2.1 Bonding Study Guide

CHAPTER 2

Strand 2: Structure and Properties of Molecules

Chapter Outline

2.1 Bonding (Chem.2.1)

2.2 Structure and Properties (Chem.2.2)

2.3 Macromolecules (Chem.2.3)

2.4 Synthetic Chemistry (Chem.2.4)

Electrical attractions and repulsions between charged particles (atomic nuclei and electrons) in matter explain the structure of atoms and the forces between atoms that cause them to form molecules via chemical bonds. Molecules can range in size from two atoms to thousands of atoms. The same forces cause atoms to combine to form extended structures, such as crystals or metals. The varied properties of the materials, both natural and manufactured, can be understood in terms of the atomic and molecular particles present and the forces within and between them. Materials are engineered to fulfill a desired function or role with desired properties.

 

2.1 Bonding (Chem.2.1)

Explore this Phenomenon

Look at this periodic table showing electronegativity values below the atomic symbol.

1. Which of the main group of elements is missing?
2. What would you expect to happen if you combine calcium and phosphorus? Would you expect it to be different if you combine carbon and phosphorus?
3. What patterns can you find? How does it change going across? How does it change going down the table?

 

Standard Chem.2.1

Analyze data to predict the type of bonding most likely to occur between two elements using the patterns of reactivity on the periodic table. Emphasize the types and strengths of attractions between charged particles in ionic, covalent, and metallic bonds. Examples could include the attraction between electrons on one atom and the nucleus of another atom in a covalent bond or between ions in an ionic compound. (PS1.A, PS2.B)

In this section, look for patterns of reactivity to explain why certain elements form certain types of bonds. How can you use the patterns you have already learned about the Periodic Table to predict why elements will form compounds and which types of bonds are likely to be formed?

There is an amazing diversity of matter in the universe, but there are only about 100 elements. How can this relatively small number of pure substances make up all kinds of matter? Elements can combine in many different ways. When they do, they form new substances called compounds.

 

Chemical Bonding

Elements form compounds when they combine chemically. Their atoms join together to form molecules, crystals, or other structures. The atoms are held together by chemical bonds (a force of attraction between atoms or ions). Chemical bonds occur when atoms share or transfer valence electrons (the electrons in the outer energy level of an atom).

Water (H 2 O) is an example of a molecule. Water molecules always consist of two atoms of hydrogen and one atom of oxygen. Like water, all other chemical compounds consist of a fixed ratio of elements. It doesn’t matter how much or how little of a molecule/compound there is, this ratio remains constant.

It is important to know that when atoms combine to form compounds, their chemical properties – (an atom’s potential to change through a chemical reaction) and physical properties – (observable, measurable properties) change.

For example, table salt, called sodium chloride, is formed by bonding a sodium atom (Na) to a chlorine atom (Cl). Elemental sodium is a soft metal that reacts with water to produce a flammable gas. Elemental chlorine is a gas at room temperature and is poisonous. These two nasty chemicals join together to form table salt (NaCl); a substance we eat most every day. Sodium (Na) and chlorine (Cl) do not have the same properties as sodium chloride (NaCl). Elements do not have the same chemical and physical properties as when they join to form compounds.

Chemical Formulas

Elements are represented by chemical symbols. Examples are H for hydrogen and O for oxygen. Compounds are represented by chemical formulas – (indicates the types and number of atoms in a chemical compound). You’ve already seen the chemical formula for water; it’s H 2 O. The subscript 2 after the H shows that there are two atoms of hydrogen in a molecule of water. The O for oxygen has no subscript. When there is just one atom of an element in a molecule, no subscript is used. The table below shows some other examples of compounds and their chemical formulas.

 

Name of Compound	Numbers of Atoms	Chemical Formula
Hydrogen chloride	H = 1 Cl = 1	HCl
Methane	C = 1 H = 4	CH 4
Hydrogen peroxide	H = 2 O = 2	H 2 O 2
Carbon dioxide	C = 1 O = 2	CO 2

 

Did you ever play the card game called Go Fish? Players try to form groups of cards of the same value, such as four sevens, with the cards they are dealt or by getting cards from other players or the deck. This give and take of cards is a simple analogy for the way atoms give and take valence electrons in chemical reactions.

 

What Are Valence Electrons?

To understand chemical bonding, we first must understand valence electrons – (the electrons in the outer energy level of an atom). Valence electrons can participate in interactions with other atoms. Valence electrons are generally the electrons that are farthest from the nucleus. As a result, they may be attracted as much or more by the nucleus of another atom than they are by their own nucleus.

 

Electron Dot Diagrams

Because valence electrons are so important, atoms are often represented by simple diagrams that show only their valence electrons. These are called electron dot diagrams, and two are shown below.

In this type of diagram, an element's chemical symbol is surrounded by dots that represent the valence electrons. Typically, the dots are drawn as if there is a square surrounding the element symbol with up to two dots per side. Eight electrons complete an element’s valence shell, so we draw up to eight dots per atom.

 

Valence Electrons and the Periodic Table

The number of valence electrons in an atom is reflected by its position in the periodic table of the elements (see the periodic table below). Across each row, or period, of the periodic table, the number of valence electrons in groups 1–2 and 13–18 increases by one from one element to the next. Within each column, or group, of the table, all the elements have the same number of valence electrons. This explains why all the elements in the same group have very similar chemical properties.

For elements in groups 1–2 and 13–18, the number of valence electrons is easy to tell directly from the periodic table. This is illustrated in the simplified periodic table in the figure below. It shows just the numbers of valence electrons in each of these groups.

For elements in groups 3–12, determining the number of valence electrons is more complicated and goes beyond the scope of this course.

Valence Electrons and Reactivity

The number of valence electrons of an atom will determine its reactivity with other atoms. Atoms are more stable when they have full outer shells. In other words, stable atoms have full octet (8 valence electrons). The noble gas elements are the most stable, and therefore the least reactive of all the elements on the periodic table, because they already have eight valence electrons. Their valence shells are full. Atoms will bond or react with other atoms to become stable with full outer shells, or full octets. Some atoms will be more reactive than other atoms depending on their valence electrons. For example, fluorine is highly reactive because it has seven valence electrons and only needs one more electron to for a full octet, which makes the atom more stable. While most atoms are most stable with eight valence electrons, hydrogen is one exception because it only needs two valence electrons to fill its valence shell.

 

Atoms Are Neutral

An atom always has the same number of electrons as protons. Electrons have an electric charge of -1 and protons have an electric charge of +1. Therefore, the charges of an atom’s electrons and protons “cancel out”. This explains why atoms are neutral in electric charge.

What would happen to an atom’s charge if it were to gain extra electrons? It would have more electrons than protons. This would give it a negative charge, so it would no longer be neutral.

 

Atoms to Ions

Atoms cannot only gain extra electrons. They can also lose electrons. In either case, they become ions – (atoms that have a charge because they have gained or lost electrons). Ions are atoms that have a positive or negative charge because they have unequal numbers of protons and electrons. Atoms will gain or lose electrons in this process, but the amount of protons will stay the same. If atoms lose electrons, they become cations – (positive ions). If atoms gain electrons, they become anions – (negative ions). Consider the example of fluorine (see Figure below). A fluorine atom has nine protons and nine electrons and is electrically neutral. If a fluorine atom gains an electron, it becomes a fluoride ion with an electric charge of -1.

 

Metals and Nonmetals

       

As mentioned above, atoms lose or gain electrons to become stable.

Which atoms gain electrons and which atoms lose electrons? Metals, the atoms found on the left side of the table, tend to lose electrons and become cations; while nonmetals tend to gain electrons and become anions. Noble gases do not form ions.

Atoms form ions by losing or gaining electrons because it makes them more stable. The most stable state for an atom is to have its outermost energy level filled. In the case of metals such as lithium, with just one valence electron in the outermost energy level, a more stable state can be achieved by losing that one outer electron. In the case of nonmetals such as fluorine, which has seven valence electrons in the outermost energy level, a more stable state can be achieved by gaining one electron and filling up the outer energy level.

 

Some Common Ions

All the metals in family 1A (shown in the figure below) have one valence electron. The entire family forms +1 ions: Li + , Na + , K + , Rb + , Cs + , and Fr + . Note that although hydrogen (H) is in this same column, it is not considered to be a metal. There are times when hydrogen acts like a metal and forms +1 ions, but most of the time it bonds with other atoms as a nonmetal. In other words, hydrogen doesn’t easily fit into any chemical family.

The metals in family 2A all have two valence electrons. This entire family will form +2: Be +2 , Mg +2 , Ca +2 , Sr +2 , Ba +2 , Ra +2 .

The elements in Family 3A each have three valence electrons. When these atoms form ions, they will almost always form +3 ions: Al +3 , Ga +3 , In +3 , Ti +3 . Notice that boron is omitted from this list. This is because boron falls on the nonmetal side of the metal/nonmetal dividing line. Boron generally doesn’t lose all of its valence electrons during chemical reactions.

Family 4A is almost evenly divided into metals and nonmetals. The larger atoms in the family (germanium, tin, and lead) are metals. Since these atoms have 4 valence electrons, they are expected to form ions with charges of +4. All three of the atoms do form such ions (Ge +4 , Sn +4 and Pb +4 ), but tin and lead also have the ability to also form +2 ions.

Like family 4A, the elements of family 5A are also divided into metals and nonmetals. The smaller atoms in this family behave as nonmetals and form -3 ions, and the larger atoms behave as metals that form +5 ions. For the nonmetals, they each have 5 valence electrons so they will need to gain 3 more hence, the -3 charge.

Most of the elements in family 6A (shown in figure below) are nonmetals that have 6 valence electrons. They form -2 ions. If you consider that each has six valence electrons, they will need to gain two more to become stable: O -2 , S -2 , Se -2 and Te -2 .

Family 7A are all nonmetals. When these atoms form ions, they form -1 ions: F - , Cl - , Br - and I - . They each have seven valence electrons therefore; they need to gain one more to be stable.

Family 8A, of course, is made up of the noble gases, which have no tendency to either gain or lose electrons.

 

Types of Chemical Bonds - Depending on where elements lie on the periodic table, it will form one of three types of bonds - ionic, covalent, or metallic.

 

• An ionic bond is the force of attraction that holds together oppositely charged ions. Ionic bonds form crystals instead of molecules.
• A metallic bond is the force of attraction between a positive metal ion and the valence electrons that surround it—both its own valence electrons and those of other ions of the same metal. The ions and electrons form a lattice-like structure. Only metals form metallic bonds.
• A covalent bond is the force of attraction that holds together two nonmetal atoms that share a pair of electrons. One electron is provided by each atom, and the pair of electrons is attracted to the positive nuclei of both atoms.

 

Formation of Ionic Bonds

All compounds form when atoms of different elements share or transfer electrons. In the strongest type of bonds called ionic compounds, the electrons actually move from one atom to another. When atoms transfer electrons in this way, they become charged particles called ions. The ions are held together by ionic bonds.

An ionic bond (the force of attraction that holds together positive and negative ions) forms when atoms of a metallic element give up electrons to atoms of a nonmetallic element. The figure shows how this happens.

An ionic bond forms when the metal sodium gives up an electron to the nonmetal chlorine.

By losing an electron, the sodium atom becomes a sodium ion. It now has one less electron than protons, giving it a charge of +1. Positive ions such as sodium are given the same name as the element. The chemical symbol has a plus sign to distinguish the ion from an atom of the element. The symbol for a sodium ion is Na+.

By gaining an electron, the chlorine atom becomes a chloride ion. It now has one more electron than protons, giving it a charge of -1. Negative ions are named  by adding the suffix -ide to the first part of the element name. The symbol for chloride is Cl-.

Sodium and chloride ions have opposite charges. Opposites attract, so sodium and chloride ions attract each other. They cling together in a strong ionic bond. You can see this in row 2 of Figure above. Brackets separate the ions in the diagram to show that the ions in the compound do not share electrons.

 

Why Ionic Bonds Form

Ionic bonds form only between metals and nonmetals. Metals become more stable by giving up electrons, and nonmetals become more stable by gaining electrons. Find sodium (Na) on the periodic table - Sodium is an alkali metal in group 1. Like other group 1 elements, it has just one valence electron. If sodium loses that one electron, it will have a full outer energy level. Now find chlorine (Cl) on the periodic table – Chlorine is a halogen in group 17. It has seven valence electrons. If chlorine gains one electron, it will have a full outer energy level. After sodium gives up its valence electron to chlorine, both atoms have a more stable arrangement of electrons.

Properties of Ionic Compounds

           

Two models of a sodium chloride crystal are shown. The purple spheres represent the Na + ions, while the green spheres represent the Cl − ions. (A) In an expanded view, the distances between ions are exaggerated, more easily showing the coordination numbers of each ion. (B) In a space filling model, the electron clouds of the ions are in contact with each other.

Ionic compounds contain ions of metals and nonmetals held together by ionic bonds. Ionic compounds do not form molecules. Instead, many positive and negative ions bond together to form a structure called a crystal. You can see an example of a crystal in the figure. It shows the ionic compound sodium chloride. Positive sodium ions (Na + ) alternate with negative chloride ions (Cl - ). The oppositely charged ions are strongly attracted to each other.

 

Writing Basic Ionic Formulas

In writing formulas for binary ionic compounds (binary refers to two elements, not two single atoms), the cation is always written first. Chemists use subscripts following the symbol of each element to indicate the number of that element present in the formula. For example, the formula Na 2 O indicates that the compound contains two atoms of sodium for every one atom of oxygen. When the subscript for an element is 1, the subscript is omitted. The number of atoms of an element with no indicated subscript is always read as 1. When an ionic compound forms, the number of electrons given off by the cations must be exactly the same as the number of electrons taken on by the anions. Therefore, if calcium, which gives off two electrons, is to be combined with fluorine, which takes on one electron, then one calcium atom must combine with two fluorine atoms. The formula would be CaF 2 .

Suppose we wish to write the formula for the compound that forms between aluminum and chlorine. To write the formula, we must first determine the charges of the ions that would be formed.

Then, we determine the simplest whole numbers with which to multiply these charges so they will balance (add to zero). In this case, we would multiply the 3+ by 1 and the 1- by 3.

You should note that we could multiply the 3+ by 2 and the 1- by 6 to get 6+ and 6- respectively. These values will also balance, but this is not acceptable because empirical formulas, by definition, must have the lowest whole number multipliers. Once we have the lowest whole number multipliers, those multipliers become the subscripts for the symbols. The formula for this compound would be AlCl 3 .

Here’s the process for writing the formula for the compound formed between aluminum and sulfur.

Al 3+ S 2−

Therefore, the formula for this compound would be Al 2 S 3 .

Another method used to write formulas is called the crisscross method. It is a quick method, but it often produces errors if the user doesn’t pay attention to the results. The example below demonstrates the crisscross method for writing the formula of a compound formed from aluminum and oxygen. In the crisscross method, the oxidation numbers are placed over the symbols for the elements just as before.

In this method, the oxidation numbers are then crisscrossed and used as the subscripts for the other atom (ignoring sign).

This produces the correct formula Al 2 O 3 for the compound. Here’s an example of a crisscross error: If you used the original method of finding the lowest multipliers to balance the charges, you would get the correct formula PbO 2 , but the crisscross method produces the incorrect formula Pb 2 O 4 . If you use the crisscross method to generate an ionic formula, it is essential that you check to make sure that the subscripts correspond to the lowest whole number ratio of the atoms involved.

 

Polyatomic Ions

Polyatomic ions (charged particles made up of more than one atom) can also be present in ionic compounds. They are a group of bonded atoms with a charge that act like a single ion. Here is a short list of some common polyatomic ions:

• Ammonium ion, NH 4⁺
• Acetate ion, C 2 H 3 O 2^-
• Carbonate ion, CO 2^-
• Chromate ion, CrO 4^2-
• Dichromate ion, Cr 2 O 7^2-
• Hydroxide ion, OH ^-
• Nitrate ion, NO 3^-
• Phosphate ion, PO 4^3-
• Sulfate ion, SO 4^2-
• Sulfite ion, SO 3^2-

 

Suppose we are asked to write the formula for the compound that would form between calcium and the nitrate ion. We begin by putting the charges above the symbols just as before.

The multipliers needed to balance these ions are 1 for calcium and 2 for nitrate. We wish to write a formula that tells our readers that there are two nitrate ions in the formula for every calcium ion. When we put the subscript 2 beside the nitrate ion in the same fashion as before, we get something strange – CaNO32 . With this formula, we are indicating 32 oxygen atoms, which is wrong. The solution to this problem is to put parentheses around the nitrate ion before the subscript is added. Therefore, the correct formula is Ca (NO 3 ) 2 .

Similarly, calcium phosphate would be Ca 3 (PO 4 ) 2 . If a polyatomic ion does not need a subscript other than an omitted 1, then the parentheses are not needed. Although including these unnecessary parentheses does not change the meaning of the formula, it may cause the reader to wonder whether a subscript was left off by mistake. Try to avoid using parentheses when they are not needed.

 

Example 1

Write the formula for the compound that will form from aluminum and acetate.

The charge on an aluminum ion is +3, and the charge on an acetate ion is -1. Therefore, three acetate ions are required to combine with one aluminum ion. This is also apparent by the crisscross method. However, we cannot place a subscript of 3 beside the oxygen subscript of 2 without inserting parentheses first. Therefore, the formula will be Al(C 2 H 3 O 2 ) 3 .

 

Example 2

Write the formula for the compound that will form from ammonium and phosphate.

The charge on an ammonium ion is +1 and the charge on a phosphate ion is -3. Therefore, three ammonium ions are required to combine with one phosphate ion. The crisscross procedure will place a subscript of 3 next to the subscript 4. This can only be carried out if the ammonium ion is first placed in parentheses. Therefore, the proper formula is (NH 4 ) 3 PO 4 .

 

Variable Charge Metals

In general, main group metal ions have only one common charge, whereas most of the transition metals have more than one. However, there are plenty of exceptions to this guideline. There are some metals with variable charges where they may form multiple different compounds with the same nonmetal. Iron, for example, may react with oxygen to form either FeO or Fe 2 O 3 . These are very different compounds with different properties.

 

The Metallic Bond

Pure metals are crystalline solids, but unlike ionic compounds, every point in the crystal lattice is occupied by an identical atom . The electrons in the outer energy levels of a metal are mobile and capable of drifting from one metal atom to another. This means that the metal is more properly viewed as an array of positive ions surrounded by a sea of mobile valence electrons . A metallic bond is the attraction of the stationary metal cations to the surrounding mobile electrons.

In a metal, the stationary metal cations are surrounded by a sea of mobile valence electrons that are not associated with any one cation.

 

Properties of Metals

The metallic bonding model explains the physical properties of metals. Metals conduct electricity and heat very well because of their free-flowing electrons. As electrons enter one end of a piece of metal, an equal number of electrons flow outward from the other end. Recall that ionic compounds are very brittle. Application of a force results in like-charged ions in the crystal coming too close to one another, causing the crystal to shatter. When a force is applied to a metal, the free-flowing electrons can slip in between the stationary cations and prevent them from coming in contact. As a result, metals are very malleable and ductile. They can be hammered into shapes, rolled into thin sheets, or pulled into thin wires.

 

Covalent Bonds

In a tennis match, two players keep hitting the ball back and forth. The ball bounces from one player to the other, over and over again. The ball keeps the players moving together on the court. What if the two players represented the nuclei of two atoms and the ball represented valence electrons ? What would the back and forth movement of the ball represent? The answer is a covalent bond.

 

Sharing Electrons

A covalent bond is the force of attraction that holds together two atoms that share a pair of valence electrons . The shared electrons are attracted to the nuclei of both atoms. This forms a molecule consisting of two or more atoms. Covalent bonds form only between atoms of nonmetals .

To understand why chemical bonds form, consider the common compound known as water , or H 2 O. It consists of two hydrogen (H) atoms and one oxygen (O) atom . As you can see on the left side of the Figure below , each hydrogen atom has just one electron, which is also its sole valence electron. The oxygen atom has six valence electrons.  These are the electrons in the outer energy level of the oxygen atom.

In the water molecule on the right in the Figure above , each hydrogen atom shares a pair of electrons with the oxygen atom. By sharing electrons, each atom has electrons available to fill its sole or outer energy level . The hydrogen atoms each have a pair of shared electrons, so their first and only energy level is full. The oxygen atom has a total of eight valence electrons , so its outer energy level is full. A full outer energy level is the most stable possible arrangement of electrons. It explains why elements form chemical bonds with each other. When atoms of different elements form covalent bonds, a new substance, called a covalent compound , results. A molecule is the smallest particle of a covalent compound that still has the properties of the compound.

 

Diatomic Elements

The diagram in the Figure below shows an example of covalent bonds between two atoms of the same element , in this case two atoms of oxygen. The diagram represents an oxygen molecule, so it’s not a new compound . Oxygen normally occurs in diatomic (“two-atom”) molecules. Several other elements also occur as diatomic molecules: hydrogen, nitrogen, and all but one of the halogens (fluorine, chlorine, bromine, and iodine).

The two oxygen atoms share two pairs of electrons, so two covalent bonds hold the oxygen molecules together.

 

Properties of Covalent Compounds

Covalent compounds have properties very different from ionic compounds. Ionic compounds have high melting points causing them to be solid at room temperature, and conduct electricity when dissolved in water. Covalent compounds have low melting points and many are liquids or gases at room temperature. Whereas most ionic compounds are capable of dissolving in water, many covalent compounds do not. Also unlike ionic compounds, when covalent compounds are dissolved in water, they are not conductors of electricity; they are nonelectrolytes.

 

Bond Polarity

Electronegativity is defined as the ability of an atom to attract electrons when the atoms are in a compound . Electronegativities of elements are shown in the periodic table below.

The degree to which a given bond is ionic or covalent is determined by calculating the difference in electronegativity between the two atoms involved in the bond.

As an example, consider the bond that occurs between an atom of potassium and an atom of fluorine. Using the table, the difference in electronegativity is equal to 4.0 - 0.8 = 3.2. Since the difference in electronegativity is relatively large, the bond between the two atoms is ionic. Since the fluorine atom has a much larger attraction for electrons than the potassium atom does, the valence electron from the potassium atom is completely transferred to the fluorine atom. The diagram below shows how difference in electronegativity relates to the ionic or covalent character of a chemical bond .

Bond type is predicated on the difference in electronegativity of the two elements involved in the bond.

 

Non-polar Covalent Bonds

A bond in which the electronegativity difference is less than 1.7 is considered to be mostly covalent in character. However, at this point we need to distinguish between two general types of covalent bonds. A non-polar covalent bond is a covalent bond in which the bonding electrons are shared equally between the two atoms. In a non-polar covalent bond, the distribution of electrical charge is balanced between the two atoms.

A nonpolar covalent bond is one in which the distribution of electron density between the two atoms is equal.

The two chlorine atoms share a pair of electrons in the single covalent bond equally, and the electron density surrounding the Cl 2 molecule is symmetrical. Also note that molecules in which the electronegativity difference is very small (<0.4) are also considered non-polar covalent. An example would be a bond between chlorine and bromine .

(ΔEN=3.0−2.8=0.2)

 

Polar Covalent Bonds

A bond in which the electronegativity difference between the atoms is between 0.4 and 1.7 is called a polar covalent bond. A polar covalent bond is a covalent bond in which the atoms have an unequal attraction for electrons and so the sharing is unequal. In a polar covalent bond, sometimes simply called a polar bond, the distribution of electrons around the molecule is no longer symmetrical.

In the polar covalent bond of HF, the electron density is unevenly distributed. There is a higher density (red) near the fluorine atom, and a lower density (blue) near the hydrogen atom.

An easy way to illustrate the uneven electron distribution in a polar covalent bond is to use the Greek letter delta

             

Use of δ to indicate partial charge.

The atom with the greater electronegativity acquires a partial negative charge, while the atom with the lesser electronegativity acquires a partial positive charge. The delta symbol is used to indicate that the quantity of charge is less than one. A crossed arrow can also be used to indicate the direction of greater electron density. Use of crossed arrow to indicate polarity.             

 

Putting It Together

Let us revisit this phenomenon:

This periodic table is showing electronegativity values for the periodic table.

1. Why is there a main group that is left out of the electronegativity values?
2. What type of bond would form when calcium and phosphorus react? What type of bond would form when carbon and phosphorus react?
3. How does electronegativity impact the type of chemical bond?