chemical-bonds

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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

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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)

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Atomic Hook-Ups - Types of Chemical Bonds: Crash Course Chemistry #22 (by crashcourse)