electron sea

Examples of electron sea in the following topics:

• Bonding in Metals: The Electron Sea Model

  • Metallic bonding may be described as the sharing of free electrons among a lattice of positively charged metal ions.
  • That is to say, instead of orbiting their respective metal atoms, they form a "sea" of electrons that surrounds the positively charged atomic nuclei of the interacting metal ions.
  • The electrons then move freely throughout the space between the atomic nuclei.
  • Metals are good conductors of electricity because the electrons in the electron sea are free to flow and carry electric current.
  • Positive atomic nuclei surrounded by a sea of delocalized electrons (the blue dots).

• General Properties of Metals

  • In a metal, atoms readily lose electrons to form positive ions (cations).
  • Those ions are surrounded by de-localized electrons, which are responsible for the conductivity.
  • Metals can be viewed as a collection of atoms embedded in a sea of electrons, which are highly mobile.
  • Metals are usually inclined to form cations through electron loss.
  • The "sea of electrons" is free to flow about the crystal of positive metal ions.

• Periodic Trends in Metallic Properties

  • When two elements are joined in a chemical bond, the element that attracts the shared electrons more strongly has more electronegativity.
  • The simplest conception of metals is a lattice of positive ions immersed in a "sea of electrons" that can migrate freely throughout the solid.
  • In effect, the electropositive nature of the metallic atoms allows their valence electrons to exist as a mobile fluid.
  • Because each ion is surrounded by the electron fluid in all directions, the bonding has no directional properties; this accounts for the high malleability and ductility of metals.

• Metallic Crystals

  • In a metal, atoms readily lose electrons to form positive ions (cations).
  • These ions are surrounded by delocalized electrons, which are responsible for conductivity.
  • The strength of a metal derives from the electrostatic attraction between the lattice of positive ions and the "sea" of valence electrons in which they are immersed.
  • Metallic solids are known and valued for these qualities, which derive from the non-directional nature of the attractions between the atomic nuclei and the sea of electrons.
  • Loosely bound and mobile electrons surround the positive nuclei of metal atoms.

• Conductors

  • In metallic conductors such as copper or aluminum, the movable charged particles are electrons.
  • Band theory models the behavior of electrons in solids by postulating the existence of energy bands.
  • The electrons of a single isolated atom occupy atomic orbitals, which form a discrete set of energy levels.
  • This produces a number of molecular orbitals proportional to the number of valence electrons.
  • For instance, the sea of electrons causes most metals to act both as electrical and thermal conductors.

• Properties of Oxygen

  • The two oxygen atoms in diatomic oxygen are chemically bonded to each other with a spin triplet electron configuration.
  • This bond has a bond order of two and is often simplified in descriptions as a double bond, or as a combination of one two-electron bond and two three-electron bonds.
  • The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals.
  • Singlet oxygen is a name given to several higher-energy species of molecular O2 in which all the electron spins are paired.
  • At 5 °C the solubility increases to 9.0 mL (50 percent more than at 25 °C) per liter for water and 7.2 mL (45 percent more) per liter for sea water.

• The Shielding Effect and Effective Nuclear Charge

  • Therefore, these electrons are not as strongly bound as electrons closer to the nucleus.
  • Ne has 10 electrons.
  • Thus the number of nonvalence electrons is 2 (10 total electrons - 8 valence).
  • Flourine has 9 electrons but F- has gained an electron and thus has 10.
  • Sodium has 11 electrons but the Na+ ion has lost an electron and thus has 10.

• Electronic Effects

  • Such charges are produced by removing (or adding) electrons from (or to) an object.
  • Electron deficient species, which may or may not be positively charged, are attracted to electron rich species, which may or may not be negatively charged.
  • Electrophiles: Electron deficient atoms, molecules or ions that seek electron rich reaction partners.
  • Nucleophiles: Electron rich atoms, molecules or ions that seek electron deficient reaction partners.

• General Rules for Assigning Electrons to Atomic Orbitals

  • An atom's electrons exist in discrete atomic orbitals, and the atom's electron configuration can be determined using a set of guidelines.
  • This nucleus is surrounded by electrons.
  • An atom's electron shell can accommodate 2n2 electrons, where n is the energy level.
  • An element's electron configuration is the arrangement of the electrons in the shells.
  • Electrons that occur together in an orbital are called an electron pair.

• The Building-Up (Aufbau) Principle

  • The Aufbau principle determines an atom's electron configuration by adding electrons to atomic orbitals following a defined set of rules.
  • As electrons are added, they assume the most stable shells with respect to the nucleus and the electrons already present.
  • When there are two electrons in an orbital, the electrons are called an electron pair.
  • If the orbital only has one electron, this electron is called an unpaired electron.
  • Put one electron in each of the three p orbitals in the second energy level (the 2p orbitals) and then if there are still electrons remaining, go back and place a second electron in each of the 2p orbitals to complete the electron pairs.