In this explainer, we will learn how to describe metallic bonding and the effect it has on the physical and chemical properties of metals.
Almost everyone is familiar with metals because metals are used all over the world. Some metals are used to make electrical wires and other metals are reshaped into cans and decorative jewelry. Metals are also used to make automotive vehicles and to provide structural support for large buildings and monuments. Metals are a mainstay of modern living, and it would be difficult to imagine life without them.
Metals are made up of relatively immobile cations and highly mobile delocalized electrons. The immobile cations are arranged as a giant three-dimensional lattice, and the lattice is held together due to attractive interactions between the lattice and delocalized electrons. Metallic bonds are usually described as being nondirectional interactions because the delocalized electrons collectively act as a glue that affects all the metallic cations equally. Metals are almost always described as being dense because the lattice is made up of cations that are held very closely together.
Definition: Metallic Bonding
Metallic bonding is the strong electrostatic attraction that exists between positively charged metal cations and delocalized electrons.
Example 1: Identifying the Dominant Force of Attraction in a Giant Metallic Lattice
In metallic bonding, what is the dominant form of attraction between the lattice of positive ions and the sea of delocalized electrons?
Metallic bonding can be defined as the electrostatic attraction that exists between relatively immobile metal cations and mobile delocalized electrons. The cations have a positive electrostatic charge, and they are attracted to surrounding negatively charged delocalized electrons. The electrostatic attraction between the oppositely charged particles is the dominant form of attraction in a giant metallic lattice structure. These statements can be used to determine that option D is the correct answer for this question.
The following figure shows that delocalized electrons come from the valence shell of metal atoms. Valence electrons have very low ionization energies, and they are the easiest electrons to remove from neutrally charged atoms. The delocalized electrons are relatively unconstrained when they decouple from metal atoms, and they can move in between the cations that make up a metallic lattice. It is important to realize here that the number of delocalized electrons will almost always be exactly equal to the number of valence electrons in a pure metal element. Sodium atoms generate single electrons because they have one valence electron and magnesium atoms generate two because they have two valence electrons.
Example 2: Identifying the Best Description of Delocalized Electrons
Which of the following is the best description of delocalized electrons in metallic bonding?
- Core electrons bound to metal ions
- Core and valence electrons that can move freely between metal ions
- Valence electrons that can move freely between metal ions
- Valence electrons bound to metal ions
- Core electrons that can move freely between metal ions
Metal atoms have at least two shells of electrons, but the valence electrons are the only electrons that decouple from metal atoms in giant metallically bonded structures. The valence electrons group together, and they effectively form one giant sea of delocalized electrons. The delocalized electrons are relatively unconstrained, and they can move in between immobile metal cations.
We can determine that options A, B, and E are all incorrect because they state that delocalized electrons are either core electrons or a combination of core and valence electrons.
We can also determine that option D cannot be the correct answer because it states that delocalized electrons are particles that are bound to metal ions. Delocalized electrons are relatively unconstrained particles that can move throughout a giant metallic lattice.
We have determined that option C must be the correct answer through the process of elimination, but we can also use logical reasoning to validate this inference. We can deduce that option C is the correct answer because we have already stated that delocalized electrons are valence electrons that can move in between immobile metal cations. Option C must be the correct answer for this question because it similarly states that delocalized electrons are valence electrons that can move freely between metal ions.
The following table lists some of the most common physical properties of pure metal elements. The listed properties can all be explained with the sea of electrons model.
|The Common Physical Properties of Pure Metal Elements|
|High melting point|
|High boiling point|
|High electrical conductivity|
|High heat conductivity|
The melting point and boiling point of any metal will depend on the size of its positively charged cations. Metals usually have high melting points and boiling points if they contain small cations. Metals usually have lower melting points and boiling points if they contain large cations. The following table shows the melting point and atomic radius values of the sodium and lithium metal elements.
|Element Name||Lithium ()||Sodium ()|
|Electron Shells||(2, 1)||(2, 8, 1)|
|Atomic Radius (pm)||152||186|
|Melting Point ()||181||98|
Lithium has a higher melting point than sodium partly because lithium atoms are smaller and they can pack closer together. Lithium atoms become highly concentrated when they pack close together in a lithium lattice, and this generates strong metallic bonding forces. It takes lots of energy to overcome the strong forces of attraction between lithium cations and surrounding delocalized electrons. It takes much less energy to overcome the forces of attraction between larger sodium cations and surrounding delocalized electrons.
Melting points and boiling points are also affected by the electrostatic charge state of cations in a metallic lattice. Cations have strong interactions with electrons if they have a high charge state and weaker interactions if they have a lower charge state. Cations have relatively strong interactions with electrons if they have a or charge state and relatively weak interactions if they have a charge state. This statement is supported by data in the following table. The table proves that sodium has a lower melting point and boiling point than magnesium. It also proves that magnesium has a lower melting point and boiling point than aluminum.
|Element Name||Sodium ()||Magnesium ()||Aluminum ()|
|Electron Shells||(2, 8, 1)||(2, 8, 2)||(2, 8, 3)|
|Melting Point ()||98||650||659|
|Boiling Point ()||890||1 110||2 470|
Sodium ions have relatively weak interactions with surrounding delocalized electrons because they have a charge state. Magnesium ions have stronger interactions with delocalized electrons because they have a charge state. Aluminum ions have even stronger interactions with delocalized electrons because they have a charge state.
It is important to understand that the boiling points and melting points are also affected by the concentration of delocalized electrons in a metallic lattice. Aluminum has such an exceptionally high boiling point partly because it contains such a high concentration of delocalized electrons. Each aluminum atom generates three delocalized electrons, and each sodium and magnesium atom can only generate one or two delocalized electrons. The following figure shows that aluminum atoms generate more delocalized electrons than sodium atoms.
The previous paragraphs state that there are lots of different particle and lattice properties that affect the melting point and boiling point of a metal. The paragraphs have explained that metals tend to have relatively high melting points and boiling points if they contain small cations. Lithium metal is supposed to have a higher melting point than sodium because lithium ions are smaller than sodium ions. The previous paragraphs also explain that metals tend to have high melting points and boiling points if they contain a high concentration of delocalized electrons and cations that have a high charge state. Aluminum is supposed to have a much higher boiling point than sodium because aluminum contains a higher concentration of delocalized electrons and it is made up of higher charge state cations.
Example 3: Identifying the System with the Highest Melting Point
Which of the following atoms would form a solid metal with the highest melting point?
The question asks us to determine which atom would form a solid metal with the highest melting point. We can immediately discount system B because it has a full valence shell of electrons and it is a noble gas. Noble gas atoms tend to form gases that have incredibly low melting points. The other options need to be carefully compared because they are all metals, and we cannot be so quick to discount them as possible answers for this question.
Metals generally have high melting points when they have a high number of valence electrons. Metals will tend to have a relatively low melting point if they have a single valence electron, and they will tend to have a higher melting point if they have two or even three valence electrons. This suggests that systems A and C have higher melting points than systems D and E. Systems A and C both have two valence electrons, and systems D and E only have one valence electron. This line of reasoning has helped us to determine that the correct answer must be system A or system C. We now have to compare other atomic properties to determine if system A or C has the higher melting point.
Metals generally have high melting points if they are made up of small atoms, and they have lower melting points if they are made up of large atoms. System A has three electron shells and system C has two electron shells. Electron shells take up space, and this suggests that system C is smaller than system A. Our line of reasoning suggests that system C must form a solid metal that has a higher melting point than systems A and B or systems D and E. We can conclude that system C has to be the correct answer for this question.
Metals are good at conducting electricity because they contain delocalized electrons. The delocalized electrons can carry charge when a voltage is applied across the metal lattice. The delocalized electrons move toward the positive terminal, and the negative terminal produces more electrons. Metals tend to be better conductors of electricity when they have lots of delocalized electrons. Aluminum is a better electrical conductor than magnesium, and magnesium is a better electrical conductor than sodium.
Definition: Electrical Conductivity
Electrical conductivity is a measure of how easy it is for an electric charge to pass through a material.
Example 4: Understanding How Valence Electrons Affect Electrical Conductivity
Shown in the bar chart are the electrical conductivities of aluminum, magnesium, and sodium.
- Using the symbols of the elements, identify the metals ,
, and .
- , ,
- , ,
- , ,
- , ,
- , ,
- Which of the following atomic properties is most responsible for the variation in conductivity between
metals , , and ?
- Atomic mass
- Bond energy
- Ionization energy
- Number of core electrons
- Number of valence electrons
The sodium, magnesium, and aluminum elements all have the same number of core electrons, but they have a different number of valence electrons. Metals tend to have higher electrical conductivity values when they contain more charge-carrying valence electrons. Sodium atoms have one valence electron, and magnesium and aluminum atoms have two and three valence electrons. System must be sodium because it has the lowest electrical conductivity value. System must be aluminum because it has the highest electrical conductivity value. System must be magnesium because it has an electrical conductivity value that is in between these two extremes. These statements can be used to determine that option D is the correct answer for this question.
Metals generally have high electrical conductivity values when they contain lots of charge-carrying valence electrons. Aluminum has three valence electrons per atom, and it has the highest electrical conductivity value of the three listed options. Sodium has just one valence electron per atom, and it has the lowest electrical conductivity value of the three listed options. We can use these statements to determine that option E is the correct answer for this question.
Metals tend to be better heat conductors than nonmetals because metallic lattices contain delocalized electrons. Heat is conducted in nonmetal materials when some of the nonmetal atoms come into contact with a heat source. The heat source makes some of the nonmetal atoms vigorously vibrate around the same lattice point. This kinetic energy is then passed onto other nonmetal atoms as one of the vibrating atoms bumps into other adjacent nonmetal atoms. The kinetic energy is slowly passed through the nonmetal material as the nonmetal atoms vibrate and bump into each other.
Metal atoms can transfer heat through similar slow vibration and collision processes. Some metal cations vibrate as they are warmed up, and these cations eventually bump into other adjacent metal cations. Heat energy can move throughout the metallic lattice as kinetic energy is passed along a long chain of vibrating metal cations.
Metals also contain highly mobile delocalized electrons that can transfer heat very rapidly. The electrons are relatively unconstrained, and they can move in between metal cations. The delocalized electrons can slip through the metallic lattice and quickly transfer heat between different parts of the giant metallic structure. The following figure compares heat (red arrows) transfer processes in giant nonmetal and metal structures.
Pure metals are usually both ductile and malleable. Metals are described as being ductile because they can be drawn out into thin wires. Metals are described as being malleable because they can be hammered or pressed into different shapes without breaking or cracking. Metals are usually both ductile and malleable because they contain delocalized electrons.
Ductile materials can be drawn out into thin wires.
Giant metallic structures are made up of layers of metal cations that are held together by a sea of delocalized electrons. This makes it easy for one layer of metal cations to slide over another layer without the lattice having to break or reform any chemical bonds. The sea of delocalized electrons will change its shape as the layers of metal cations slide over each other.
Malleable materials can be hammered or pressed into shape without breaking or cracking.
Pure metal elements can almost always be described as being lustrous (shiny) when they exist as a large block or a medium-to-large sized stone. The lustrous appearance of a metal can be explained in terms of its delocalized electrons. It is said that the delocalized electrons at the surface of the metal are continuously absorbing and then reflecting packets of light energy. The reflection of light by a delocalized sea of electrons is what tends to make a metallic surface look so shiny or lustrous. Some metals tend to be highly valuable because they are lustrous. Gold tends to be sold for a high price partly because it is so lustrous and rare.
Lustrous materials can reflect light evenly and efficiently without glitter or sparkle.
It is not always desirable to have metals that are malleable or ductile and we sometimes want to reform pure metals into alloys to make them stiffer. We can produce alloys by mixing one pure metal with another metal or nonmetal element. The following figure shows how an alloy can be made by mixing one pure metal element with an altogether different substance. Alloys tend to have unusual physical properties because they have such an irregular arrangement of atoms and ions.
An alloy can be made by combining one metal element with one or more other elements.
Example 5: Identifying the Correct Classification for a Composite Material
A metal’s properties can be changed by mixing that metal with another element. Which of the following is the name given to the resulting mixture?
- A gel
- An alloy
- An emulsion
- A suspension
- A metalloid
The atoms of any one metal element can be combined with the atoms of a different element to make a new composite material. The new composite material usually has an altogether different set of physical properties. The new composite material might be less malleable, or it might be more hard and more brittle. The process of generating new composite materials from metal elements is now well established, and people have developed terms to describe the process of making new composite materials and terms for the composite materials themselves. The process is usually called alloying, and the end product of the alloying process is almost always called an alloy. We can use these statements to determine that option B must be the correct answer for this question.
- Metallic bonding is the strong electrostatic attraction that exists between relatively immobile cations and mobile delocalized electrons.
- Most metals have high melting and boiling points because of the strong electrostatic attraction that exists between their cations and delocalized electrons.
- Metals are generally good conductors of heat and electricity because they have delocalized electrons.
- Pure metals are usually malleable and ductile because they contain layers of atoms that can easily pass over each other.
- Pure metals are usually lustrous because they contain a sea of delocalized electrons.
- Pure metals can be mixed with other elements to produce metal alloys.