It is vital for you to study the electron sea theory if you want to understand metallic bonding. Almost 3/4th of all the chemical elements in the Periodic Table are metals. These are predominantly located in Groups 1-12 of the Periodic Table. Metal atoms present in Groups 1 and 2 are known as alkali and alkaline earth metals, respectively. Contrarily, the metals present in Groups 3-12 are called transition metal elements.
What is metallic bonding?
Metal reactivity is determined by its electropositivity, i.e., the ability to lose electrons. Positively charged metal ions called cations are formed due to this electron loss. What do you think is the force of attraction between oppositely charged cations and electrons called in metals? The name is obvious, i.e., metallic bonding.
Then the question is, where does the name electron sea theory come from? Let’s find out through this article.
What is electron sea theory of metallic bonding – Definition
The electron sea theory, also known as the electron gas theory of metallic bonding, was introduced by Paul Drude back in the early 1900s. Henrik Lorentz further developed it in the later years. This theory discusses the bonding present between identical metal atoms.
As per the electron sea theory, highly electropositive metal atoms lose all their valence electrons. This results in positively charged cations embedded in a sea of delocalized electrons. In a metallic crystal, the cations occupy fixed positions at measurable distances from each other. Contrarily, the electrons are free to move throughout the crystal, just as water moves freely in the sea. The strong force of attraction that bind the positively charged ions and the negatively charged electrons in the crystal lattice is thus known as metallic bonding.
Sea of electrons and metallic bonding
The greater the number of electrons lost by the metal atoms to the sea of delocalized electrons, the smaller the size of a metal cation and the higher its charge. As a result, a stronger metallic bond is formed.
The strength of a metallic bond depends on:
- Number of electrons lost by the metal atom.
- Charge present on the metal cations.
- Cation size.
A large number of metal ions arrange in a three-dimensional lattice arrangement known as a metallic crystal. You may note that as per the electron sea model, the delocalized electrons are not completely lost but are still present within the structure, even though they are not associated with any particular atom. Therefore, the sodium or magnesium metal is still represented using their atomic symbols; Na or Mg, not as Na+ or Mg2+, respectively.
Envisaging metal cations floating in a sea of electrons, as per the electron gas theory, helps us determine numerous valuable metallic properties.
What are the properties of metals in the electron sea model?
Metals possess unique physicochemical properties such as:
- Metals possess a low electronegativity and/or electron affinity.
- Metals are highly electropositive.
- Metal atoms have a low ionization energy.
- Metals are solid with high melting points.
- Metals are good conductors of heat and electricity.
- Metals are malleable and ductile.
- Metals exhibit a lustrous (shiny) appearance.
- Metals have a high elasticity.
- Metals emit electrons when heated.
What is interesting for us to note is that the chemical properties of metals facilitate ‘metallic bond formation” as per the electron sea model. In contrast, their physical properties can be explained as an ‘outcome” of the electron sea model.
Let us explain it to you more clearly by discussing each of the above properties, one at a time.
Electronegativity of an elemental atom is defined as its ability to attract electrons. In other words, it is the affinity of an atom towards electrons, thus also known as electron affinity. Low metal electronegativity means metal atoms do not attract electrons; rather, they are ready to lose some or all of their valence electrons into the sea of delocalized electrons.
Electropositivity is the exact opposite of an atom’s electronegativity. The lower the electronegativity of a metal atom, the higher its electropositivity; thus, it can readily lose electrons and form positively charged metal ions.
Low ionization energy
Ionization energy is referred to as the amount of energy released when a metal atom changes to an ion (M = M+ + e–). Low ionization energy denotes it is easier for a metal atom to lose an electron and change into an ion.
High melting point
The high melting point of metals (for e.g. copper m.pt = 1084°C) is accredited to the strong metallic bonding present in them. The higher the charge density (greater charge, smaller radius) in a metallic crystal, the higher its melting point.
For instance, magnesium metal has a much higher melting point (650°C) than the sodium metal (m.pt = 97.8°C). This is because Mg loses 2 electrons per atom while Na loses only a single valence electron available. Additionally, atomic size decreases across the period. Thus, Mg has a higher overall charge density than Na. Consequently, more energy is required to break metallic bonding in magnesium as compared to sodium metal.
High thermal conductivity
The presence of numerous delocalized electrons in a metallic crystal allows a rapid thermal or heat conduction. When one part of the metal is heated, the electrons present in that part gain kinetic energy. Moving rapidly from the hotter region of the metal to colder, electrons conduct heat.
High electrical conductivity
Mobile electrons present in the sea of delocalized electrons also help the conduction of electrical current through the metal.
Malleability and ductility
Malleability means metals can be beaten into sheets, while ductility refers to their ability to be drawn into wires.
The electron sea or electron gas present around the metal cations acts as a support or cushioning agent. When the metal is hammered or when shear stress is applied to it, the metal crystal does not break. Rather, the cation layers slip past one another. These are re-arranged, supported by the sea of delocalized electrons. In short, the metal crystal is not rigid; it only gets deformed, unlike an ionic crystal lattice that falls apart on appealing slight pressure.
An example of metallic malleability and ductility lies in the different shapes and forms in which we use metals, such as gold, silver, tin, copper, lead, etc., in our everyday lives.
The shine in metal surfaces can also be explained using the electron gas model. When irradiated with a light beam, the delocalized electrons absorb certain wavelengths of light energy. This energy absorption leads to electronic excitation followed by de-excitation in the metal lattice. The de-excitation results in visible radiation emission at all angles; thus, metals appear shiny when light falls on their surfaces.
The ability of metals to change their shape without fracturing also accounts for their high elasticity. When a temporary external force is removed, the layers of metal cations and delocalized electrons readjust at their original positions. This property to readily gain their original shape and size is known as high metal elasticity.
The photoelectric effect
A metal surface ejects electrons when heated or when light falls on it. This is known as the photoelectric effect. The emitted electrons are known as photoelectrons. As the electrons move freely in all directions in a metal, they can thus easily leave the crystal surface when provided with the right amount of energy. However, the energy input should be greater than a specific threshold frequency, which differs for different types of metals.
The photoelectric effect exhibited by metals is widely utilized in manufacturing electronic equipment.
If you found this article helpful, you may also like our other articles related to metallic bonding:
Last but not least, here is a video for you to revise your concepts on metallic bonds.
1. Chaudhary, Ghulam Rasool. 2006. Inorganic chemistry
2. Sanaullah. 2016. Inorganic Chemistry.
3. Steurer, W. 1996. ‘CHAPTER 1 – CRYSTAL STRUCTURE OF THE METALLIC ELEMENTS.’ in Robert W. Cahn and Peter Haasen† (eds.), Physical Metallurgy (Fourth Edition) (North-Holland: Oxford).