Comparison between Covalent and Ionic Compounds

Covalent and ionic compounds have distinct physical properties.


Key Takeaways

valence electronsoctet ruleelectronegativity

Two Classes of Compounds

The definition of a compound is a substance which contains two or more chemical elements together.The chemical structures of proteins are characterized by a fixed ratio of atoms held together by chemical bonds.We take a look at two major types of compounds based on the bond type that holds the atoms together: ionic and covalent.

Covalent Compounds

In covalent bonds, atoms share electrons.The majority of these bonds occur between nonmetals and elements with similar electronegativity. Two atoms with similar electronegativity will not exchange electrons from their outermost shell, rather they will share electrons and their valence electron shell will be filled.

Examples of compounds that contain only covalent bonds are methane (CH4), carbon monoxide (CO), and iodine monobromide (IBr).


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Covalent bonding between hydrogen atoms: Since each hydrogen atom has one electron, they are able to fill their outermost shells by sharing a pair of electrons through a covalent bond.


Ionic Compounds

The ionic bond forms when two atoms have different electronegativity levels.Consequently, the less electronegative atom loses an electron and the more electronegative atom gains that electron, giving rise to two ions.Ionic bonds are formed by electrostatic attraction between oppositely charged ions.

An ionic bond occurs when a nonmetal, which is an electron acceptor, and a metal, which is an electron donor, act together.The metals have few valence electrons, while the nonmetals have closer to eight. Therefore, the nonmetal, in order to satisfy the octet rule, accepts an electron from the metal.An ionic bond can have more than one electron being sent and received.

Some examples of compounds with ionic bonding include NaCl, KI, MgCl2.


NaF is formed by the transfer of an electron from a sodium atom to a fluorine atom, resulting in two oppositely charged ions: Na+ and F–.It is the attraction of the oppositely charged ions that forms an ionic bond between Na and F.


Effect on Physical Properties

.The following differences exist:


Single Covalent Bonds

Single covalent bonds are sigma bonds, which occur when one pair of electrons is shared between atoms.


Learning Objectives

Identify the four orbital types used in covalent bond formation


Key Takeaways

sigma bondcovalent bondatomic orbital

Hierarchical Structure of the Atom

The position and energy of the electrons in an atom is described in four hierarchical levels.Here are some of the possible values (letters) they can have:

Principal energy levels are made out of sublevels, which are in turn made out of orbitals, in which electrons are found.

Atomic Orbitals

The probability that an electron can be found around a nucleus of an atom is called an atomic orbital.In general, orbital shapes serve to describe regions in space where electrons are likely to exist.The “electron density” indicates how likely a region in space is to contain electrons.


These are the shapes of the first five atomic orbitals, in order: 1s, 2s, and the three 2p orbitals.There are blue- and orange-shaded regions in space where electrons (which belong to these orbitals) can be found.


Sigma Bonds

In covalent bonding, two atomic orbitals meet in close proximity and the electron densities overlap.Sigma bonds, which are composed of orbitals of each bonded atom, are the strongest type of covalent bond.Sigma bonds can occur regardless of atomic orbital type as long as the orbitals directly overlap between the nuclei of the atoms.


Sigma bonds are formed by atomic orbital overlaps that cause a sigma bond to form between two atoms when they occur.The overlap always occurs between the atom nuclei of the two bonded atoms.


When one pair of electrons are shared between the atoms in a compound or molecule, a single covalent bond is formed.Covalent bonds can be modeled by comparing two atoms with a single line.The diatomic hydrogen molecule, H2, occurs as H—H because of its single covalent bond between its two hydrogen atoms.


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Sigma bond in the hydrogen molecule: Higher intensity of the red color indicates a greater probability of the bonding electrons being localized between the nuclei.


Double and Triple Covalent Bonds

Double and triple bonds, comprised of sigma and pi bonds, increase the stability and restrict the geometry of a compound.


Learning Objectives

Describe the types of orbital overlap that occur in single, double, and triple bonds


Key Takeaways

bond strengthbond lengthorbital hybridizationatomic orbitals

Double and Triple Covalent Bonds

The covalent bond occurs when electrons are shared between atoms. .During multiple bond formation, only atoms that have to gain or lose at least two valence electrons will be able to participate.

Bonding Concepts

Hybridization

.Essentially, hybridization describes a bonding situation from the point of view of the individual atom.As a result of combining s and p orbitals, hybrid orbitals are formed.Observation of the bonded geometry of molecules agrees with the geometrical arrangement of the newly formed hybrid orbitals, which have the same energy.The hybrid orbitals are designated by spx, where s and p denote the orbitals used for mixing. The value of the superscript x varies from 1 to 3, depending on how many p orbitals are required to explain the observed bonding.


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Schematic of hybrid sp3 orbitals as they result in space.Summarizing the superscripts (1 for s, and 3 for p) will give you the total number of hybrid orbitals.A tetrahedron is created when four orbitals point in the same direction as the vertices.


Pi Bonds

Two atoms can form a pi bond if their unhybridized p orbitals overlap.There is no overlap between the nuclei of the atoms, which is the key difference between sigma and pi bonds.To form the bond efficiently, unhybridized p orbitals must be in the same plane.


.Unlike sigma bonds, the electron density corresponding to shared electrons does not lie along the internuclear axis.


Multiple bonds between atoms always consist of a sigma bond, with any additional bonds being of the π type.

Examples of Pi Bonds

.Two carbon atoms form a double bond comprised of a sigma bond and a π bond.


The molecule ethylene is an example of a simple molecule with a double bond between carbon atoms.Indications are given of the lengths and angles of the bonds (which represent the molecular geometry).


.In a plane with 120-degree angles, three sp2 orbitals are arranged.In carbon, as the orbitals of the carbon atom approach one another, a bond forms.A bond forms as the p orbitals converge on each other.Due to this bond, it is impossible to rotate the p atoms as they must remain parallel.

With a triple bond, three electrons share six electrons with two pi bonds and one sigma bond.As with acetylene, C2H2 is the simplest triple-bonded organic compound.Triple bonds are stronger than double bonds because there are two pi bonds compared to one.Two carbon atoms have sp hybrid orbitals, and one of them overlaps the corresponding sp orbital from the other carbon atom to form a sp-sp sigma bond.In the remaining four unhybridized p orbitals, two pi bonds overlap with each other.Triple bonds cannot be rotated around their axis, just like double bonds can't.

Observable Consequences of Multiple Bonds

Bond Strength

Bonds can be classified according to the amount of energy required to break them.Since O2 requires more energy to break a bond than H2 does to break a bond between two hydrogen atoms, the oxygen atoms must be held together more tightly.Two oxygen atoms have a stronger bond than two hydrogen atoms, so we say the bond between them is stronger.

.It would therefore take more energy to break the triple bond in N2 than to break the double bond in O2.The O2 molecule is broken by 497 kcal/mol while the N2 molecule is broken by 945 kJ/mol.

Bond Length

Moreover, there is a difference between the distances between nuclei of bonded atoms when there is the presence of multiple bonds between them.A double bond is shorter than a single bond, and a triple bond is shorter than a double bond.


Physical Properties of Covalent Molecules

The covalent bonding model helps predict many of the physical properties of compounds.


Learning Objectives

Discuss the qualitative predictions of covalent bond theory on the boiling and melting points, bond length and strength, and conductivity of molecules


Key Takeaways

bond lengthintermolecular forcesbond strengthoctet rule

Gilbert Lewis first described a covalent bond as when electrons of two different atoms are shared.Using the octet rule, we can anticipate these cases of electron sharing.A rule of chemistry that applies to atoms with a low atomic number (2, and halogens (F2, Cl2, Br2, I2)) is the octet rule. .It is believed that the Lewis bond theory explains why these atoms will share their valence electrons and thus create their own octet.

Several physical properties of molecules/compounds are related to the presence of covalent bonds:

Some observations of compounds in nature are not explained by the Lewis theory of covalent bonding.According to the theory, two atoms will be more bonded if more electrons are shared.Triple bonds are stronger than double bonds, and double bonds are stronger than single bonds, according to theory.I concur.But according to the theory, the bond strength of double bonds is twice that of single bonds, which isn't really true.Even though the covalent bonding model explains many physical observations, it has its limitations.