How To Draw Lewis Dot Diagrams For Ionic Compounds

Learning Objectives

• Land the octet rule.
• Define ionic bond.
• Draw Lewis structures for ionic compounds.

In Section 4.7, we demonstrated that ions are formed by losing electrons to make cations, or by gaining electrons to course anions. The astute reader may have noticed something: many of the ions that form have 8 electrons in their valence shell. Either atoms proceeds plenty electrons to accept eight electrons in the valence shell and become the appropriately charged anion, or they lose the electrons in their original valence shell; the lower shell, now the valence shell, has 8 electrons in information technology, and so the atom becomes positively charged. For whatever reason, having eight electrons in a valence beat is a especially energetically stable arrangement of electrons. Theoctet ruleexplains the favorable trend of atoms having viii electrons in their valence crush. When atoms class compounds, the octet dominion is not always satisfied for all atoms at all times, but it is a very good rule of pollex for understanding the kinds of bonding arrangements that atoms can make.

It is not incommunicable to violate the octet rule. Consider sodium: in its elemental form, it has i valence electron and is stable. It is rather reactive, however, and does non require a lot of energy to remove that electron to make the Na ⁺ ion. Nosotros could remove another electron by calculation fifty-fifty more free energy to the ion, to brand the Na ²⁺ ion. However, that requires much more free energy than is ordinarily bachelor in chemical reactions, and so sodium stops at a 1+ charge afterwards losing a single electron. It turns out that the Na ⁺ ion has a consummate octet in its new valence shell, the n = 2 beat out, which satisfies the octet dominion. The octet rule is a event of trends in energies and is useful in explaining why atoms grade the ions that they practise.

Now consider an Na atom in the presence of a Cl atom. The two atoms have these Lewis electron dot diagrams and electron configurations:

\[\mathbf{Na\, \cdot }\; \; \; \; \; \; \; \; \; \; \mathbf{\cdot }\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}\]

\[\left [ Ne \right ]3s^{1}\; \; \; \; \left [ Ne \right ]3s^{2}3p^{five}\]

For the Na cantlet to obtain an octet, it must lose an electron; for the Cl atom to gain an octet, it must proceeds an electron. An electron transfers from the Na atom to the Cl cantlet:

\[\mathbf{Na\, \cdot }\curvearrowright \mathbf{\cdot }\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}\]

resulting in ii ions—the Na ⁺ ion and the Cl ⁻ ion:

\[\mathbf{Na}^{+}\; \; \; \; \; \; \; \; \mathbf{:}\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}^{-}\]

\[\left [ Ne \right ]\; \; \; \; \; \left [ Ne \right ]3s^{two}3p^{6}\]

Both species at present have complete octets, and the electron shells are energetically stable. From basic physics, we know that opposite charges attract. This is what happens to the Na ⁺ and Cl ⁻ ions:

\[\mathbf{Na}^{+}\; + \; \mathbf{:}\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}^{-}\rightarrow Na^{+}Cl^{-}\; \; or\; \; NaCl\]

where we have written the final formula (the formula for sodium chloride) as per the convention for ionic compounds, without listing the charges explicitly. The allure between oppositely charged ions is called an ionic bail, and it is one of the main types of chemical bonds in chemistry. Ionic bonds are caused by electrons transferring from one atom to another.

In electron transfer, the number of electrons lost must equal the number of electrons gained. We saw this in the formation of NaCl. A similar process occurs between Mg atoms and O atoms, except in this case two electrons are transferred:

The ii ions each accept octets as their valence crush, and the 2 oppositely charged particles attract, making an ionic bond:

\[\mathbf{Mg\,}^{2+}\; + \; \left[\mathbf{:}\mathbf{\ddot{\underset{.\: .}O}}\mathbf{\: :}\right]^{2-}\; \; \; \; \; Mg^{2+}O^{2-}\; or\; MgO\]

Remember, in the final formula for the ionic compound, nosotros do not write the charges on the ions.

What about when an Na atom interacts with an O cantlet? The O atom needs 2 electrons to complete its valence octet, just the Na atom supplies only one electron:

\[\mathbf{Na\, \cdot }\curvearrowright \mathbf{\cdot }\mathbf{\ddot{\underset{.}O}}\mathbf{\: :}\]

The O atom still does non have an octet of electrons. What we demand is a 2d Na cantlet to donate a second electron to the O atom:

These three ions attract each other to give an overall neutral-charged ionic compound, which we write as Na _two O. The demand for the number of electrons lost beingness equal to the number of electrons gained explains why ionic compounds accept the ratio of cations to anions that they practise. This is required past the police of conservation of matter as well.

Example \(\PageIndex{one}\): Synthesis of Calcium Chloride from Elements

With arrows, illustrate the transfer of electrons to form calcium chloride from \(Ca\) atoms and \(Cl\) atoms.

Solution

A \(Ca\) cantlet has two valence electrons, while a \(Cl\) atom has 7 electrons. A \(Cl\) atom needs simply one more to complete its octet, while \(Ca\) atoms have two electrons to lose. Thus nosotros need two \(Cl\) atoms to accept the two electrons from one \(Ca\) atom. The transfer process looks as follows:

The oppositely charged ions attract each other to brand CaCl ₂ .

Exercise \(\PageIndex{i}\)

With arrows, illustrate the transfer of electrons to form potassium sulfide from \(Thou\) atoms and \(South\) atoms.

Answer

Summary

• The trend to form species that accept eight electrons in the valence shell is called the octet rule.
• The attraction of oppositely charged ions caused by electron transfer is called an ionic bond.
• The force of ionic bonding depends on the magnitude of the charges and the sizes of the ions.

Contributions & Attributions

This page was constructed from content via the following contributor(south) and edited (topically or extensively) by the LibreTexts development squad to meet platform fashion, presentation, and quality:

• Marisa Alviar-Agnew (Sacramento City College)

• Henry Agnew (UC Davis)

Source: https://chem.libretexts.org/Courses/College_of_Marin/CHEM_114:_Introductory_Chemistry/10:_Chemical_Bonding/10.03:_Lewis_Structures_of_Ionic_Compounds-_Electrons_Transferred

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