2.2 Structure and Properties Study Guide

2.2 Structure and Properties (Chem.2.2)

Explore this Phenomenon

Imagine you have a hot drink that you want to sweeten with sugar. You can see two different white crystals. You know that one is table salt and one is sugar.

1. If you weren’t allowed to taste the crystals, how would you figure out which one was which?
2. As you read the following section, think of possible tests you could conduct (besides tasting) to determine which white crystal is table salt and which white crystal is sugar.

 

Standard Chem.2.2

Plan and carry out an investigation to compare the properties of substances at the bulk scale and relate them to molecular structures . Emphasize using models to explain or describe the strength of electrical forces between particles. Examples of models could include Lewis dot structures or ball and stick models. Examples of particles could include ions, atoms, molecules, or networked materials (such as graphite). Examples of properties could include melting point and boiling point, vapor pressure, solubility, or surface tension. (PS1.A)

In this section, see if you can recognize how the structures of ionic substances and the structures of covalent substances could explain the properties and function of the substance.

 

Properties of Salt, NaCl

Sodium chloride, NaCl, is an ionic compound. It is made of one positively charged Sodium atom and one positively charged Chlorine atom. In the following sections, you will read about some of the other properties of ionic compounds.

 

Physical Properties of Ionic Compounds

In nature, the ordered arrangement of ionic solids gives rise to beautiful crystals. (A) Amethyst – a form of quartz, SiO 2 , whose purple color comes from iron ions. (B) Cinnabar – the primary ore of mercury is mercury(II) sulfide, HgS. (C) Azurite – a copper mineral, Cu 3 (CO 3 ) 2 (OH) 2 . D) Vanadinite – the primary ore of vanadium, Pb 5 (VO 4 ) 3 Cl.

What produces colored crystals?

The figure above shows just a few examples of the color and brilliance of naturally occurring ionic crystals. The regular and orderly arrangement of ions in the crystal lattice is responsible for the various shapes of these crystals, while transition metal ions give rise to the colors.

 

Physical Properties of Ionic Compounds

Melting Points

Because of the many simultaneous attractions between cations and anions that occur, ionic crystal lattices are very strong. The process of melting an ionic compound requires the addition of large amounts of energy in order to break all of the ionic bonds in the crystal. For example, sodium chloride has a melting point of about 800°C.

 Shattering

Ionic compounds are generally hard, but brittle. Why? It takes a large amount of mechanical force, such as striking a crystal with a hammer, to force one layer of ions to shift relative to its neighbor. However, when that happens, it brings ions of the same charge next to each other (see Figure below ). The repulsive forces between the like-charged ions cause the crystal to shatter. When an ionic crystal breaks, it tends to do so along smooth planes because of the regular arrangement of the ions.

(A) The sodium chloride crystal is shown in two dimensions. (B) When struck by a hammer, the negatively-charged chloride ions are forced near each other and the repulsive force causes the crystal to shatter.

 

Conductivity

Another characteristic property of ionic compounds is their electrical conductivity. The figure below shows three experiments in which two electrodes that are connected to a light bulb are placed in beakers containing three different substances.

(A) Distilled water does not conduct electricity. (B) A solid ionic compound also does not conduct. (C) A water solution of an ionic compound conducts electricity well.

In the first beaker, distilled water does not conduct a current because water is a molecular compound . In the second beaker, solid sodium chloride also does not conduct a current. Despite being ionic and thus composed of charged particles, the solid crystal lattice does not allow the ions to move between the electrodes. Mobile charged particles are required for the circuit to be complete and the light bulb to light up. In the third beaker, the NaCl has been dissolved into the distilled water. Now the crystal lattice has been broken apart and the individual positive and negative ions can move. Cations move to one electrode, while anions move to the other, allowing electricity to flow (see Figure below ). Melting an ionic compound also frees the ions to conduct a current . Ionic compounds conduct an electric current when melted or dissolved in water .

In an ionic solution, the A + ions migrate toward the negative electrode, while the B – ions migrate toward the positive electrode.

 

What Are Nonmetals?

Nonmetals are elements that generally do not conduct electricity. They are one of three classes of elements (the other two classes are metals and metalloids .) Nonmetals are the second largest of the three classes after metals. They are the elements located on the right side of the periodic table.

Q: From left to right across each period (row) of the periodic table, each element has atoms with one more proton and one more electron than the element before it. How might this be related to the properties of nonmetals?

A: Because nonmetals are on the right side of the periodic table, they have more electrons in their outer energy level than elements on the left side or in the middle of the periodic table. The number of electrons in the outer energy level of an atom determines many of its properties.

 

Properties of Nonmetals

As their name suggests, nonmetals generally have properties that are very different from the properties of metals . Properties of nonmetals include a relatively low boiling point, which explains why many of them are gases at room temperature . However, some nonmetals are solids at room temperature, including the three pictured above, and one nonmetal—bromine—is a liquid at room temperature. Other properties of nonmetals are illustrated and described in the Figure below .

 

Reactivity of Nonmetals

Reactivity is how likely an element is to react chemically with other elements. Some nonmetals are extremely reactive, whereas others are completely nonreactive. What explains this variation in nonmetals? The answer is their number of valence electrons. These are the electrons in the outer energy level of an atom that are involved in interactions with other atoms. Let’s look at two examples of nonmetals, fluorine and neon. Simple atomic models of these two elements are shown in the Figure below .

Q: Which element , fluorine or neon, do you predict is more reactive?

A: Fluorine is more reactive than neon. That’s because it has seven of eight possible electrons in its outer energy level , whereas neon already has eight electrons in this energy level.

Although neon has just one more electron than fluorine in its outer energy level, that one electron makes a huge difference. Fluorine needs one more electron to fill its outer energy level in order to have the most stable arrangement of electrons. Therefore, fluorine readily accepts an electron from any element that is equally “eager” to give one up, such as the metal lithium or sodium. As a result, fluorine is highly reactive. In fact, reactions with fluorine are often explosive. Neon, on the other hand, already has a full outer energy level. It is already very stable and never reacts with other elements. It neither accepts nor gives up electrons. Neon doesn’t even react with fluorine, which reacts with all other elements except helium.

 

Why Most Nonmetals Cannot Conduct Electricity

Like most other nonmetals, fluorine cannot conduct electricity, and its electrons explain this as well. An electric current is a flow of electrons. Elements that readily give up electrons (the metals ) can carry an electric current because their electrons can flow freely. Elements that gain electrons instead of giving them up cannot carry electric current. They hold onto their electrons so they cannot flow.

The burner on a gas stove burns with a pretty blue flame like the one pictured in the opening image. The fuel burned by most gas stoves is natural gas, which consists mainly of methane. Methane is a compound that contains only carbon and hydrogen. Like many other compounds that consist of just these two elements, methane is used for fuel because it burns very easily. Methane is an example of a covalent compound.

 

What Are Covalent Compounds?

Compounds that form from two or more nonmetallic elements, such as carbon and hydrogen, are called covalent compounds. In a covalent compound , atoms of different elements are held together in molecules by covalent bonds. These are chemical bonds in which atoms share valence electrons . The force of attraction between the shared electrons and the positive nuclei of both atoms holds the atoms together in the molecule. A molecule is the smallest particle of a covalent compound that still has the properties of the compound.

The largest, most complex covalent molecules have thousands of atoms. Examples include proteins and carbohydrates , which are compounds in living things. The smallest, simplest covalent compounds have molecules with just two atoms. An example is hydrogen chloride (HCl). It consists of one hydrogen atom and one chlorine atom, as you can see in the Figure below .

 

Properties of Covalent Compounds

The covalent bonds of covalent compounds are responsible for many of the properties of the compounds. Because valence electrons are shared in covalent compounds, rather than transferred between atoms as they are in ionic compounds, covalent compounds have very different properties than ionic compounds.

• Many covalent compounds, especially those containing carbon and hydrogen, burn easily. In contrast, many ionic compounds do not burn.
• Many covalent compounds do not dissolve in water , whereas most ionic compounds dissolve well in water.
• Unlike ionic compounds, covalent compounds do not have freely moving electrons, so they cannot conduct electricity.
• The individual molecules of covalent compounds are more easily separated than the ions in a crystal, so most covalent compounds have relatively low boiling points. This explains why many of them are liquids or gases at room temperature . You can compare the boiling points of some covalent and ionic compounds in the Table below.

 

Name of Compound (Chemical Formula)	Type of Compound	Boiling Point (°C)
Methane (CH 4 )	covalent	-164
Nitrogen oxide (NO)	covalent	-152
Sodium chloride (NaCl)	Ionic	1413
Lithium fluoride (LiF)	ionic	1676

 

Q: The two covalent compounds in the table are gases at room temperature , which is 20°C. For a compound to be a liquid at room temperature, what does its boiling point have to be?

A: To be a liquid at room temperature , a covalent compound has to have a boiling point higher than 20°C. Water is an example of a covalent compound that is a liquid at room temperature. The boiling point of water is 100°C.

 

Summary

• Covalent compounds contain two or more nonmetallic elements held together by covalent bonds, in which atoms share pairs of valence electrons . A molecule is the smallest particle of a covalent compound that still has the properties of the compound.
• Covalent bonds are responsible for many of the properties of covalent compounds. Covalent compounds have relatively low boiling points, cannot conduct electricity, and may not dissolve in water .

 

Putting It Together

Let us revisit the phenomenon: We now know that table salt is an ionic compound and sugar is a covalent compound.

1. Using your knowledge of the properties of ionic and covalent compounds design an experiment on how could you test if the white crystal is table salt or sugar.