Lecture Notes:

Ions and Ionic Compounds

• Stability and the Octet Rule
  •  Recall that in discussing periodicity we found that the Noble Gases were very unreactive; this is a result of completely filled outermost electron energy levels
  • All of the Noble Gases have eight electrons in their outmost shell (except He which only has two electrons); this is the origin of the octet rule of stability
  • It is an observed fact that elements are more stable when they have eight electrons in their outermost energy level (remember that we are only talking about Main Group Elements: Groups 1, 2, and 13-18; the transition metals are another story altogether because their outermost energy levels are d orbitals, which hold more than 8 electrons)
  • Other Main Group elements gain or lose electrons to be ‘more like’ the Noble Gases; in doing so, they become more stable
  • In order to gain or lose electrons and remain stable, elements undergo chemical reactions; for this reason almost no elements except the noble gases are found uncombined in nature
  • This is the origin of reactivity
    • Take K (potassium) as an example: it has one more electron in its outermost energy levels (a 4s electron) than Ar (argon)
    • By losing the 4s electron, K can achieve the electron configuration of the noble gas Ar
    • The process of removing an electron is (you recall) called ionization and it requires an input of energy; the process is favorable because of the increased stability of the product
    • The driving force in the ionization process is the release of energy that comes about because of the formation of a chemical bond; the formation of bonds is an exothermic process
  • Chemicals react in ways that increase their stability; measuring stability is difficult
  • Instead of trying to measure stability, chemists measure the changes in energy that accompany a chemical reaction; energy is easier to measure
  • Remember that objects with a high potential energy tend to lose that energy: rocks fall off of the top of cliffs onto our heads, not the other way around
  • The same is true on the atomic level: when potassium loses an electron it releases chemical potential energy; if you try to reverse a reaction such as the reaction of potassium with water it requires that you add energy
  • For this reason, different compounds are thought of as being at different energy levels
    • It takes energy to break chemical bonds
    • Bonds that require a lot of energy to break are low energy bonds because they are more stable
    • Bonds that require a small amount of energy to break are high energy bonds because they are less stable
• Electron Transfers
  •  When you subtract an electron, an atom becomes positively charged; this is because the number of protons (the number of positive charges) remains the same while the number of negative charges is reduced; an atom or group of atoms with a positive charge is called a cation
  • When you add an electron, an atom becomes negatively charged; this is because the number of protons (the number of positive charges) remains the same while the number of negative charges is increased; an atom or group of atoms with a negative charge is called an anion
  • When ₁₁Na (sodium) loses an electron, it gains the stable octet of the element Ne (neon); it does not become neon because it still has 11 protons, not 10 (₁₀Ne); you write the symbol for this cation as Na⁺
  • When ₁₇Cl (chlorine) gains an electron, it gains the stable octet of Ar (argon); remember though that it still only has 17 protons! you write the symbol for this anion as Cl^-
  • Opposite charges attract; when Na reacts with Cl they exchange electrons and form an ionic bond
  • Ionic compounds do not exist as individual molecules because they are more stable when they form crystals
  • Crystalline shapes or arrangements of atoms allow (in the case of NaCl) 6 Cl^- anions to be next to each Na⁺ cation
  • Even though ionic compounds exist because of the attraction of positive for negative charges they are themselves electrically neutral; the charges must balance
• The Properties of Salts
  •  Solid salts are poor conductors of electricity
  • To conduct electricity a substance must have charged particles that can move around; solid salts’ charged particles cannot move at all except to vibrate in place
  • Molten and dissolved salts are excellent conductors of electricity because their ions can move toward the electrode that has the opposite charge
  • Differences in electronegativity between the ions in a salt play a large role in determining the properties of the salt; large differences in electronegativity mean stronger bonds than small differences in electronegativity
  • The stronger the bonds between ions in a salt, the higher the melting point and boiling point of that salt; ionic compounds are characterized by having very high melting points and boiling points
    • NaCl melts at 801°C and boils at 1413°C; water, which has a different kind of chemical bond, melts at 0°C and boils at 100°C
  • Salts are hard and brittle because it takes a lot of energy to push the ions apart and once they are pushed far enough apart that like charges end up next to one another they break abruptly
• Energy and Ionic Bonding
  •  The process of removing an electron (ionization) require the input of energy
  • Adding an electron to an atom (electron affinity) usually releases some energy but not usually enough to cause the ionization of another atom
  • Ionic compounds form as a result of a larger process than just the exchange of electrons
    • Energy must be added to the system to get things started; just like a spark is needed to start a fire even though the overall reaction releases energy
    • The formation of a crystal lattice ultimately releases more energy than was added to the system because it is so much more stable than the starting materials; the more energy is released in this process, the stronger the ionic bond
  • Metals exhibit crystal structure just like ionic compounds but the forces holding them together work differently; here are some points of contrast:
    • Electrons are not tied to the atoms but can move throughout the crystal structure
    • Because the electrons can move, metals are good conductors of electricity and heat
    • Also, because the electrons can move without disrupting the forces holding the atoms together, metals can be stretched or pounded flat without breaking