Structure 2.1: Ionic Bonding and Structure

Ions

• Neutral atoms possess an equal count of protons and electrons.
• Cations, or positive ions, are created via the loss of electrons by atoms.
• Anions, or negative ions, are created via the gain of electrons by atoms.

Positive Ions (Cations)

• Positive ions result from atoms shedding electrons.
• For instance, a sodium atom gives up its solitary valence electron to become a positive ion with a charge of 1+, denoted as $Na^+$.
• This process reduces the atom's occupied energy levels, giving it the same electron configuration as neon (Ne): $1s^2 2s^2 2p^6$.

Negative Ions (Anions)

• Negative ions are formed when atoms gain electrons.
• For example, a chlorine atom gains an electron to form a negative ion with a 1- charge, denoted as $Cl^-$.
• This results in a full outer shell of electrons with the same electron configuration as argon (Ar): $1s^2 2s^2 2p^6 3s^2 3p^6$.

Exercises: Ion Formation and Properties

1. Radius Comparison: Sodium Atom vs. Sodium Ion

   • The sodium atom has three occupied energy levels, whereas the sodium ion has only two.
   • The sodium ion experiences a stronger effective nuclear charge because it has more protons (11) than electrons (10), increasing the electrostatic attraction and shrinking the ion.

2. Radius Comparison: Chloride Ion vs. Chlorine Atom

   • The chloride ion, having gained an electron, has more electrons than protons.
   • The increased electron count intensifies electron repulsion, causing the ion to expand. The valence electrons experience a weaker electrostatic attraction, making the ion larger.

3. Ion Formation Tendencies: Metals vs. Non-metals

   • Metals, characterized by low ionization energy values, readily lose electrons to form positive ions.
   • Non-metals, conversely, have high ionization energy values, favoring electron gain and the formation of negative ions.

Ionic Bonding

• Ionic bonding arises from the electrostatic attraction between oppositely charged ions.
• Typically, it occurs when there's an electronegativity difference of 1.8 or greater between two elements (though exceptions exist).
• Ionic bonds usually form between metal and non-metal elements located on opposite sides of the periodic table.
• The formation of an ionic bond is often described as the transfer of electrons.

Exercises: Understanding Ionic Bonds

1. Describe an Ionic Bond

   • An ionic bond is the electrostatic attraction between oppositely charged ions.

2. Bond Formation Between Elements on Opposite Sides of the Periodic Table

   • Elements on the far left of the periodic table (metals) and those on the far right (non-metals) typically form ionic bonds.
   • The considerable electronegativity difference (1.8 or greater) between metals and non-metals facilitates the formation of ionic bonds.

Writing Formulae of Ionic Compounds

• Common Positive Ions (Cations)
  • Hydrogen: $H^+$
  • Sodium: $Na^+$
  • Silver: $Ag^+$
  • Potassium: $K^+$
  • Lithium: $Li^+$
  • Ammonium: $NH_4^+$
  • Barium: $Ba^{2+}$
  • Calcium: $Ca^{2+}$
  • Copper(II): $Cu^{2+}$
  • Magnesium: $Mg^{2+}$
  • Zinc: $Zn^{2+}$
  • Mercury(I): $Hg^{+}$
  • Lead: $Pb^{2+}$
  • Iron(II): $Fe^{2+}$
  • Iron(III): $Fe^{3+}$
  • Aluminum: $Al^{3+}$
• Common Negative Ions (Anions)
  • Fluoride: $F^-$
  • Chloride: $Cl^-$
  • Bromide: $Br^-$
  • Iodide: $I^-$
  • Hydrogencarbonate: $HCO_3^-$
  • Hydroxide: $OH^-$
  • Nitrate: $NO_3^-$
  • Oxide: $O^{2-}$
  • Sulfate: $SO_4^{2-}$
  • Carbonate: $CO_3^{2-}$
  • Phosphate: $PO_4^{3-}$
  • Nitride: $N^{3-}$
  • Sulfide: $S^{2-}$
  • Phosphide: $P^{3-}$
  • Nitrite: $NO_2^-$
  • Sulfite: $SO_3^{2-}$

Exercises: Formulae of Ionic Compounds

• Potassium bromide: KBr
• Calcium fluoride: $CaF_2$
• Beryllium sulfide: BeS
• Strontium iodide: $SrI_2$
• Magnesium nitride: $Mg<em>3N</em>2$
• Aluminum oxide: $Al<em>2O</em>3$
• Sodium carbonate: $Na<em>2CO</em>3$
• Copper(II) phosphide: $Cu<em>3P</em>2$
• Zinc phosphate: $Zn<em>3(PO</em>4)_2$
• Ammonium nitrate: $NH<em>4NO</em>3$
• Ammonium sulfate: $(NH<em>4)</em>2SO_4$
• Iron(III) sulfite: $Fe<em>2(SO</em>3)_3$
• Copper(II) nitrite: $Cu(NO<em>2)</em>2$
• Potassium hydrogencarbonate: $KHCO_3$
• Aluminum sulfate: $Al<em>2(SO</em>4)_3$
• Mercury(I) nitride: $Hg_3N$
• Iron(II) nitrite: $Fe(NO<em>2)</em>2$
• Barium nitrate: $Ba(NO<em>3)</em>2$
• Iron(II) phosphide: $Fe<em>3P</em>2$
• Calcium hydrogencarbonate: $Ca(HCO<em>3)</em>2$

Polyatomic Ions

• Polyatomic ions (or molecular ions) consist of two or more atoms covalently bonded together and carrying an overall charge.
• Atoms within a polyatomic ion are held together by covalent bonds, while the ions in a compound containing a polyatomic ion are bonded ionically.
• The geometry of a polyatomic ion depends on the number of electron domains around the central atom.
• Polyatomic ions with multiple positions for multiple bonds exhibit resonance structures, leading to equal bond lengths and strengths – intermediate between single and double bonds.

Exercise: Bonding Types in Ionic Compounds with Polyatomic Ions

• Ionic compounds containing polyatomic ions exhibit two types of bonding: covalent bonds within the polyatomic ion and ionic bonds between the polyatomic ion and the other constituent ion.

Structure of Ionic Compounds

• Ionic compounds form three-dimensional lattice structures represented by empirical formulas.
• Key properties include volatility, electrical conductivity, and solubility.

Properties of Ionic Compounds

• Electrical Conductivity
  • Ionic compounds do not conduct electricity in the solid-state because ions are locked in fixed positions due to strong electrostatic attractions.
  • They conduct electricity when molten (melted) or dissolved in water because the ions are free to move and carry electric current.
• Melting and Boiling Points
  • Ionic compounds have high melting and boiling points due to strong electrostatic attractions between oppositely charged ions (e.g., NaCl melts at 800°C).
  • Higher charge and smaller ionic radius increase electrostatic attraction and thus, the melting point.
• Solubility
  • Ionic compounds are soluble in polar solvents like water. Polar water molecules separate ions from the lattice structure and surround them (hydration).

Exercises: Structure and Properties of Ionic Compounds

1. Structure of Ionic Compounds
   • Ionic compounds form a lattice structure where oppositely charged ions are held by electrostatic forces.
2. Conductivity of Ionic Compounds
   • Solid ionic compounds do not conduct electricity because ions are fixed in the lattice structure.
   • Molten or dissolved ionic compounds conduct electricity because ions are free to move.
3. High Melting Point of Ionic Compounds
   • The high melting point is attributed to strong electrostatic attractions between oppositely charged ions.
4. Melting Point Comparison: NaF vs. KF
   • NaF has a higher melting point than KF because the sodium ion has a smaller ionic radius than the potassium ion, resulting in stronger electrostatic attraction.

Lattice Enthalpy ($\Delta H_{lattice}^{\ominus}$

• Lattice enthalpy ($\Delta H<em>{lattice}^{\ominus}$) is the enthalpy change when one mole of a solid ionic compound is separated into its gaseous ions under standard conditions. For example: $NaCl(s) \rightarrow Na^+(g) + Cl^-(g)$$\Delta H</em>{lattice}^{\ominus} = +790 \text{ kJ mol}^{-1}$
• Lattice enthalpy indicates ionic bond strength; higher values mean stronger bonds.
• Factors affecting lattice enthalpy are the charge on the ions and the ionic radius of the ions.
  • Higher ion charge leads to stronger attraction and higher lattice enthalpy.
  • Smaller ions also have stronger attraction and higher lattice enthalpy.

Exercises: Lattice Enthalpy Comparison

1. Lattice Enthalpy Comparison: $MgCl_2$ vs. NaCl

   • $MgCl_2$ has a higher lattice enthalpy due to the higher charge of the $Mg^{2+}$ ion compared to $Na^+$.

2. Lattice Enthalpy Comparison: NaBr vs. KBr

   • NaBr has a higher lattice enthalpy because the $Na^+$ ion has a smaller ionic radius than the $K^+$ ion.

3. Arranging Compounds by Magnitude of Lattice Enthalpy

   • Increasing order of lattice enthalpy: KCl < NaCl < MgS < MgO