In high school chemistry, we learn that Covalent bonds involve the sharing of one or more pairs of electrons between two atoms. We can represent these bonding pair of electrons in a diagram. The electrons are shown to be midway between the two atoms . For a homonuclear bond (two atoms are the same elements), this is approximately true as the electrons on average will be found in the center of the bond. For a heteronuclear bond, for instance, CO, carbon monoxide, the electrons do not necessarily average in the center of the bond. This is because different elements attract the bonding pair of electrons differently.
We use the Pauling scale as a guide to determine whether a bond is covalent or ionic.
| Bond Type | Electronegativity Difference | Characteristics |
|---|---|---|
| Ionic Bonds | Greater than 1.7 | Significant electronegativity difference One atom donates electrons, the other accepts Formation of cations and anions Na → Cl (Sodium Chloride) – Sodium (Na) loses an electron while chlorine (Cl) gains that electron. The ions (Na⁺) and (Cl⁻) are formed. |
| Polar Covalent Bonds | Between 0.5 and 1.7 | Moderate electronegativity difference Unequal sharing of electrons, leading to partial positive and negative charges May exhibit dipole moments (H-O) is a polar covalent bond, with H more positive and O more negative |
| Non-Polar Covalent Bonds | Less than 0.5 | Minimal electronegativity difference Nearly equal sharing of electrons No significant charges on atoms (Cl-Cl) is a nonpolar covalent bond, where electrons are shared equally between the two chlorine atoms |
Bonding spectrum
Electronegativity is an important concept in chemistry when elements bond together. When elements bond, it always involves electrons. We’ve learned that in some bonds, electrons are shared (covalent bonds), while in others, they’re transferred from one element to another (ionic bonds). But real bonding is more complex than just these two types.
In reality, there’s a spectrum of bonding possibilities. Bonds aren’t strictly just covalent or purely ionic; they can be a mix of both. Sometimes, electrons are pulled so strongly toward one element that they’re considered transferred, and we call this an ionic bond. A classic example of an ionic bond is the salt you put on your food, sodium chloride (NaCl). So, bonding is a bit like a sliding scale between sharing and transferring electrons, with many variations in between.
In summary, the categorization of these compounds as ionic or covalent is determined by the electronegativity difference between the elements involved. Compounds with a large electronegativity difference tend to be ionic, while compounds with a smaller electronegativity difference or similar electronegativities tend to be covalent.
Exploring Bonding Patterns Involving Metalloids
Metalloids can form both ionic and covalent bonds. Metalloids, which are elements found along the boundary between metals and non-metals on the periodic table, can exhibit a variety of bonding patterns depending on the specific element and the elements they are bonding with. The bonding behavior of metalloids can be quite diverse.
Generally, when metalloids form bonds, they tend to form covalent bonds rather than ionic bonds. This preference for covalent bonding is due to their intermediate electronegativity values, which lie between those of metals and non-metals. As a result, metalloids can share electrons with other elements in a more balanced manner, typical of covalent bonding. However, the bonding nature of metalloids can vary depending on the elements they interact with and the specific environmental conditions. In some cases, metalloids can form ionic bonds, especially when they react with elements that have significantly different electronegativities. This versatility in bonding allows metalloids to participate in a wide variety of chemical compounds and reactions, reflecting their position as a bridge between metals and nonmetals in the periodic table.
| Compound | METALLOID + METAL |
|---|---|
| Gallium Arsenide (GaAs) | GaAs exhibits a covalent character in its bonding, but it is not purely covalent or purely ionic. It has a mix of both ionic and covalent bonds, with a majority (69%) being covalent in nature and a smaller portion (31%) being ionic. |
| Indium Phosphide (InP) | Indium phosphide is a binary semiconductor material composed of phosphorus and indium. It has a high degree of covalent bonding. |
| Boron Nitride (BN) | Boron nitride (BN) is a binary compound of boron (B) and nitrogen (N) atoms. It exhibits covalent bonds, where boron and nitrogen atoms share electrons, similar to carbon in organic compounds. The bonding in BN varies with its structural forms, including boron nitrides and boron nitride nanotubes. |
| Compound | METALLOID + NON-METAL |
|---|---|
| Silicon Dioxide (SiO₂) | SiO₂ is a covalent compound. Silicon and oxygen share electrons to form strong covalent bonds, creating a molecular structure. |
| Germanium Tetrachloride (GeCl₄) | GeCl₄ is a covalent compound. Germanium and chlorine atoms share electrons through covalent bonds. |
| Arsenic Trichloride (AsCl₃) | AsCl₃ is also a covalent compound. Arsenic and chlorine atoms share electrons in a covalent bonding arrangement. |
| Arsenic(III) Oxide (As₂O₃) or Arsenic Trioxide | Arsenic(III) oxide, also known as arsenous oxide – it contains arsenic and oxide ions. Arsenic(III) oxide is sometimes named as arsenic trioxide due to its covalent nature. This dual bonding nature of As2O3 is a reflection of the intermediate properties of metalloids like arsenic, capable of displaying both metallic and nonmetallic bonding characteristics depending on the environmental context and the nature of their interaction with other elements. |
Metallic
Metallic bonds are a unique type of chemical bonding that exists in metals and metalloids. In metallic bonds, electrons are delocalized or free to move throughout the entire crystal lattice of the metal. Unlike in ionic bonds, where electrons are transferred from one atom to another, and covalent bonds, where electrons are shared between atoms, metallic bonds involve a “sea” of electrons shared by all atoms in the metal lattice.
This delocalization of electrons in metallic bonds gives metals their unique properties, such as electrical conductivity, malleability, and ductility. It also allows metals to form lustrous, shiny surfaces.
In summary, metallic bonds involve the sharing of a pool of electrons among all atoms in a metal, giving rise to the characteristic properties of metals. These bonds are distinct from both ionic and covalent bonds in their electron-sharing behavior.
Summary
| Property | Ionic Bonds | Covalent Bonds | Metallic Bonds |
| Type of element involved | Metal + Nonmetal | Non-metal + Non-metal or Metalloid + Non-metal | Metals |
| Electron Behavior | Transfer of electrons | Sharing of electrons | Sea of delocalized electrons |
| Strength of Bonds | Generally Strong | Varies (Strong in network covalent, weaker in molecular) | Generally strong |
| Melting and Boiling Points | High | Lower than ionic | High |
| Electrical conductivity | Good when molten or dissolved | Poor | Very Good |
