ghschemistry Ms. Vervoort
Glastonbury High School  
Chapter 13, Ionic Bonds
The purpose of this chapter is to explain why certain atoms of certain elements combine or react to each other the way they do.  It will further serve to answer some questions regarding chemical and physical properties of such elements and compounds.

13.1 Valence Electrons
Valence electrons determine the chemical properties of an element.  First, a reminder of chemical properties: the ability of a substance to undergo a chemical reaction and form a new substance (use words like rust, decompose, decay, and combustion are used to describe chemical properties of an element).  Second, the definition of valence electrons: the electrons in the outermost orbit or energy level of the atom.

The periodic table is organized with valence electrons in mind.  For example, all the elements in group one have one valence electron in their highest energy level.  For the group A elements, the number of the group will tell you how many valence electrons the atom contains.  The only exception to the rule is found in the Noble gas group.  He has 2 valence electrons and the rest of the noble gases have 8.  They are non-reactive. 

Do problem 1 on page 300. 

We focus on valence electrons because we believe they are the only ones involved with chemical reactions.  In order to describe how atoms combine, it is customary to learn how to draw them.  We use electron dot structures to illustrate the atom’s valence electron composition.  View Table 13.1 on page 300 for examples.

13.2 Stable Electron Configurations for Cations
As far as atoms are concerned, stability is directly dependent on low energy.  The lower the energy level of the atom, the more stable it is.  High-energy atoms are unstable and tend to react more vigorously with other atoms of other elements.  They do this with the intent to achieve the lowest possible energy level, thus the greatest stability.  In chemistry, low energy and high stability are achieved when the atom has 8 valence electrons (like the noble gases). 

Gilbert Lewis, in 1916, proposed the octet rule.  He said atoms of different elements react with each other by changing the number of their valence electrons so as to acquire a stable electron structure.  This means a stable atom must have 8 electrons in its highest energy level, like the noble gases.

Lets look at the formation of cations (an atom that is positively charged due to a loss of electrons).

Atoms of most metals lose electrons when they react with other atoms because that is the easiest way for them to obey the octet rule and reach maximum stability and the lowest level of energy.  For example, Na with one valence electron will lose it when it reacts with other atoms and become a cation.  With only 10 electrons, Na+ looks like Ne in terms of electron configurations.  Look at the top of page 302 for an illustration of this point. 

Group 1A elements always have a charge of 1+ and group 2A have a charge of 2+.  This information was used in chapter 5 when you were asked to write formulas for ionic compounds and to criss-cross the charges.  Now you know why those cations have those charges.

As we have seen before, there are exceptions to every rule.  The octet rule has its exceptions (Ag+, Cu+, Au+, Cd++, and Hg++) but we will focus on the elements that obey the rule.  FYI – rather than losing valence electrons, these atoms lose electrons from lower energy levels in order to reach greater stability.

13.3 Stable Electron Configurations for Anions (atoms that gain electrons and are negatively charged)
Most non-metals will gain valence electrons in a chemical reaction to obey the octet rule and achieve stability.  Look at page 303 for a visual representation of how the Cl atom reacts with other atoms.  As you can see, it will gain one electron to obey the octet rule.  In fact, all group 7A atoms will gain one electron and become anions with –1 charge.  Cl- appears to have the electron configuration of Ar, a noble gas.  When group 7A atoms (the halogens) gain electrons they are called halide ions. 

Do problems 3 and 4 on page 304.   Writing a formula for an ion is as simple as Cl- or Na+

13.4 Ionic Compounds

Attractive forces between oppositely charged ions form ionic bonds.  Once two ions are joined together the resulting compound is neutral.  Remember from chapter 5, that we criss-cross the charges of the ions to form the compound’s formula.  Ionic compounds are also called salts.

Two examples of ionic compounds or salts are NaCl and Al Br3.  Look on the top of page 305 for an illustration of the bonding of these ions.  Notice that Na will give up one of its valence electrons to Cl and Al will give up 3 of its valence electrons to Br.

Do problems 5,6, and 7.

13.5 Properties of Ionic Compounds

At room temperature, they are solids and crystalline in composition (a repeating 3-D pattern). 

Related to the crystalline nature, is the coordination number.  This number tells you how many ions of opposite charge surround each ion in a crystal.  For example, in NaCl, the coordination number is 6.  This means there are 6 Cl- ions surrounding each Na+ ion and 6 Na+ ions surrounding each Cl- ion.  This pattern repeats and is 3-D, which forms the salt crystal.  Look at the bottom of page 307 for an illustration of this.

CsCl (cesium chloride) has a coordination number of 8.  Eight Cs+ ions surround each Cl- ion and 8 Cl- ions surround each Cs+ ion.  The pattern repeats and is 3-D.

Not all ions in a salt compound will have the same coordination number.  Titanium dioxide, for example, has a coordination number of 6 for Ti4+, and the coordination number of 3 for O2-.  This means each O2- ion is surrounded by 6 Ti4+ ions and each Ti4+is surrounded by 3 O2- ions.

When heated to the molten state, ionic compounds will conduct electricity because their ions are freed form their ionic bonds and they will move about freely.  When oppositely charged electrodes are introduced to the molten salt, the ions migrate towards the oppositely charged electrode and this causes current.  Refer to the diagram on page 308.

Breaking ionic bonds is caused by the application of a force, which causes like ions to slide beside each other.  At that moment, the like charges repel and the crystal breaks.

13.6 Metallic Bonds

Metals contain only one kind of atom, like Fe, Ni, or Cu.

The bonds between atoms of metals are unlike the bonds between ionic compounds.  In ionic compounds the attraction is between two oppositely charged ions.  In metals the attraction is between positive ions and free floating valence electrons.  Because of the free-floating valence electrons, metals are good conductors of electricity (because the valence electrons migrate from one end of the metal to the other.

Metals are good conductors of electricity, malleable (forced into different shapes), and ductile (drawn into wires). 

Breaking metallic bonds is not easy because the application of force only causes the ions to slide past each other but there is no repulsion of like charges because the negative to positive attractions remain intact.   Refer to the bottom of page 309.

There are 3 different crystalline structures for metals, the body-centered cubic, face-centered cubic, and hexagonal close-packed arrangements.  These can be found on the bottom of page 310.

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