Name: ionic compounds and crystal lattices
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 48 s
Channel: JoVE

K12 Standards

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North Carolina K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.1

Ionic Bonding and Electron Transfer

Ionic bonds are formed by the transfer of electrons from metals to nonmetals. In comparison to covalent bonds where electrons are shared, in ionic bonds metals tend to lose and nonmetals tend to accept electrons &mdash; but why?&nbsp;

The most stable electron configuration of an atom is with a full octet. Achieving this state in conjunction with lowering the potential energies is a major driving force for bond formation.&nbsp;

But how to predict when electrons will be transferred and not shared? For this, examine the ionization energies and electron affinities of atoms.

On the periodic table, s-block metals have low ionization energies. This facilitates easier loss of electrons to reach an octet while concurrently forming cations.&nbsp;

On the contrary, nonmetals from the p-block, except for noble gases, have high electron affinities and readily accept electrons to form anions. The oppositely charged cations and anions experience strong electrostatic interactions, attracting each other to form ionic bonds.

Ionic compounds often form hard and well-defined crystal structures with high melting points, due to strong electrostatic attraction between ions.

Consider lithium fluoride, an ionic compound made of lithium and fluorine. When lithium loses an electron to become a lithium cation with a duet, the electron is being taken up by fluorine to form the anion with an octet.

According to Coulomb&#39;s law, these two charged particles are attracted to each other to form lithium fluoride, a neutral compound.&nbsp;

The Lewis model can be used to describe and predict the chemical formula of ionic compounds. The Lewis symbol shows charges in the upper-right hand corner and depicts anions and cations with the symbol written within brackets.&nbsp;

Consider strontium chloride, a commonly used salt to get a bright red color in fireworks. The Lewis symbol for strontium has two, and chlorine has one unpaired electron.

Strontium must lose two electrons, while chlorine must gain one electron to reach an octet. Therefore, one strontium cation binds to two chloride anions, in the ratio of 1:2.&nbsp;

ionic bonding and electron transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence shell, and an anion (a negative ion) forms when a neutral atom gains one or more electrons in its valence shell. Compounds composed of ions are called ionic compounds (or salts), and their constituent ions are held together by ionic bonds: electrostatic forces of attraction between oppositely charged cations and anions.&nbsp;

<h4>Properties of Ionic Compounds</h4>

The properties of ionic compounds shed some light on the nature of ionic bonds.

<ul>
	<li>Ionic solids exhibit a crystalline structure and tend to be rigid and brittle; they also tend to have high melting and boiling points, which suggests that ionic bonds are very strong.&nbsp;</li>
	<li>Ionic solids are also poor conductors of electricity for the same reason &mdash; the strength of ionic bonds prevents ions from moving freely in the solid-state.&nbsp;</li>
	<li>Most ionic solids, however, dissolve readily in water. Once dissolved or melted, ionic compounds are excellent conductors of electricity and heat because the ions can move about freely.</li>
</ul>

<h4>The Formation of Ionic Compounds</h4>

Many metallic elements have relatively low ionization potentials and lose electrons easily. These elements lie to the left in a period or near the bottom of a group on the periodic table. Nonmetal atoms have relatively high electron affinities and thus readily gain electrons lost by metal atoms, thereby filling their valence shells. Nonmetallic elements are found in the upper-right corner of the periodic table.

As all substances must be electrically neutral, the total number of positive charges on the cations of an ionic compound must equal the total number of negative charges on its anions. The formula of an ionic compound represents the simplest ratio of the numbers of ions necessary to give identical numbers of positive and negative charges.&nbsp;

<h4>Ionic Compounds Form Regularly Arranged 3-Dimensional Structures</h4>

It is important to note, however, that the formula for an ionic compound does not represent the physical arrangement of its ions. It is incorrect to refer to sodium chloride (NaCl) &ldquo;molecule&rdquo; because there is not a single ionic bond, per se, between any specific pair of sodium and chloride ions. The attractive forces between ions are isotropic &mdash; the same in all directions &mdash; meaning that any particular ion is equally attracted to all of the nearby ions of opposite charge. This results in the ions arranging themselves into a tightly bound, three-dimensional lattice structure. Sodium chloride, for example, consists of a regular arrangement of equal numbers of Na+ cations and Cl&ndash; anions. The strong electrostatic attraction between Na+ and Cl&ndash; ions holds them tightly together in solid NaCl. It requires 769 kJ of energy to dissociate one mole of solid NaCl into separate gaseous Na+ and Cl&ndash; ions.&nbsp;

<h4>Electronic Structures of Cations</h4>

When forming a cation, an atom of the main group element tends to lose all of its valence electrons, thus assuming the electronic structure of the noble gas that precedes it in the periodic table.&nbsp;

<ul>
	<li>For groups 1 (the alkali metals) and 2 (the alkaline earth metals), the group numbers are equal to the numbers of valence shell electrons and, consequently, to the charges of the cations formed from atoms of these elements when all valence shell electrons are removed.&nbsp;</li>
	<li>For example, calcium is a group 2 element whose neutral atoms have 20 electrons and a ground-state electron configuration of 1s22s22p63s23p64s2. When a Ca atom loses both of its valence electrons, the result is a cation with 18 electrons, a 2+ charge, and an electron configuration of 1s22s22p63s23p6. The Ca2+ ion is, therefore, isoelectronic with the noble gas Ar.</li>
	<li>For groups 13&ndash;17, the group numbers exceed the number of valence electrons by 10 (accounting for the possibility of full d subshells in atoms of elements in the fourth and greater periods). Thus, the charge of a cation formed by the loss of all valence electrons is equal to the group number minus 10. For example, aluminum (in group 13) forms 3+ ions (Al3+).</li>
</ul>

<h4>Exceptions</h4>

<ul>
	<li>Exceptions to the expected behavior involve elements toward the bottom of the groups.&nbsp;</li>
	<li>In addition to the expected ions Tl3+, Sn4+, Pb4+, and Bi5+, a partial loss of these atoms&rsquo; valence shell electrons can also lead to the formation of Tl+, Sn2+, Pb2+, and Bi3+ ions. The formation of these 1+, 2+, and 3+ cations is ascribed to the inert pair effect, which reflects the relatively low energy of the valence s-electron pair for atoms of the heavy elements of groups 13, 14, and 15.&nbsp;</li>
	<li>Mercury (group 12) also exhibits an unexpected behavior: it forms a diatomic ion, Hg22+ (an ion formed from two mercury atoms, with a Hg-Hg bond), in addition to the expected monatomic ion Hg2+ (formed from only one mercury atom).</li>
	<li>Transition and inner transition metal elements behave differently than the main group elements. Most transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s electron(s) first, sometimes followed by the loss of one or two d electrons from the next-to-outermost shell.</li>
	<li>Although the d orbitals of the transition elements are&mdash;according to the Aufbau principle&mdash;the last to fill when building up electron configurations, the outermost s electrons are the first to be lost when these atoms ionize. When the inner transition metals form ions, they usually have a 3+ charge, resulting from the loss of their outermost s electrons and a d or f electron.</li>
</ul>

<h4>Electronic Structures of Anions</h4>

Most monatomic anions form when a neutral nonmetal atom gains enough electrons to completely fill its outer s and p orbitals, thereby reaching the electron configuration of the next noble gas. Thus, it is simple to determine the charge on such a negative ion: The charge is equal to the number of electrons that must be gained to fill the s and p orbitals of the parent atom. Oxygen, for example, has the electron configuration 1s22s22p4, whereas the oxygen anion has the electron configuration of the noble gas neon (Ne), 1s22s22p6. The two additional electrons required to fill the valence orbitals give the oxide ion the charge of 2&ndash; (O2&ndash;).

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-1-ionic-bonding">Openstax, Chemistry 2e, Section 7.3: Ionic Bonding.</a>

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

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Chemical Bonding: Basic Concepts

KEY TERMS AND CONCEPTS

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 melting and boiling points. They are also poor conductors of electricity in the solid state because the ions cannot move freely. Most ionic solids dissolve readily in water. When an ionic compound is melted or dissolved, it conducts electricity and heat well because the ions can move about freely.

Ionic compounds form most easily when metals lose electrons and nonmetals gain them. Many metals have low ionization potentials, so they lose electrons easily. These elements are found on the left side of the periodic table or near the bottom of a group. Nonmetals have high electron affinities, so they readily gain the electrons lost by metals. They are found in the upper-right corner of the periodic table.

Ionic compounds must remain electrically neutral. That means the total positive charge from the cations must equal the total negative charge from the anions. The formula of an ionic compound gives the simplest ratio of ions needed to balance charge. The formula does not show the physical layout of the ions. In solid sodium chloride, for example, the ions form a tightly bound three-dimensional lattice. The attraction between Na+ and Cl- ions is strong, and it takes 769 kJ of energy to separate one mole of solid NaCl into gaseous ions.

Main-group cations often form by losing all valence electrons and taking on the electron structure of the noble gas before them. For groups 1 and 2, the group number matches the number of valence electrons and also the charge of the ion formed. Calcium is a group 2 element. A neutral calcium atom has 20 electrons, and Ca2+ has 18 electrons with the same electron configuration as argon. For groups 13 through 17, the cation charge equals the group number minus 10. Aluminum, in group 13, forms Al3+.

Some heavy main-group elements do not follow the expected pattern exactly. Thallium, tin, lead, and bismuth can also form lower-charge ions such as Tl+, Sn2+, Pb2+, and Bi3+. This is explained by the inert pair effect, which reflects the relatively low energy of the valence s-electron pair in these heavy atoms. Mercury is another exception. It can form the diatomic ion Hg2^2+, which contains a Hg-Hg bond, as well as the monatomic ion Hg2+.

Transition metals and inner transition metals form ions in a different way. Most transition metal cations have 2+ or 3+ charges. These usually come from losing the outermost s electron or electrons first, and sometimes one or two d electrons after that. Even though d orbitals fill later according to the Aufbau principle, the outermost s electrons are lost first when transition metals ionize. Inner transition metals usually form 3+ ions by losing their outermost s electrons and one d or f electron.

Most monatomic anions form when a nonmetal atom gains enough electrons to fill its outer s and p orbitals. The charge on the ion equals the number of electrons needed to reach the next noble-gas electron configuration. Oxygen is a clear example. A neutral oxygen atom has the configuration 1s2 2s2 2p4. The oxide ion, O2-, has the neon configuration, 1s2 2s2 2p6, after gaining two electrons.

Ionic Compounds and Crystal Lattices

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 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

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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...

bond polarity, dipole moment, and percent ionic character

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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...

bonding in metals

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...

key terms and concepts

Explore how ions form, why ionic compounds make crystal lattices, and how charge, conductivity, and electron configurations shape ionic behavior in chemistry.

Ionic Compounds and Crystal Lattices | JoVE K12

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Physical Science

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Pennsylvania State Academic Standards

Structure and Properties of Matter

3.2.9-12.B

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Louisiana Student Standards for Science

High School

Matter And Its Interactions

HS-PS1-3

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Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Bonding and compounds

Structure And Properties Of Matter

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Matter and Its Interactions

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New Jersey Student Learning Standards for Science

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SC.912.P.8.6

Texas Essential Knowledge and Skills

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New York State Science Learning Standards

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Chemical Bonds

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Bonding Types: From Nonpolar Covalent to Ionic

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Ionic Compounds and Crystal Lattices

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