Basics of Interaction and Energetics
Any chemical transformation is based on the electrostatic attraction between positive (cations) and negative (anions) charges. The main criterion for the formation of any bond is the lowering of the system's energy.
For example, when an isolated hydrogen atom ($H$) and a free proton ($p^+$) approach each other, the simplest molecular system—the $H_2^+$ ion—is formed. If the initial energy of the isolated particles is taken as zero, their combination causes the potential energy to drop to –270 kJ/mol.
This energy gain is called the chemical bond energy ($E_{bond}$). It is precisely what ensures the stability of the new structure compared to isolated particles. The value of $E_{bond}$ is measured in kJ/mol (per $6.02 \cdot 10^{23}$ bonds).
Covalent bond
It arises between uncharged atoms through the formation of a shared bonding electron pair. From the standpoint of quantum mechanics, a bond is formed if a single electron is capable of simultaneously holding two atomic nuclei.
There are two main pathways of formation:
- Exchange mechanism: each atom provides an unpaired electron. Mutual penetration (overlapping) of electron clouds occurs. For example, in a hydrogen molecule ($H_2$), the distance between the nuclei decreases to 75 pm, which is less than twice the radius of an isolated atom (106 pm). The energy of the system decreases by 436 kJ/mol in the process.
- Donor-acceptor mechanism: one participant (donor, nucleophile, base) provides a lone electron pair, while the second (acceptor, electrophile, acid) provides a vacant valence orbital. A classic example is the reaction of ammonia ($:NH_3$) with a proton ($H^+$) to form the ammonium ion ($NH_4^+$).
Only outer (valence) electrons participate in cloud overlapping. Electrons in inner orbitals practically do not change their state.
Basics of Molecular Orbital Theory
During atomic interaction, a new molecular orbital is formed, shared by the nuclei of the interacting atoms. Electron density is concentrated between them, and the energy of such electrons is lower than in separate atoms.
However, as the number of electrons increases, different types of levels begin to fill:
- If additional electrons appear in the system (such as in the $H_2^-$ ion, where there are three), the third electron occupies an energetically unfavorable antibonding orbital. This does not strengthen but weakens the bond.
- In a hypothetical helium molecule ($He_2$), four electrons would be distributed equally between the bonding and antibonding levels. There is no net energy gain, so a chemical bond is not formed.
Conclusion: For successful bond formation, the combining atoms must have unpaired electrons to avoid filling antibonding levels.
Ionic and Metallic Bonds
- Ionic bond (Ionic bond) is formed exclusively due to electrostatic attraction forces between ready-made ions. Atoms completely give up or accept electrons. As a result, instead of discrete molecules, durable ionic crystal lattices are formed. Most characteristic of salts (e.g., $NaCl$).
- Metallic bond is characteristic of elements with a small number of outer electrons and vacant orbitals. During crystal formation, outer electrons become delocalized—they occupy closely spaced orbitals, forming a conduction band. Electrons move freely within it ("electron gas"). The bond is due to the interaction of this electron gas with the positively charged ion core of the crystal structure, which also ensures high electrical conductivity.
Element Valency
Valency of an atom is determined by the number of its unpaired electrons in the ground or excited state. Excitation (transition of an electron to a higher sublevel) is only possible within a single energy level, since a transition to the next level is energetically unfavorable.
Elements of the 2nd period exhibit clear trends:
- Carbon (C) has the highest valency. The number of its valence orbitals equals the number of valence electrons, and all of them can be unpaired.
- In the series N $\rightarrow$ O $\rightarrow$ F $\rightarrow$ Ne, the total number of valence electrons increases, but the number of unpaired electrons decreases. For neon, valency drops to zero.
- Lone pairs in valence orbitals allow an atom to form additional bonds (as in nitrogen in the ammonium ion). Because of this, the total number of bonds of an atom can exceed its formal valency.
Briefly: answers to topical questions
What types of chemical bonds are found in bioorganic compounds?
Bioorganic compounds contain both covalent and non-covalent interactions.
- Disulfide bond is a covalent bond formed through the oxidation of SH-groups of cysteine residues and their interaction with each other.
- Ionic bond is an electrostatic interaction between positively and negatively charged functional groups.
- Hydrogen bond is involved in stabilizing the secondary structure of a protein; it occurs between the oxygen atom of one peptide group and the hydrogen atom of another peptide group. Hydrogen bonds also form between complementary DNA nitrogenous bases.
- Hydrophobic interactions occur between hydrophobic amino acid radicals, as well as between stacked DNA nitrogenous bases.
Which organogen atoms act as electron pair donors in biological molecules?
Nitrogen and oxygen atoms can act as electron pair donors.
- Nitrogen atom provides a lone electron pair and acts as a donor, for example, in ammonia and amines when forming donor-acceptor bonds with metal ions.
- Oxygen atom can also participate in the formation of donor-acceptor bonds via lone electron pairs, for example, during the complexation of $Mg^{2+}$ and $Ca^{2+}$ ions with complexones.
When such bonds are formed, the donor provides an electron pair and the acceptor accepts it into a vacant valence orbital.
How does a carbon atom transition into an excited state?
The transition of a carbon atom into an excited state occurs due to the unpairing of $s$-electrons and the transition of one electron to the $p$-sublevel.
In the ground state, group IV element atoms have 2 unpaired electrons in the $p$-sublevel and exhibit a valency of II. In the excited state, the transition occurs according to the scheme: $$ns^2np^2 \xrightarrow{E} ns^1np^3$$ As a result, 4 unpaired electrons are formed, and carbon exhibits a valency of IV in compounds.
The energy expended on electron excitation is compensated by the energy released during the formation of additional chemical bonds. Upon bond formation in a molecule, the atom receives additional electrons to populate orbitals, preventing a return to the ground state.
What is a hydrogen bond and between which atoms does it form?
A hydrogen bond is an interaction in which hydrogen unites two molecules or parts of a molecule, acting as a "bridge" between two atoms.
It forms during an incomplete proton transfer: a bond is created with a new atom without completely breaking the previous bond.
In organic compounds, a hydrogen bond arises between the hydrogen atom of one polar group (such as a hydroxyl group) and an electronegative atom (such as oxygen) of another molecule. The condition for its formation is the presence of a hydrogen atom directly bonded to oxygen in the –OH group.
In proteins, hydrogen bonds form between the oxygen atom of one peptide group and the hydrogen atom of another peptide group, stabilizing the secondary structure.
To remember the roles of participants in bond formation: Donor Gives (an electron pair), Acceptor Rents (an empty orbital).
Class and exam questions
Does a chemical bond necessarily form when atoms collide?
No. A bond is formed only if this process is accompanied by the release of energy (a decrease in the system's energy) without forcing the population of antibonding orbitals.
How do the bonds in the ammonium ion (NH₄⁺) differ?
Formally, one bond is formed via the donor-acceptor mechanism, while the remaining three are formed via the exchange mechanism. However, in reality, all four bonds in the resulting ion are completely identical and indistinguishable.
Why is the H₂⁺ ion more stable than an isolated proton and hydrogen atom?
When a proton and a hydrogen atom approach each other to the optimal distance (106 pm), the electron density is distributed between the nuclei. The potential energy of the system decreases by 270 kJ/mol, which makes the H₂⁺ ion more stable.
What else to study on this topic
- Quantum-mechanical model of molecular orbitals
- Properties of the hypothetical molecular ion H₂⁻
- Energetics of electron transitions during atom excitation
- Relationship between carbon valence and the number of valence orbitals
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