The Metallic Model

Bonding and Structure of Metals

    Metallic Lattice Structure

  • In metallic substances, metal atoms arrange themselves into a lattice structure, which is similar to ions
  • In this lattice, the valence electrons of the metal atoms become delocalized, meaning they are not bound to a single atom
    • This happens because the electrons are attracted to neighbouring atoms, causing them to move around freely
  • This loss of electrons forms metal cations, which find themselves situated in a "sea" of delocalized electrons

    Metallic Bonds:

  • Hold metals together, which stem from the strong electrostatic attraction between the lattice of positive metal ions and the sea of delocalized electrons
  • This type of bond is unique, as it does not occur between atoms or ions, but rather between cations and delocalized valence electrons
  • Metallic bonds are very strong, and they are non-directional because electrons pull on the cations from all directions

    Alloys

  • Metals can have impurities added to them in the form of other elements, where the modified metals are known as alloys
  • Alloys can be much stronger than regular metals, where the addition of elements with different sizes disrupts the metallic lattice structure and can make it more difficult for cations to slide over each other
  • Example: steel is a common alloy formed by adding small amounts of carbon to iron, which can make steel up to 1000x stronger than pure iron

Metallic Properties

    Lustre:

  • The shiny appearance of a substance, due to light being reflected off of it
  • Metals are generally shiny/lustrous, as light interacts with their delocalized electrons causing them to vibrate, where the vibrating electrons then radiate light outwards which creates the shiny appearance

    Sonority:

  • The ability of a substance to produce a deep or ringing sound when hit
  • Metals are generally sonorous, where they make a ringing sound when struck due to their delocalized electrons propagating sound energy throughout the metal
  • Metals with low densities are more sonorous than high-density metals, as they have more space between cations which allows electrons to move more freely

    Malleability:

  • The ability of a substance to be bent and reshaped under compression or stress
  • Metals are malleable, as the non-directionality of their bonds means that layers within the metallic lattice can shift under stress without affecting the bonds (the cations are still surround by delocalized electrons on all sides)

    Ductility:

  • The ability of a substance to be drawn out into a wire when stretched
  • Metals are ductile for the same reason that they are malleable, where they can easily be deformed without affecting their bonds and structure

    Electrical Conductivity:

  • The ability of charged particles to move through a region of space in a substance
  • Metals tend to be very electrically conductive, as they have many charged particles in the form of delocalized electrons which are free to move around the substance
  • Metals with more valence electrons will end up having more delocalized electrons, which increases the electrical conductivity of the metal

    Thermal Conductivity:

  • The ability of substance to transfer heat under a temperature difference
  • Metals tend to be very thermally conductive, as hotter regions of the substance result in delocalized electrons gaining kinetic energy, and their ability to move freely allows them to easily transfer this energy/heat to colder regions of the substance
  • Similar to electrical conductivity, metals with more valence/delocalized electrons tend to be more thermally conductive

Metallic Bond Strength

    Ionic Radius of Metals

  • Metal ions with a smaller radius will experience stronger metallic bonds
    • This is the same with ionic bonding, where electrostatic forces are inversely proportional to distance squared
  • Going down a group of metals on the periodic table, we expect the metallic bond strength to decrease as these metals will have more occupied energy levels and hence a larger ionic radius than metals higher up in the group

    Ionic Charge of Metals

  • Metal ions with a greater ionic charge will experience stronger metallic bonds
  • When metal cations have a more positive charge, they are more strongly attracted to delocalized electrons and experience a stronger metallic bond
  • A greater ionic charge on a metal cation indicates that it lost more valence electrons, which means there are more delocalized electrons for the cations to be electrostatically attracted to, thus increasing the metallic bond strength
  • Going across a period of metals on the periodic table, we expect the metallic bond strength to increase as these metals will have more valence electrons which increases their cationic charge and the number of delocalized electrons

    Melting Points

  • The melting point of metals depends on the strength of their metallic bonds, where stronger bonds result in higher melting points
  • As metallic bonds are typically very strong, metals can generally be found as solids at room temperature (except mercury)
  • Going across a period and up a group of metals in the s and p blocks of the periodic table will generally show an increase in the melting point, due to stronger metallic bonds

    Metal Hardness

  • The hardness of metals, or their resistance to deformation, increases with stronger metallic bonds
  • While most metals are hard and cannot be easily deformed, many metals in group 1 of the periodic can be cut with a knife due to their weaker metallic bonds

Transition Metals

    What are Transition Metals?

  • Transition metals are the metallic elements located in groups 3-12 on the periodic table
  • These metals are different from other types of metals, as they have valence electrons occupying d orbitals
  • As mentioned in ionic bonding, transition metals can form ions with different valence charges, giving them unique properties

    Melting Points of Transition Metals

  • Transition metals typically have much higher melting points than other metals
    • This is because they can have a lot more delocalized electrons from the d orbital, creating stronger bonds
  • There is no general trend across the period for their melting points, as there are many inconsistencies

    Hardness of Transition Metals

  • Transition metals typically are much harder/stronger than other metals
    • Like with melting points, this is due to stronger metallic bonds stemming from more delocalized electrons
  • While group 1 metals are soft and can be cut with a knife, transition metals are hard and usually cannot be cut this way

    Electrical Conductivity of Transition Metals

  • Like other metals, transition metals tend to be very good electrical conductors
    • This is again due to the high number of delocalized electrons, which are free to move around the substance
  • This properties of transition metals makes them useful for electrical wiring, where copper (a transition metal) is one of the most conductive metals and is used extensively in wiring