Bonding: Covalent, Ionic, and Metallic Bonds

Intramolecular Bonds vs Intermolecular Forces

Bonding occurs because atoms wish to arrange themselves in the most stable patterns possible. Stability is achieved through the completion of the atom's outermost electron orbits, thus obtaining the octet configuration in the valence shell and satisfying the octet rule. Atoms can do this by joining other atoms via chemical bonds.

  • Intramolecular bonds: bond atoms to other atoms, creating compounds.

  • Intermolecular forces: attract atoms and molecules to other atoms and molecules. This video will focus on intramolecular bonds, covalent, ionic, and metallic.

Covalent Bonds (Molecular Bonds)

A covalent bond, also known as a molecular bond, is formed when a pair or pairs of electrons are shared between two atoms to form a covalently bonded species or a molecular compound. As a result, the pair of shared electrons forms a new molecular orbital extending around the nuclei of both atoms to form a molecule. Covalent bonds are the most common bond in organic molecules, and you can get compounds with high molecular mass in this way.

  • Macromolecules can be linear, branched, or cross linked. You can also get a crystal network or lattice in which each atom is covalently bonded to its neighbors to form one large molecule.

  • Sometimes you can have different allotropes, different forms of an element in its natural state. For example, graphite and diamonds are both allotropes of carbon. In graphite, carbon forms three covalent bonds, leaving one electron delocalized, and forms planes. In diamond, the hardest material known, bonding occurs in a tetrahedral geometry.

  • Covalent bonds come in two forms, polar and nonpolar, depending on the electronegativities of the atoms involved. Electronegativity is the tendency of an atom to attract a shared pair of bonded electrons in its combined state. Electropositivity is the opposite.

Polar Covalent Bonds

A polar covalent bond is formed between two different nonmetal atoms with different electronegativities, which results in them sharing electrons unequally.

  • The bonding electron pair is closer to one of the nuclei depending on the relative electronegativities of the two atoms. So although the overall molecule is neutral, this results in one end of the molecule being slightly negatively charged and the other slightly positively charged, and this charge distribution is denoted by a dipole arrow and a lowercase delta with a charged superscript: ext{—}
    ightarrow ext{(}oldsymbol{oldsymbol{oldsymbol{bd}}^+ / oldsymbol{oldsymbol{bd}}^- ext{)}, where appropriate notations use oldsymbol{oldsymbol{}7d^+ and oldsymbol{bd^-}.

  • Substances with polar covalent bonds have higher melting and boiling points. They are also soluble in polar compounds, such as water. Examples include ext{HCl} and ext{HF}.

Nonpolar Covalent Bonds

A nonpolar covalent bond is formed between two atoms with the same or very similar electronegativity. They share their electrons equally.

  • These substances tend to exist as gases and rarely as liquids. They have low melting and boiling points, and are soluble in nonpolar solvents. Examples include ext{H}2, ext{O}2, and ext{N}_2.

Why Covalent or Ionic Bonds Form

In covalent bonding, the difference between the electronegativities of the two atoms is insufficient for an electron transfer to occur and to form ions. The atoms involved have high ionization energies.

  • Ionization energy is the energy needed to remove electrons from a neutral atom, forming a positively charged ion. The first ionization energy of an atom is the energy needed to remove the outermost, highest energy electron when it is neutral and in gas phase.

  • This ionization energy is inversely proportional to atomic size. As you travel right on the periodic table, you have more protons, which attract the electrons in closer to the nucleus and make it harder for them to escape. As you go down, with each row, you get a new shell of electrons, and so the valence electrons are further from the nucleus and easier to take away.

  • After covalent bonds are formed, covalent bonds rarely break spontaneously. They are stronger than ionic bonds (as stated in the transcript).

Ionic Bonds

Ionic bonds hold atoms together via strong electrostatic attraction between charged ions which differ significantly in electronegativity. As a result, the less electronegative ion transfers electrons to the more electronegative ion. The result is an anion with a negative charge, and a cation with a positive charge. These opposite charges attract to form a compound.

  • Example: ext{NaCl}. Chloride needs an electron to complete its octet. The easiest way for sodium to have a full octet valence shell is to lose its one valence electron, so sodium donates an electron to chloride. As a result, sodium is now a ext{Na}^+, and chloride is now a ext{Cl}^-. Note that ionic compounds can dissociate into ions in solution.

  • The likelihood of an ionic bond forming depends on the radius of the atoms. A larger radius increases the likelihood of ionic bonding by decreasing the ionization energy. Because ionic bonds form between atoms with big differences in electronegativity, they form between a metal and a nonmetal.

Metallic Bonds

Metallic bonds are formed between metals, metalloids, and alloys. The bond is formed between positively charged atoms that share electrons. The valence electrons go from one atom to the next, continuously moving throughout the entire space. These free electrons have been described as a sea of electrons. Examples of metallic bonds include ext{Au}, ext{Ag}, ext{Fe}.

  • Different types of bonds result in different properties. It is important to remember that covalent bonds form between two nonmetals, ionic bonds between a metal and a nonmetal, and metallic bonds form with metals.

  • Molecules with covalent bonds exist as solids, liquids, and gases at room temperature, while substances with ionic and metallic bonds exist in solid state at room temperature.

Connections to Real-World Relevance and Implications

  • allotropes (e.g., carbon: graphite vs diamond) illustrate how the same element can form very different structures with distinct properties depending on bonding.

  • polar vs nonpolar bonds help explain solubility trends (polar solvents dissolve polar substances; nonpolar solvents dissolve nonpolar substances).

  • understanding bond types informs predictions about melting/boiling points and the physical state of substances at room temperature.

Quick Summary

  • Covalent bonds: form between two nonmetals; electrons are shared; can be polar or nonpolar.

  • Ionic bonds: form between a metal and a nonmetal; electrons transfer, producing cations and anions; ionic compounds can dissociate in solution.

  • Metallic bonds: bonding in metals and alloys; electrons are delocalized as a sea of electrons; leads to conductivity and other metallic properties.

Next Video

Check out the next video to learn about intermolecular forces.