Name: electron sea model in metals
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 32 s
Channel: JoVE

K12 Standards

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North Carolina K-12 Science Standard Course of Study

9-12

Physical Science

Matter and its Interactions

PS.PSc.2.1

Bonding in Metals

A metallic bond is a bond between two metal atoms. Compared to nonmetals, metals have low ionization energies, allowing them to lose valence electrons easily. This gives the metallic bond distinct properties in contrast to ionic and covalent bonds.&nbsp;

Metallic bonds and most of their properties can be explained using the simple electron sea model. Consider the metal potassium. Due to the low ionization energy, each potassium atom can easily lose its valence electron to become a cation.&nbsp;

These potassium cations are held together in close-packing because of their attraction to the negatively charged sea of electrons. These electrons are not confined to any single ion but are evenly distributed and relatively free to move within the metal. &nbsp;

The electron sea model accounts for several important characteristics of metals. For example, when a voltage difference is applied to a metal wire, like copper wire, the negatively charged electrons move freely towards the wire&rsquo;s positive end, generating an electric current. This is why most metals are excellent conductors of electricity.&nbsp;

In contrast, ionic compounds are nonconductors of electricity in their solid form but can conduct electricity when dissolved in water. This is because, in a crystalline ionic bond, electrons are transferred from the metal to the nonmetal, but remain localized to one ion.&nbsp;

However, when dissolved in water, the cations and anions dissociate and can move when subjected to a potential difference, creating an electric current.

Metals are also excellent thermal conductors. According to the electron sea model, when heat is applied to one end of the metal, the electrons move freely and quickly disperse the heat throughout the metal.

Metals can easily be pounded into sheets due to their malleability or into wires due to the ductility property. Since there are no localized bonds in metals, the metal atoms can slide past each other allowing easy deformity. Electrons then flow into the new shape to accommodate the deformity.&nbsp; 
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bonding in metals

Metallic bonds are formed between two metal atoms. A simplified model to describe metallic bonding has been developed by Paul Dr&uuml;de called the &ldquo;Electron Sea Model&rdquo;.&nbsp;

<h4>Electron Sea Model</h4>

Most metal atoms do not possess enough valence electrons to enter into an ionic or covalent bonding. However, the valence electrons in metal atoms are loosely held due to their low electronegativity or attraction with the nucleus. The ionization energy of metal atoms (energy required to remove an electron from the atom) is low, facilitating the easy removal of valence electrons from the parent atom. The atom forms a positively charged metal ion, while the free outer electrons exist as negatively charged delocalized electron clouds. These electrons can be shared by multiple neighboring metal-cations through a strong, attractive force between these negatively and positively charged species. Such an attractive force between the negatively charged electrons and metal cations is called metallic bonds, holding the atoms together. This electron sea model accounts for most physical properties of metals such as conductance to heat and electricity, high melting and boiling points, malleability, and ductility.

<h4>Metallic Solids</h4>

The electron sea model accounts for several metallic properties, including high thermal and electrical conductivity, metallic luster, ductility, and malleability. The delocalized electrons can conduct both electricity and heat from one end of the metal to another with low resistance. The metallic bond is not between two specific metal atoms, but between metal ions and many delocalized electrons, allowing metals to deform under pressure and heat without shattering or breaking. Different metals, such as iron, mercury, or copper, differ in their physical properties, reflecting the difference in metallic bond&nbsp;strength among the metals.

Metallic solids such as crystals of copper, aluminum, and iron are formed by metal atoms: all exhibit high thermal and electrical conductivity, metallic luster, and malleability. Many are very hard and quite strong. Because of their malleability (the ability to deform under pressure or hammering), they do not shatter and, therefore, make useful construction materials. The melting points of the metals vary widely. Mercury is a liquid at room temperature, and the alkali metals melt below 200 &deg;C. Several post-transition metals also have low melting points, whereas the transition metals melt at temperatures above 1000 &deg;C. These differences reflect differences in the strength of metallic bonding among the metals.

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter">Openstax, Chemistry 2e, Section 10.5: The Solid State of Matter.</a>

Metallic bonds are formed between two metal atoms. A simplified model to describe metallic bonding has been developed by Paul Dr&uuml;de called the ...

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Chemistry

Chemical Bonding: Basic Concepts

KEY TERMS AND CONCEPTS

electron sea model in metals

Electron sea model explains how metallic bonding holds metals together. In this model, most metal atoms do not have enough valence electrons to form ionic or covalent bonds in the usual way. Their valence electrons are held loosely because metals have low electronegativity, and their ionization energy is low, so electrons are removed easily.

When a metal atom loses these outer electrons, it becomes a positively charged metal ion. The freed electrons form a delocalized electron cloud around many neighboring metal cations. Metallic bonds are the attractive forces between the negative electrons and the positive ions. These bonds do not link just two atoms. They connect metal ions to many shared electrons throughout the solid.

This electron sea model helps explain the main properties of metallic solids. Delocalized electrons carry heat and electricity with low resistance, so metals are good conductors. Metals also have metallic luster, high melting and boiling points, ductility, and malleability. Because the bonding is spread through the structure, metals can deform under pressure or heat without shattering or breaking.

Different metals can have different physical properties because their metallic bond strength is not the same. Examples of metallic solids include copper, aluminum, and iron crystals, and all show high thermal and electrical conductivity, luster, and malleability. Many metals are very hard and strong, which makes them useful as construction materials. Their melting points also vary widely. Mercury is liquid at room temperature, alkali metals melt below 200 °C, several post-transition metals melt at low temperatures, and transition metals often melt above 1000 °C. These differences reflect changes in metallic bonding strength.

Electron Sea Model in Metals

types of chemical bonds

Chemical Bonding and Bond Properties

Chemical bonding explains how atoms and ions join to form molecules and compounds. American chemist Gilbert N. Lewis pioneered chemical bonding theory with the Lewis model. This model helps explain why bonds form, why some are stronger than others, and why compounds have fixed formulas like H2O instead of H4O. Ionic bonding involves ions, which are atoms or molecules with an electrical charge. A cation is a positive ion formed when a neutral atom loses one or more electrons from its valence...

lewis symbols and the octet rule

Valence Electrons, Lewis Symbols, and Bonding

Chemical bonds form when atoms or ions interact in ways that lower the potential energy of a molecule. Gilbert N. Lewis created a simple model to show bond formation using valence electrons, the electrons in the outer shell. In this model, electrons are shared or transferred to help atoms reach a stable electron configuration. The Lewis model helps describe two common bond types. An ionic bond forms when electrons move from a metal to a nonmetal. A covalent bond forms when two nonmetals share...

ionic bonding and electron transfer

Ionic Compounds and Crystal Lattices

Ionic compounds form when positive and negative ions attract each other. A cation is a positive ion, and it forms when an atom loses one or more valence electrons. An anion is a negative ion, and it forms when an atom gains one or more valence electrons. Compounds made of these ions are called ionic compounds, or salts, and the force between them is an ionic bond. The properties of ionic solids help show how strong this attraction is. They usually form rigid, brittle crystals and have high...

the born-haber cycle

Predicting Ionic Stability with Lattice Energy

Lattice energy helps explain why ionic compounds are stable. It measures the strength of attraction between positive and negative ions in a crystal. In one common definition, lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions, so it has a positive value. Another definition uses the opposite sign and describes the energy released when gaseous ions form a lattice. When checking data, it is important to confirm which convention is being used. In either...

trends in lattice energy: ion size and charge

trends-in-lattice-energy-ion-size-and-charge

Lattice Energy Patterns in Ionic Solids

Lattice energy shows how strongly ions hold an ionic solid together. It is the energy required to separate one mole of a solid into its gaseous ions. For sodium chloride, this process requires 769 kJ per mole. The opposite convention may also be used, where lattice energy is the energy released when gaseous ions form the solid. In either case, a larger magnitude means a more stable ionic compound. For ionic crystals, lattice energy follows a pattern based on charge and ion size. The...

covalent bonding and lewis structures

Shared Electrons in Covalent Molecules

Covalent molecules form when atoms share electrons instead of transferring them. This bonding usually happens between nonmetal atoms with similar tendencies to attract electrons. Because the atoms pull on the electrons in a shared way, the bond is held together by mutual attraction. Covalent compounds often have physical properties that differ from ionic compounds. The attraction between neutral molecules is weaker than the attraction between charged ions. As a result, many covalent compounds...

electronegativity

Electronegativity and Polar Bonds

Electronegativity helps explain why some covalent bonds are nonpolar and others are polar covalent. It is a measure of how strongly an atom attracts shared electrons toward itself. That pull controls how electron density is distributed in a bond. Linus Pauling proposed the first electronegativity values by studying bond energies in heteronuclear molecules such as hydrogen fluoride. He expected the bond energy of HF to be about the average of the bond energies of H2 and F2, or 296 kJ/mol. The...

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How Bond Polarity Shapes Molecular Dipoles

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic. Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it...

lewis structures of molecular compounds and polyatomic ions

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Drawing Lewis Structures with Valence Electron Counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign. The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine...

resonance

Resonance Forms in Lewis Structures

Resonance forms in Lewis structures help explain molecules and ions that cannot be shown correctly with just one drawing. In nitrite ion, NO2−, the Lewis structure can be drawn in two ways. The N-O and N=O bonds switch positions in the two drawings, but the atoms stay in the same arrangement. If nitrite really had one single bond and one double bond, the two bond lengths should be different. A double bond is shorter and stronger than a single bond between the same atoms. Experiments show that...

formal charges

Lewis Structures with Formal Charge

Lewis structures can sometimes show more than one reasonable arrangement for a molecule or polyatomic ion. Formal charge helps chemists choose the best Lewis structure when that happens. It is a useful way to compare possible structures in chemistry. Formal charge is the hypothetical charge an atom would have if bonding electrons were shared evenly between the atoms. It can also be found by starting with the valence electrons of the neutral atom, subtracting the nonbonding electrons, and then...

exceptions to the octet rule

Octet Rule Exceptions in Lewis Structures

Some covalent molecules do not follow the octet rule in their Lewis structures. Their central atoms can fall into three groups: odd-electron molecules, electron-deficient molecules, and hypervalent molecules. These cases show why a Lewis structure may not always give every atom eight valence electrons. Odd-electron molecules have an odd number of valence electrons, so one electron remains unpaired. Molecules with an odd number of electrons are called radicals. Nitric oxide, NO, is one example,...

bond energies and bond lengths

How Bond Energy Predicts Reaction Heat

Bond energy helps explain why molecules stay together and how much heat a reaction may release or absorb. A covalent bond is stronger when more energy is needed to separate the atoms. The energy needed to break one specific covalent bond in one mole of gaseous molecules is called bond energy, or bond dissociation energy. For a diatomic molecule, bond energy is the standard enthalpy change for an endothermic reaction that breaks that bond. The H-H bond is a common example. Its bond energy is...

Electron sea model explains how metallic bonding holds metals together. In this model, most metal atoms do not have enough valence electrons to form ionic or covalent bonds in the usual way. Their valence electrons are held loosely because metals have low electronegativity, and their ionization energy is low, so electrons are removed easily. When a metal atom loses these outer electrons, it becomes a positively charged metal ion. The freed electrons form a delocalized electron cloud around...

key terms and concepts

Learn how the electron sea model explains metallic bonding, conductivity, malleability, luster, and why metals like mercury, copper, and iron differ.

Electron Sea Model in Metals | JoVE K12

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Pennsylvania State Academic Standards

Structure and Properties of Matter

3.2.9-12.B

3.2.9-12.N

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Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Bonding and compounds

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Chemical bonding

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Chemical bonding (ionic, covalent, metallic)

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HS.P1U1.2

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New Jersey Student Learning Standards for Science

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Motion and Stability: Forces and Interactions

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Matter and Its Interactions

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Florida Department of Education

SC.912.P.8

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Chemistry and Physics

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Phase Transitions: Sublimation and Deposition

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Heating and Cooling Curves

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Phase Diagrams

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PS.PSc.1.2

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Physical and Chemical Properties of Matter

The characteristics that enable us to distinguish one substance from another are called properties.

Physical Properties of Matter

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Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

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PS.PSc.1.6

Understand types, properties, and structure of matter.

Electron Sea Model in Metals

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