8.1 Ionic Bonding
PRE-HEALTH PREP
Ionic Bonding
Lattice Energy Trends and Bond Strength
General Chemistry 8.1
Ionic bonding involves the transfer of electrons between metals and non-metals, while covalent bonding involves electron sharing between non-metals. The electronegativity difference between bonded atoms determines whether a bond is classified as ionic (ΔEN > 1.7), polar covalent (0.5–1.7), or non-polar covalent (< 0.5), with greater differences corresponding to greater ionic character. Lattice energy — the measure of ionic bond strength defined as the energy required to separate an ionic solid into gaseous ions — depends primarily on ion charge and secondarily on ion size, with higher charges and smaller ions producing stronger electrostatic attractions and higher lattice energies.
Ionic Bonding vs. Covalent Bonding
At its simplest, ionic bonding involves the transfer of electrons from one atom to another, while covalent bonding involves the sharing of electrons between atoms. In reality, bonding exists along a continuous spectrum rather than falling neatly into two categories, but these working definitions provide a useful starting point for classifying bond type.
Consider the most famous ionic compound: NaCl (sodium chloride). Sodium has one valence electron, and chlorine has seven. Sodium loses its valence electron, effectively achieving a filled octet from its previous shell. Chlorine gains that electron, completing the octet of its current shell. This electron transfer gives sodium a positive charge and chlorine a negative charge, and the resulting electrostatic attraction between these oppositely charged ions holds the ionic compound together.
Example: Sodium Chloride Ionic Bond Formation
Na loses 1 electron → Na+ (cation)
Cl gains 1 electron → Cl− (anion)
The oppositely charged ions attract each other, forming the ionic compound NaCl.
In contrast, covalent bonding occurs when atoms share electrons rather than transferring them. The shared electrons form what we call covalent bonds, and neither atom completely gains or loses electrons in the process.
| IONIC vs COVALENT BONDING | Ionic | Covalent |
|---|---|---|
| Atoms involved | Between metal and nonmetal | Between 2 Nonmetals |
| Electron behavior | "Transfer" of electrons | "Sharing" of electrons |
| Ionic (ΔEN > 1.7) | Electronegativity Difference > 1.7 Larger Difference = Greater Ionic Character | Nonpolar — Electronegativity Difference < 0.5 Usually between 2 of the same atom or C–H |
| Polar covalent (0.5–1.7) | Polar — Electronegativity Difference 0.5 to 1.7; have partial ionic character |
Identifying Ionic vs. Covalent Bonds
The simplest method for identifying bond type relies on the identity of the atoms involved. A bond between a metal and a non-metal is generally classified as an ionic bond, while a bond between two non-metals is typically a covalent bond. Sodium is a metal and chlorine is a non-metal, so NaCl fits the ionic pattern perfectly.
Important Caveat: The metal/non-metal rule is not foolproof. Certain metal–non-metal combinations do not behave as purely ionic bonds. However, this guideline works well for the vast majority of cases you will encounter and is the rule you should follow on exams.
Electronegativity and Bond Type: 3 Classifications
Electronegativity provides a more precise way to classify ionic and covalent bonds. On the Pauling scale, fluorine holds the highest electronegativity at 4.0, followed by oxygen at 3.5. In general, the closer an element sits to fluorine on the periodic table, the more electronegative it is.
The key metric is the electronegativity difference (ΔEN) between the two bonded atoms. This single value determines which of the 3 bond type classifications applies:
| Electronegativity Difference (ΔEN) | Bond Type | Electron Behavior |
|---|---|---|
| Greater than 1.7 | Ionic | Essentially complete transfer |
| 0.5 to 1.7 | Polar covalent | Unequal sharing |
| Less than 0.5 | Non-polar covalent | Approximately equal sharing |
Example: Classifying NaCl by Electronegativity Difference
Sodium has an electronegativity of 0.9 and chlorine has an electronegativity of 3.0. The electronegativity difference is 3.0 − 0.9 = 2.1, which exceeds 1.7, confirming that NaCl is an ionic compound.
Exam Tip: Most courses do not require you to memorize Pauling electronegativity values. However, if you are provided a periodic table with electronegativity values, you may be asked to use the ΔEN cutoffs to classify bonds. Some sources use 1.8 instead of 1.7 as the ionic threshold — follow whichever value your instructor specifies.
Ionic Character and Periodic Trends
Because bonding exists on a continuous spectrum, chemists describe the degree to which a bond resembles a pure ionic bond using the term ionic character. The larger the electronegativity difference between two bonded atoms, the greater the ionic character of that bond.
Example: Comparing Ionic Character in NaCl and NaF
NaCl: ΔEN = 3.0 − 0.9 = 2.1
NaF: ΔEN = 4.0 − 0.9 = 3.1
Because sodium fluoride has a larger electronegativity difference, it has greater ionic character than sodium chloride.
Predicting Ionic Character from Periodic Trends
Even without a table of electronegativity values, you can compare ionic character using the periodic trend: electronegativity increases as you move up a group and to the right across a period. This makes it possible to rank compounds by ionic character using only your knowledge of periodic trends.
Example: Which Has the Greatest Ionic Character — LiF, NaF, KF, RbF, or CsF?
All five compounds share fluoride as the anion, so the comparison depends entirely on the cation. The greatest ionic character arises from the largest electronegativity difference, which means we need the least electronegative cation. Electronegativity decreases going down Group 1, making cesium the least electronegative alkali metal. Therefore, CsF has the greatest ionic character.
Non-Polar vs. Polar Covalent Bonds
Covalent bonds — those with a ΔEN below 1.7 — are further divided into two categories based on how equally the electrons are shared between the bonded atoms.
Non-polar covalent bonds (sometimes called “pure covalent” bonds) occur when the electronegativity difference is less than 0.5. In these bonds, the electrons are shared approximately equally between the two atoms, resulting in no separation of charge.
Polar covalent bonds occur when ΔEN falls between 0.5 and 1.7. The electrons are still shared, but one atom pulls them closer than the other, creating an unequal distribution of electron density. The more electronegative atom develops a partial negative charge (δ−), while the less electronegative atom develops a partial positive charge (δ+).
Recall: Electronegativity is the ability of an atom to attract shared electrons toward itself within a bond. When this pull is uneven, the bond becomes a polar covalent bond with measurable bond polarity.
2 Shortcuts for Recognizing Bond Polarity Without a Table
You do not always need exact electronegativity values to classify a covalent bond. Two reliable shortcuts make identification straightforward even without a reference table:
- Two identical atoms always form a non-polar covalent bond. For example, F₂ is non-polar because both atoms have the same electronegativity, resulting in zero electronegativity difference.
- Carbon–hydrogen bonds are considered non-polar covalent. Carbon has an electronegativity of 2.5 and hydrogen has 2.1, giving a ΔEN of only 0.4. Because C–H bonds appear so frequently in organic molecules and biomolecules, this is a classification you should commit to memory.
Beyond these two shortcuts, virtually any bond between two different non-metals will be polar covalent.
Example: Classifying a Carbon–Chlorine Bond
Carbon has an electronegativity of 2.5 and chlorine has an electronegativity of 3.0, giving a ΔEN = 0.5. This falls right at the boundary and is classified as a polar covalent bond.
Summary of Bond Identification: Metal + non-metal → ionic bond. Two non-metals → covalent bond (then determine whether non-polar or polar based on ΔEN or the two shortcuts above).
What Is Lattice Energy?
Lattice energy is a measure of ionic bond strength in an ionic compound. A higher lattice energy indicates stronger ionic bonds. This concept applies exclusively to ionic compounds — it is not used to describe covalent bond strength.
The Technical Definition of Lattice Energy
Lattice energy is defined as the energy required to completely separate one mole of an ionic solid into its gaseous ions. For sodium chloride, this process is represented as:
Lattice Energy of NaCl
NaCl(s) → Na+(g) + Cl−(g)
The energy change associated with this reaction is the lattice energy. Because you must overcome the strong electrostatic attraction between oppositely charged ions, this process is highly endothermic.
The stronger the ionic bonds holding the solid together, the more energy is required to pull the ions apart — and therefore the greater the lattice energy. Lattice energy trends directly reflect the strength of the electrostatic forces within an ionic compound.
Note: Lattice energies cannot be measured directly in the laboratory. Instead, they are determined indirectly using a thermodynamic method called the Born-Haber cycle, which you may encounter in a separate lesson.
Exam Tip: You may be asked to identify which chemical equation correctly represents the lattice energy of a given ionic compound. Remember: it always starts with the ionic solid and produces gaseous ions.
2 Factors That Affect Lattice Energy
Two factors determine the magnitude of lattice energy: the charge on the ions and the size of the ions. Mathematically, lattice energy is proportional to the product of the ion charges divided by the distance between the ion centers:
Lattice Energy Relationship
Lattice Energy ∝ (q+ × q−) / r
where q+ and q− are the magnitudes of the ion charges, and r is the distance between ion centers (which depends on ion size).
| LATTICE ENERGY TREND |
|---|
| 1. Higher for greater charges (more significant) |
| 2. Higher for smaller ions (less significant) |
Factor 1: Ion Charge (More Significant)
Greater charges on the cation and anion produce stronger electrostatic attraction and therefore higher lattice energy. A +2 cation is attracted more strongly to an anion than a +1 cation. Likewise, a −2 anion creates a stronger attraction than a −1 anion. Ion charge is typically the dominant factor when comparing lattice energies across different ionic compounds.
Factor 2: Ion Size (Less Significant)
Smaller ions can get closer together, shortening the distance between their nuclei. A shorter internuclear distance means a stronger electrostatic attraction and a higher lattice energy. Since lattice energy is inversely proportional to r, even small decreases in ion size can have a meaningful effect on ionic bond strength.
Strategy for Ranking Lattice Energies: Always examine ion charges first — charge is the more significant of the 2 factors. Only compare ion sizes when the charges are identical across all compounds in the set.
Lattice Energy Comparison Examples and Trends
The following two worked examples illustrate how to apply these 2 factors to rank lattice energies — a common question type on general chemistry exams.
When Ion Charges Differ
Example: Which Has the Highest Lattice Energy — LiF, NaF, MgO, or MgCl₂?
First, determine the ion charges for each ionic compound:
- LiF: Li+ and F− (charges: +1 and −1)
- NaF: Na+ and F− (charges: +1 and −1)
- MgO: Mg2+ and O2− (charges: +2 and −2)
- MgCl₂: Mg2+ and Cl− (charges: +2 and −1)
MgO has the highest charges on both ions (+2 and −2), so it wins on charge considerations alone. MgO has the highest lattice energy.
When Ion Charges Are the Same
Example: Which Has the Highest Lattice Energy — NaF, NaCl, NaBr, or NaI?
All four compounds contain Na+ paired with a halide ion carrying a −1 charge. Since the ion charges are identical, the lattice energy comparison comes down to ion size.
The halide ions increase in size going down the group: F− < Cl− < Br− < I−. Fluoride is the smallest halide, producing the shortest internuclear distance with sodium and therefore the strongest ionic bond.
NaF has the highest lattice energy because its smaller fluoride ion allows for a shorter distance between ion centers and a stronger electrostatic attraction.
Exam Strategy: On exams, instructors typically design lattice energy comparison questions so that one compound clearly wins on charge, or all charges are the same and you compare ion sizes. You rarely need to weigh both factors against each other simultaneously.
Key Takeaways
- Ionic bonds form through the transfer of electrons (typically metal to non-metal), while covalent bonds form through the sharing of electrons (typically between two non-metals).
- The electronegativity difference (ΔEN) provides a quantitative way to classify bonds: greater than 1.7 is ionic, 0.5–1.7 is polar covalent, and less than 0.5 is non-polar covalent.
- Ionic character increases with larger electronegativity differences — use periodic trends to predict which compound in a set has the greatest ionic character.
- Two identical atoms bonded together and carbon–hydrogen bonds are the most common examples of non-polar covalent bonds.
- Lattice energy measures ionic bond strength and is defined as the energy needed to separate one mole of an ionic solid into its gaseous ions.
- Lattice energy is proportional to ion charge (the dominant factor) and inversely proportional to ion size — higher charges and smaller ions both increase lattice energy and reflect stronger ionic bonds.
- When comparing lattice energies, always examine ion charges first; only compare ion sizes when all charges in the set are identical.