4.3 Ionic bonding and electronegativity

Ionic bonding and electronegativity are key concepts in understanding chemical bonds. They explain how atoms interact to form stable compounds through electron transfer and attraction. These principles are crucial for predicting compound formation and properties.

Electronegativity differences determine bond types, from ionic to polar covalent. This knowledge helps us predict how elements will combine and behave in various chemical reactions, forming the foundation for understanding more complex molecular interactions.

Ionic bond formation

Complete transfer of valence electrons

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• Ionic bonds form when valence electrons are completely transferred from a metal to a nonmetal
  • Results in the formation of positively charged cations (metals) and negatively charged anions (nonmetals)
  • Example: In sodium chloride (NaCl), the sodium atom loses its valence electron to the chlorine atom, forming Na+ and Cl- ions

Electrostatic attraction between ions

• The primary force holding ionic compounds together is the electrostatic attraction between oppositely charged ions
  • Cations and anions are strongly attracted to each other due to their opposite charges
  • This attraction is the basis for the stability of ionic compounds

Properties of ionic compounds

• High melting and boiling points
  • Strong electrostatic forces between ions in the solid state lead to high melting and boiling points
  • Example: Sodium chloride (NaCl) has a melting point of 801°C and a boiling point of 1,413°C
• Brittleness
  • Ionic compounds are typically brittle and can shatter upon impact
  • This is due to the rigid arrangement of ions in the crystal lattice, which does not allow for easy deformation
• Electrical conductivity
  • In the solid state, ionic compounds do not conduct electricity due to the fixed positions of the ions
  • When molten or dissolved in water, ionic compounds become excellent conductors
    • Mobile ions can carry electrical charge
    • Example: Molten sodium chloride (NaCl) is an excellent conductor of electricity

Electronegativity and bond polarity

Electronegativity

• Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond
  • Atoms with higher electronegativity have a greater tendency to attract electrons
  • Example: Fluorine (F) has the highest electronegativity of all elements
• Electronegativity determines bond type and polarity
  • The difference in electronegativity between two atoms determines the type of bond formed
    • Large difference (> 1.7): Ionic bond
    • Small difference (< 0.4): Nonpolar covalent bond
    • Difference between 0.4 and 1.7: Polar covalent bond

Polar covalent bonds

• In a polar covalent bond, the electron density is unevenly distributed between the bonded atoms
  • The more electronegative atom has a higher electron density and a partial negative charge (δ-)
  • The less electronegative atom has a lower electron density and a partial positive charge (δ+)
  • Example: In a water molecule (H2O), oxygen is more electronegative than hydrogen, resulting in a polar covalent bond
• The degree of polarity in a bond depends on the difference in electronegativity between the bonded atoms
  • Larger differences result in more polar bonds
  • The polarity of a bond can be represented by a dipole moment, which is a vector quantity

Predicting ionic compound formation

Periodic trends in electronegativity

• Electronegativity generally increases from left to right across a period in the periodic table
  • Effective nuclear charge increases and atomic radius decreases, leading to a greater attraction for electrons
  • Example: In period 2, electronegativity increases from lithium (Li) to fluorine (F)
• Electronegativity generally decreases from top to bottom within a group in the periodic table
  • Atomic radius increases and effective nuclear charge decreases, leading to a weaker attraction for electrons
  • Example: In group 1 (alkali metals), electronegativity decreases from lithium (Li) to francium (Fr)

Formation of ionic compounds

• Ionic compounds are most likely to form between elements with a large difference in electronegativity
  • Typically between metals from the left side of the periodic table (low electronegativity) and nonmetals from the right side (high electronegativity)
  • Example: Sodium chloride (NaCl) forms between sodium (Na), a metal, and chlorine (Cl), a nonmetal
• The charge of the ions formed in an ionic compound can be predicted based on the electron configuration of the elements involved
  • Elements tend to achieve a stable octet configuration by losing or gaining electrons
  • Example: Sodium (Na) forms a +1 cation (Na+) by losing its single valence electron, while chlorine (Cl) forms a -1 anion (Cl-) by gaining one electron

Crystal structures and lattice energies

Crystal lattices

• Ionic compounds form crystal lattices, which are regular, repeating arrangements of cations and anions in three-dimensional space
  • The most common crystal structures for ionic compounds are:
    • Sodium chloride (NaCl) structure
    • Cesium chloride (CsCl) structure
    • Zinc blende (ZnS) structure
• The type of crystal structure adopted by an ionic compound depends on several factors
  • Relative sizes of the cations and anions
  • Charge ratio between the ions
  • Coordination number of the ions (the number of nearest neighbors)

Lattice energy

• Lattice energy is the energy required to separate a mole of an ionic solid into its constituent gaseous ions
  • It is a measure of the stability of an ionic compound
  • Higher lattice energies indicate stronger electrostatic forces between the ions and greater stability
• Factors influencing lattice energy
  • Charge of the ions: Higher charges lead to higher lattice energies
  • Size of the ions: Smaller ions result in higher lattice energies due to closer packing
  • Crystal structure: Different crystal structures have different lattice energies
• Properties related to lattice energy
  • Ionic compounds with higher lattice energies tend to have:
    • Higher melting points
    • Higher boiling points
    • Greater hardness
  • These properties are a result of the stronger electrostatic forces holding the ions together
• Born-Haber cycle
  • The Born-Haber cycle is used to calculate the lattice energy of an ionic compound
  • It considers the various energy changes involved in the formation of the compound from its constituent elements
  • Example: The Born-Haber cycle for sodium chloride (NaCl) includes the sublimation of sodium, ionization of sodium, dissociation of chlorine, and formation of the ionic solid