. You are probably familiar with metals like chromium [Cr], nickel [Ni], silver [Ag] and aluminum [Al], for example, but many other elements are also technically metals. In the periodic table below, every element left of the black line as well as the lanthanides and actinides are metals. This does not mean, however, that all of these elements form metallic bonds (I'll describe them in a moment) under normal conditions. Hydrogen, for example, is a gas (H2) under normal pressure. It forms polar covalent bonds.
Metal reactivity tends to decrease as you move right on the periodic table shown below. Tin [Sn] is not very reactive but sodium [Na] is.
Metallic bonds, in which atoms of one type of element bond together, tend to be quite strong. That's why most metals have high melting and boiling points. Chromium's melting point is 1907°C, nickel's is 1455°C, silver's is 962°C, and aluminum's is 660°C.
Mercury [Hg], shown below, is a very interesting exception. It is a liquid at room temperature. It freezes into a solid below -38°C (that's its melting point) and it boils at 360°C.
Mercury has a unique electron shell configuration. It's a transition metal, but unlike most transition metals, which have partially filled d orbitals, mercury's electrons fill up all of its orbitals - [Xe] 4f14 5d10 6s2 (we haven't covered the f orbital yet, it can contains a maximum of 14 electrons in it; notice that the d orbital with 10 electrons is full).
Mercury behaves more like a noble gas as a result of it s filled orbitals, so when its atoms bond with each other, they are unusually weak bonds. Mercury's valence electrons are not very reactive so its atoms do not share electrons readily with one another. Metallic bonding occurs but it is weak, so mercury melts far more easily than most other metals.
Metals (except mercury) form tightly packed lattices of atoms, much like ionic solids do, but these solids, unlike ionic solids, conduct electricity very well. Valence electrons are very mobile in these lattices because they are delocalized. This means that their orbitals extend over many adjacent atoms.
Sodium [Na], shown below, is a good example.
(Dnn87 at en.wikipedia)
It has the electron configuration 1s22s22p63s1. When two sodium atoms approach each other, the lone 3s orbital electrons pair up, just like how other covalent bonds form. The difference here is that each sodium atom is surrounded by eight other sodium atoms in a tightly packed lattice. The 3s orbital electron in the central atom overlaps with the 3s orbitals of all the surrounding atoms. Each orbital can only hold two electrons so this means there is a vast number of shared orbitals extending all over the sodium metal. There is no single metallic bond. They are all interconnected, so electrons in these orbitals can move freely all over the metal, within the shared orbitals.
Magnesium [Mg], shown below, has two 3s orbital electrons.
Its delocalized electron "sea" has twice as the electron density of sodium. It has even stronger bonds and a higher melting point (650°C versus sodium's 98°C) as a result. It is also about twice as electrically conductive as sodium.
The delocalized electrons in transition metals have a greater density still, because both their 3d and 4s orbital electrons can take part. These metals are located in the light peachy pink central square in the periodic table above, called the d-block. The melting and boiling points of these metals tend to be very high. My earlier examples of chromium, nickel and silver are transition metals but aluminum is not. Some of these metals are very electrically conductive but the trend based on electron density isn't always simple. Silver is the most electrically conductive metal of all, but iron [Fe] and platinum [Pt], also transition metals, are not very electrically conductive at all.
When metals melt, the metallic bond is still there - the electron orbital sharing still occurs - but the tight lattice structure breaks down and the bonds are weakened.
Under pressure, elements that are technically metals begin to show their metal character. The hydrogen atom I mentioned earlier is a gas in our atmosphere. It shows no metallic bonding at all. The two electrons are completely localized in their orbitals and the two hydrogen atoms bond covalently. But under extreme pressure, again scientists think this might happen deep inside Jupiter, hydrogen is a metal. The hydrogen atoms become so tightly packed together they begin to from a liquid metal and then with increasing pressure they form a solid lattice arrangement, and in this tight arrangement, electrons begin to delocalize. They are unbound to their proton nuclei and are free to move about in shared orbitals, making crystalline hydrogen metal very electrically conductive. (This could be the reason Jupiter has such an enormously powerful magnetosphere.)
Metals are strong, malleable, lustrous and opaque and they conduct electricity, all thanks to metallic bonding.
Science teachers Jonathan Bergmann and Aaron Sams explore metals and how metallic bonding works in this six-minute video:
Next, we'll draw some brief conclusions about atoms, their electrons and their chemistry in Atoms Part 4E.