Covalent Bonding!

Before I even begin talking about what Covalent bonds are I want to clear up some possible confusion. If you’re learning general chemistry now you’ll thank me later.

The last two graphics above are typically what are used to illustrate what a covalent bond is between atoms. However this picture is not quite accurate. In reality atoms don’t align themselves in a Bohr model form of atoms and electron orbiting around like planets and moons.

Reality it turns out is just a bit weirder than that. Because quantum mechanics is really weird. The reality is that what exists are orbitals, which are a probability distribution of where an electron may be around an atom. These take on particular shapes depending on the element (heavier elements have more and different orbitals). This is because according to Heisenberg’s Uncertainty Principle, it is not possible to know both where an electron is and its exact momentum. Because of this we can only rely on probabilities on where electrons may be.

The first image is a closer representation of the way electrons are distributed than the circular models used in Basic/General Chemistry.

But this isn’t about orbitals. This is about Covalent bonds, so lets begin.

Covalent bonds are essentially what takes place when the valence electrons (the outermost electrons in an atom) are shared between atoms forming a bond. These occur between non-metals. They can form single, double or triple bonds depending on the amount of pairs of electrons that are shared. In more detail, these bonds consist of sigma and pi bonds, but I won’t go into those today.

These covalent bonds lead to stable molecules if they share electrons in such a way to create a noble gas configuration (where the electrons amounts are the same as those for the noble gas elements) for each atom. Essentially, when the ns2 and np6 orbitals are filled, this leads to stable orientation for the molecules. This is what is known as the Octet Rule, as it involves filling 8 electrons up in the mentioned ns2 and np6 orbitals.

There are different types of Covalent Bonds too.

Nonpolar covalent bonds occur between identical, non-metal atoms. These are the simplest types of covalent bonds. Because the atoms are the same, the electrons are shared equally between the atoms.

Examples of Nonpolar covalent bonds are: Oxygen (O2), Iodine (I2) and Fluorine (F2).

Polar Covalent Bonds are covalent bonds in which the shared electrons are not shared equally. Different atoms pull differently on the shared electrons, which will lead to an increased electron density to one side of the molecule rather than the other. The atom that pulls more strongly on the electrons becomes slightly more negative, and the other slightly more positive. This positive and negative end of a molecule is known as a dipole (like a magnet)  How much this occurs is related in the numerical value of Electronegativity. Electronegativity in simpler terms is an atoms affinity for attracting electrons to itself.

Examples of Polar Covalent bonds are: Water (H2O), Ammonia (NH3) and Hydrogen Fluoride (HF).

Whether a molecule is polar or nonpolar covalent can be determined by the electronegativity difference of the atoms that compose it. From 0.0-0.2 is nonpolar covalent. From 0.3-1.4 is Polar Covalent. From >1.5 is Ionic (but I won’t talk about that now). Covalent bonds are stronger than Hydrogen bonds, but are in general weaker than Ionic bonds.

If I’ve made any errors please let me know and I’ll correct them ASAP!

Sources: [1][2][3]

Other articles by me:

Hydrogen Bonding

Whales are Really, Really Fascinating


Note: As written, this activity requires that students hold hands. Younger students may not have any problems with this, however, the self-consciousness of adolescents may hinder the spontaneous movement and physical contact required for this activity. If you think this will be problematic in your classroom, cut 12-inch lengths of string for the students to hold to make the ‘bonds.’


Ionic Bonds

Ionic Bonds are a type of chemical bond that typically forms between metal and non-metal atoms. I say typical because it is also possible between certain non-metal molecular ions such as Ammonia (NH3). Unlike Covalent bonds, the electrons in Ionic bonds are not shared equally. They are instead transferred between species. Ionic bonds are formed by the electrostatic attraction between oppositely charged ions.

The images above show beautifully, that when a metallic species gives up its electron(s) it becomes more positively charged (what we call a cation), and the atom receiving the extra electron(s) becomes negative (what we call an anion). It is these positive and negative charges on these ions that form ionic bonds. For the most part these compounds follow the octet rule of filling electron shells to became stable, but this is a general rule, not absolute. There are many exceptions (particularly with metals).

Properties of Ionic bonded compounds are that they conduct electricity while molten or in a solution, they generally have a high melting point and tend to be soluble in water. In the real world, ionic compounds exist not as pairs of bonded atoms, but in arrangements of ions set in very specific shapes and forms. What we call a crystal lattice. This specific form repeated millions of times gives ionic compounds their 3 dimensional shape (salt crystals for example). Ionic bonds are the strongest of all bonds (Ionic > Covalent > Hydrogen).

Pure ionic bonds do not exist in nature. They are all covalent to some degree. The greater the electronegativity difference between atoms, the more ionic the bond. For reference again:

>1.7 EN difference = Ionic bond

0.4-1.7 EN difference = Polar Covalent

< 0.4 EN difference = Covalent

The example of Sodium Chloride (NaCl), or table salt is a good one to explain ionic bonds. When Na and Cl are combined, the sodium atoms lose an electron becoming cations (Na+), while the Chlorine atoms each gain an electron to form anions (Cl-). These oppositely charged ions are then attracted to each other in a 1:1 ratio to form Sodium Chloride (NaCl). Depending on the charges on the atoms, this ratio can vary. The compound Potassium Oxide (K2O) for example, the potassium exists in a 2:1 relationship with Oxygen because Oxygen’s charge is -2; so two K atoms (charge of +1) are needed per O atom to balance the charges. The greater the charge, the greater the attraction, and the stronger the ionic bond.

To remove electrons from atoms in this process requires the input of energy or heat, and is therefore endothermic. However the formation of ionic compounds is exothermic and creates heat or energy as a byproduct. Ionic bonding however will only occur if the reaction is what we call favourable. Essentially the bonded atoms must have a lower energy requirement than the free atoms. The larger the energy change, the stronger the ionic bond. Because metals have low electronegativity, and non-metals have high electronegativity, the resulting energy change of the reaction is more favourable when metals lose electrons and non-metals gain them.

This can be explained easier by reviewing the 3 steps in ionic bonding, in this case again with NaCl:

Please note that this is a simplification that does not take into effect a solid product forming into a crystal lattice (as occurs in reality) or beginning with gaseous Cl (Cl2). See here for a more detailed picture.

Ionic bonding can be broken down into 3 steps: ex. sodium chloride

Formation of the sodium ions (ionization)

Na(g) + energy —–> Na+(g) (DH = +496 kJ/mol)

Requires 496kj/mol input to remove electrons.

Formation of the chloride ions (electron affinity)

Cl(g) + e- —–> Cl-(g) + energy (DH = -348 kJ/mol)

Releases 348kJ/mol of heat/energy in this reaction.

Formation of the ion pair (bond energy) 

Na+(g) + Cl-(g) —–> NaCl(g) + energy (DH = -504 kJ/mol)

Releases 504kJ/mol of heat/energy in the final formation of NaCl.

DH = ΔH = Change in Enthalpy (heat energy).

We use Hess’s Law, which simply put means that the total energy change in a reaction is equal to the sum of the energy changes that occur.

So adding up each reaction we get:

DHrxn = DH1+ DH2 + DH3 = -356 kJ/mol

Sources: [1][2][3][4][5]

Read my other science writing posts here.

First Hi-Res Images Taken of a Molecule Breaking and Forming Chemical Bonds

Felix R. Fischer of the U.S. Department of Energy’s Lawrence Berkeley National Laboratory used nc-AFM techniques to yield direct imaging of covalent bond structure in single-molecule chemical reactions using .

Fischer originally set out with the goal of finding a better way to mass produce graphene nanostructures for use in transistors, logic gates, and other pieces in electronic equipment. To create these exact bonds from the bottom up, the laboratory at UC Berkeley needed a clear view under controlled reactions in order to make sure they were getting it right.

Enter a clever imaging technique called “noncontact atomic force microscopy,” which basically works the same way that a phonograph does: Scratching or probing a surface with a sharp tip in order to read it. Unlike a phonograph, though, the AFM ‘needle’ — really a single oxygen atom — is deflected by small electronic forces that create a readable pattern. This type of microscopy ended up being so precise that it was able to detect not only the atoms, but the actual forces that create the bonds between them. (more)

Liquids versus Solids

     I was curious as to the actual differences internally at the molecular level between solids and liquids beyond simply their temperatures in determining their precise state of matter.  So I did a quick study.  First, molecular bonds are dependent on how electrons interact between the adjacent atoms that they’re composed of, with electrons thereby producing the forces involved in forming molecules.  So what forces, then, are left to explain the differences between these two states of matter?  I tried to disregard gases in this study because they’re easy to understand, they’re simply behaving like a complete loss of attraction between molecules has occurred, this thereby resulting in complete freedom and expansion of the substance as a gas.  And solids aren’t too hard to understand either, being that with enough lowering of temperature, the molecules then collapse into rigid relationships.  But liquids are weird!  I mean, why stay both somehow bound together while at the same time being able to freely move all around?

     So this is the answer that I discovered.  Although molecules are indeed held together in bonds between atoms due to how electrons overall behave and interact, these electrons also have a similar but lower attractive effect between different molecules.  The electron forces involved inside a single molecule, however, are generally much stronger than how they are in-between differing molecules.  Nevertheless, it’s still all centered on electron forces regardless of whether in forming a molecule or attracting different molecules to one another.

     What’s actually happening involving the weaker electron attractions between different molecules is that each molecule has a physical shape, a specific construction.  And based on the specific construction of a given molecule, each different region of each molecule will carry a different level of electrostatic voltage because of how the electrons are uniquely involved within the various bonds between its atoms.  And it is because of these resulting differences in electrostatic voltages around each molecule that they in turn then attract other molecules with these differing electrostatic forces that are position-dependent on their exteriors!

     So, when the temperature is low enough so that the molecules move slowly enough for these secondary electrostatic attractions to come into play, then the substance forms either a solid or a liquid rather than a gas due to its molecules attracting each other to the point of actually being able to come into contact with each other and stay connected as such.  And the precise difference between whether it behaves as a solid or a liquid in this condition is then further determined by how energetic that these molecules are that are now touching each other due to these weaker electrostatic attractions that have come into play between the different molecules.  Once the molecules slow down too much, then they form a solid because they find their best positions between adjacent molecules, such as two magnets attaching to each other, and then they just stay there because of there being too little energy to bump them out of place.

     But if there’s still too much energy involved in their movements due to there being too high a temperature, then even though they’re currently acting upon this weaker attraction between molecules that presses them next to each other, they’re nevertheless still too fidgety to stay settled down into any particular spots of connectedness between the now-joined molecules.  So they simply keep shuffling around each other all over the place…kind of like country dancers who are always holding onto other people’s hands yet continually swapping the specific person’s hands that they’re holding onto.  And this is what then amounts to they’re behaving like freely flowing liquid at the time.  And even while in this freely-moving liquid state of joining between molecules, the individual molecules still maintain their original formation as well because of the bonds between the atoms within each of the molecules having a much stronger electrostatic force that holds them together than is exhibited on the exteriors of the molecules involving the formation of the state of matter for the substance at the time.

     So this is what I’ve come to understand as to the specific difference in terms of what is internally going on at the molecular level between solids and liquids, and how electrons are therefore the same force that’s behind both the specific state of matter as well as the formation of its molecules for any given compound or substance.  I know that this explanation tends to generalize a good deal when it comes to all of the many details involved with all of these things such as the different types of chemical bonds that are involved, but they all still involve the interactions of electrons, and therefore the specific varieties of these situations aren’t really what I considered as being very important to the overall understanding of this topic.

     — Mark Ober, 08/30/2015