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Published: Thursday, 17 October 2024 01:11
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Atomic Bonds
So far we saw atoms in isolation. Atoms combine together to form molecules which take place via some chemical reaction. The individual atoms form bonds with other atoms, which makes these molecules. These bonds take energy to form, and hence will require energy to break them apart to release individual atoms. Atoms want to have full orbit (or 2 or 8 electrons in their outermost shell), so whatever reaction lets them achieve it, they go for it.
Before we learn Bonds, let's look at electronegativity.
Electronegativity (EN):
Electronegativity refers to the relative ability of an atom to attract it's electrons. An atom which has more +ve charge in nucleus (i.e more protons) will attract it's electrons stronger, meaning it's more electronegative. However, if the atom is bigger (due to higher Z), then electrons will be further out from protons, meaning the force is reduced (as 1/R^2). So, based on this Fluorine is arbitrarily assigned EN=4, which is the highest in periodic table. EN increases as we go from left to right in a row, and decreases as we go from top to bottom in a column. Lowest EN is for Caesium=0.8, with Hydrogen's EN=2.2
Link => https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Electronegativity
So, based on EN trend, elements which are in the diagonal line going from top left to bottom right, have almost same EN< which results in many of their chemical properties being very similar. Since no 2 dif atoms have same EN, if bonds are formed b/w the, the more EN element wil attract the electrons more than the less EN element does.
As it is usually calculated, electronegativity is not a property of an atom alone, but rather a property of an atom in a molecule. Even so, the electronegativity of an atom is strongly correlated with the 1st IE (read in IE in periodic table section).
There are 2 kind of Bonds that exist b/w atoms. One is the bond that exist within a molecule b/w diff atoms of a molecule, and other is the one that exists b/w atoms of different molecules of same compound.
Intramolecular Bonds:
Intramolecular forces are the forces that hold atoms together within a molecule. They are responsible for forming intramolecular bonds.
There are 3 different kinds of intramolecular bonds:
1. Ionic Bond: These are bonds that are formed when one of the atoms loses electrons while other gains them, resulting in +ve and -ve ions, which are attracted to each other and form bonds. Atoms Ex: Na-Cl bond is ionic. Certain ions are referred to in physiology as electrolytes (including sodium, potassium, and calcium). These ions are necessary for nerve impulse conduction, muscle contractions and water balance.
2. Covalent Bond: These are bonds that are formed when the atoms share electrons b/w each other. Here electrons are shared equally b/w the 2 atoms involved to complete their outer shell (require 2 or 8, electrons depending on the outermost shell number, known as the octet rule). These bonds are usually formed b/w atoms of same element, i.e O atoms form covalent bond with each other to form O2. Similarly for H2 . Many bonds commonly found are covalent bonds.ex: H2O. Sharing of 1 electron results in 1 covalent bond, similarly 2 or 3 forms double or triple covalent bond. More the electrons shared, stronger the bond is. Ideally there shouldn't be any charge on any atom as they didn't give or take any charge. 2 types of covalent Bonds exist:
- Non Polar Covalent Bonds: What we know from EN is that, no 2 elements have same EN. When same elements forms bond as in H, the electrons will be right in middle of the bond, and no H atom has more force over these shared electrons. These bonds are called non polar covalent bonds (or pure covalent bonds). Thesebonds are also formed b/w atoms of different elements when their EN is very close to each other.
- Polar Covalent Bonds: Most atoms have different EN. When bonds are between diff atoms as in HCl, then the more EN atom will pull electrons closer to it than less EN one. Here Cl is more EN, so it gets shared electrons closer to itself, resulting in slightly -ve charge on Cl and slightly +ve on H. These kind of covalent bonds are polar, since they have developed a charge separation due to EN. Similar is the case for H2O where H has slight +ve charge and O has slight -ve charge. Non polar Covalent bonds have no charge develop as in O2. If the EN diff is too big, the electrons may get so much closer to more EN atom, that the bond doesn't look like covalent bond, but rather an ionic bond. So, an ionic bond is just a highly polar covalent bond. Thus there is no clear cut separartion b/w Ionic and Covalent bond. They are the same bond, but when charge separation is too large, we refer the covalent bond as ionic bond.
3. Metallic Bond: This is a type of covalent bond that specifically occurs between atoms of metals, in which the valence electrons are free to move through the lattice. This bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions. Metallic bonds are present in samples of pure elemental metals, such as gold or aluminum, or alloys, like brass or bronze.
These are the 3 rules of these bonds:
- No electronegativity difference between two atoms leads to a pure non-polar covalent bond. Ex: CH4 .
- A small electronegativity difference leads to a polar covalent bond. Ex: H2O
- A large electronegativity difference leads to an ionic bond.Ex: NaCl
Metallic bond is the strongest, followed by Ionic, then polar covalent bond, and lastly non polar covalent bond.
Intermolecular Bonds:
Intermolecular forces are forces that exist between molecules. They are responsible for forming intermolecular bonds. They are much weaker than intramolecular bonds. They are important because they determine the physical properties of molecules like their boiling point, melting point, density, and enthalpies of fusion and vaporization.
Intermolecular forces: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/11%3A_Liquids_and_Intermolecular_Forces/11.02%3A_Intermolecular_Forces
Weak Bonds: What we saw above are strong bonds. We can also have weak bonds, which are not chemical bonds, but they are formed b/w +ve side charge of an individual molecule to -ve side charge of different molecule of same compound. These are important as they determine the phase of the compound (solid, liquid, etc), as well as within same phase, they may hold molecules close together to give them desired proporties. These are also referred as Van der Walls forces. 2 kinds of weak bonds:
- Hydrogen Bond: This is a a bond that is usually b/w +ve charged H atom and another -ve charged atom. This bond forms due to EN diff b/w the 2 diff atoms that form the molecule. Water is one of the ex of H bond.
- In general, these are the requirements for a H bond:
- A polar covalent bond needs to exist b/w H atom and a highly EN atom as N(2p3), O(2p4), and F(2p5), These are the only 3 that have large EN difference to form a highly polar bond.
- There should be at least one active lone pair of electrons available in highly EN atom. Lone pairs at the 2-level have electrons contained in a relatively small volume of space, resulting in a high negative charge density.
- Once these 2 conditions are satisfied, a bond starts forming b/w lone pair of electrons in 1 molecule which is highly EN with highly positive hydrogen atoms of another molecule. This bond has 1/10 the strength of an avg covalent bond, so it is strong enough to change properties of that molecule.
- NH3, H2O and HF are 3 examples of compounds which form strong enough H bonds that their boiling points are lot higher, compared to what is expected in the absence of H bond.
- Water is the perfect ex of how H bond causes higher boiling pt. There are 2 pair of lone electrons available on Oxygen to form 2 H bonds with neighboring Hydrogen atoms of another molecule. Even though H2O has covalent bonds and hydrogen and oxygen atoms share their electrons, they end up developing a polarity due to higher electronegativity of oxygen compared to hydrogen. Oxygen ends up getting a slightly negative charge, while hydrogen a slightly positive charge. This allows different molecules of H2O to form a lattice structure with -ve polarity of oxygen of 1 molecule forming polar bonds with +ve polarity of hydrogen atoms of another molecule. These bonds are called Hydrogen bonds (and NOT polar bonds). When temperature are low, the vibrational energy of each molecule is very low, and not enough to break these hydrogen bonds between hydrogen and oxygen of different molecules. This makes it a solid, where the molecules can't slide past each other. But as temperature rises, the vibrational energy of each molecule increases, causing these hydrogen bonds to get weaker to a point where the molecules can slide past each other. This forms a liquid. If we keep on increasing the temperature, these hydrogen bonds break completely free, and different molecules become independent of each other. This forms the vapor or gas state. This video from Khan Academy explains it => https://www.khanacademy.org/science/chemistry/states-of-matter-and-intermolecular-forces/states-of-matter/v/states-of-matter
- London dispersion forces : Like hydrogen bonds, London dispersion forces are weak attractions between molecules. However, unlike hydrogen bonds, they can occur between atoms or molecules of any kind, and they depend on temporary imbalances in electron distribution.
