Metals & Non-Metals

As shown on the periodic table of the elements below, the majority of the chemical elements in pure form are classified as metals. It seems appropriate to describe what is meant by "metal" in general terms. This general description is adapted from Shipman, et al.

Chemical Properties
Metals	Nonmetals
Usually have 1-3 electrons in their outer shell. Lose their valence electrons easily. Form oxides that are basic. Are good reducing agents. Have lower electronegativities.	Usually have 4-8 electrons in their outer shell. Gain or share valence electrons easily. Form oxides that are acidic. Are good oxidizing agents. Have higher electronegativities.
Physical Properties
Metals	Nonmetals
Good electrical conductors andheat conductors. Malleable - can be beaten into thin sheets. Ductile - can be stretched into wire. Possess metallic luster. Opaque as thin sheet. Solid at room temperature (except Hg).	Poor conductors of heat and electricity. Brittle - if a solid. Nonductile. Do not possess metallic luster. Transparent as a thin sheet. Solids, liquids or gases at room temperature.

Chemical Reactions

A chemical reaction is a process in which one set of chemical substances (reactants) is converted into another (products). It involves making and breaking chemical bonds and the rearrangement of atoms. Chemical reactions are represented by balanced chemical equations, with chemical formulas symbolizing reactants and products. For specific chemical reactants, two questions may be posed about a possible chemical reaction. First, will a reaction occur? Second, what are the possible products if a reaction occurs? This

Sulfur reacting to heat.

entry will focus only on the second question. The most reliable answer is obtained by conducting an experiment—mixing the reactants and then isolating and identifying the products. We can also use periodicity, since elements within the same group in the Periodic Table undergo similar reactions. Finally, we can use rules to help predict the products of reactions, based on the classification of inorganic chemical reactions into four general categories: combination, decomposition, single-displacement, and double-displacement reactions.

Reactions may also be classified according to whether the oxidation number of one or more elements changes. Those reactions in which a change in oxidation number occurs are called oxidation–reduction reactions . One element increases its oxidation number (is oxidized), while the other decreases its oxidation number (is reduced).

Combination Reactions

In combination reactions, two substances, either elements or compounds, react to produce a single compound. One type of combination reaction involves two elements. Most metals react with most nonmetals to form ionic compounds. The products can be predicted from the charges expected for cations of the metal and anions of the nonmetal. For example, the product of the reaction between aluminum and bromine can be predicted from the following charges: 3+ for aluminum ion and 1− for bromide ion. Since there is a change in the oxidation numbers of the elements, this type of reaction is an oxidation–reduction reaction:

2Al ( s ) + 3Br ₂ ( g ) → 2AlBr ₃ ( s )

Similarly, a nonmetal may react with a more reactive nonmetal to form a covalent compound. The composition of the product is predicted from the common oxidation numbers of the elements, positive for the less reactive and negative for the more reactive nonmetal (usually located closer to the upper right side of the Periodic Table). For example, sulfur reacts with oxygen gas to form gaseous sulfur dioxide:

S ₈ ( s ) + 8O ₂ ( g ) → 8SO ₂ ( g )

A compound and an element may unite to form another compound if in the original compound, the element with a positive oxidation number has an accessible higher oxidation number. Carbon monoxide, formed by the burning of hydrocarbons under conditions of oxygen deficiency, reacts with oxygen to form carbon dioxide:

2CO ( g ) + O ₂ ( g ) → 2CO ₂ ( g )

The oxidation number of carbon changes from +2 to +4 so this reaction is an oxidation–reduction reaction.

Two compounds may react to form a new compound. For example, calcium oxide (or lime) reacts with carbon dioxide to form calcium carbonate (limestone):

CaO ( s ) + CO ₂ ( g ) → CaCO ₃ ( s )

Decomposition Reactions

When a compound undergoes a decomposition reaction, usually when heated, it breaks down into its component elements or simpler compounds. The products of a decomposition reaction are determined largely by the identity of the anion in the compound. The ammonium ion also has characteristic decomposition reactions.

A few binary compounds decompose to their constituent elements upon heating. This is an oxidation–reduction reaction since the elements undergo a change in oxidation number. For example, the oxides and halides of noble metals (primarily Au, Pt, and Hg) decompose when heated. When red solid mercury(II) oxide is heated, it decomposes to liquid metallic mercury and oxygen gas:

2HgO ( s ) → 2Hg ( l ) + O ₂ ( g )

Some nonmetal oxides, such as the halogen oxides, also decompose upon heating:

2Cl ₂ O ₅ ( g ) → 2Cl ₂ ( g ) + 5O ₂ ( g )

Other nonmetal oxides, such as dinitrogen pentoxide, decompose to an element and a compound:

2N ₂ O ₅ ( g ) → O ₂ ( g ) + 4NO ₂ ( g )

Many metal salts containing oxoanions decompose upon heating. These salts either give off oxygen gas, forming a metal salt with a different nonmetal anion, or they give off a nonmetal oxide, forming a metal oxide. For example, metal nitrates containing Group 1A or 2A metals or aluminum decompose to metal nitrites and oxygen gas:

Mg(NO ₃ ) ₂ ( s ) → Mg(NO ₂ ) ₂ ( s ) + O ₂ ( g )

All other metal nitrates decompose to metal oxides, along with nitrogen dioxide and oxygen:

2Cu(NO ₃ ) ₂ ( s ) → 2CuO ( s ) + 4NO ₂ ( g ) + O ₂ ( g )

Salts of the halogen oxoanions decompose to halides and oxygen upon heating:

2KBrO ₃ ( s ) → 2KBr ( s ) + 3O ₂ ( g )

Carbonates, except for those of the alkali metals, decompose to oxides and carbon dioxide.

CaCO ₃ ( s ) → CaO ( s ) + CO ₂ ( g )

A number of compounds—hydrates, hydroxides, and oxoacids—that contain water or its components lose water when heated. Hydrates, compounds that contain water molecules, lose water to form anhydrous compounds, free of molecular water.

CaSO ₄ · 2H ₂ O ( s ) → CaSO ₄ ( s ) + 2H ₂ O ( g )

Metal hydroxides are converted to metal oxides by heating:

2Fe(OH) ₃ ( s ) → Fe ₂ O ₃ ( s ) + 3H ₂ O ( g )

Most oxoacids lose water until no hydrogen remains, leaving a nonmetal oxide:

H ₂ SO ₄ ( l ) → H ₂ O ( g ) + SO ₃ ( g )

Oxoanion salts that contain hydrogen ions break down into the corresponding oxoanion salts and oxoacids:

Ca(HSO ₄ ) ₂ ( s ) → CaSO ₄ ( s ) + H ₂ SO ₄ ( l )

Finally, some ammonium salts undergo an oxidation–reduction reaction when heated. Common salts of this type are ammonium dichromate, ammonium permanganate, ammonium nitrate, and ammonium nitrite. When these salts decompose, they give off nitrogen gas and water.

(NH ₄ ) ₂ Cr ₂ O ₇ ( s ) → Cr ₂ O ₃ ( s ) + 4H ₂ O ( g ) + N ₂ ( g )

2NH ₄ NO ₃ ( s ) → 2N ₂ ( g ) + 4H ₂ O ( g ) + O ₂ ( g )

Ammonium salts, which do not contain an oxidizing agent, lose ammonia gas upon heating:

(NH ₄ ) ₂ SO ₄ ( s ) → 2NH ₃ ( g ) + H ₂ SO ₄ ( l )

Single-Displacement Reactions

In a single-displacement reaction, a free element displaces another element from a compound to produce a different compound and a different free element. A more active element displaces a less active element from its compounds. These are all oxidation–reduction reactions. An example is the thermite reaction between aluminum and iron(III) oxide:

2Al ( s ) + Fe ₂ O ₃ ( s ) → Al ₂ O ₃ ( s ) + 2Fe ( l )

The element displaced from the compound is always the more metallic element—the one nearer the bottom left of the Periodic Table. The displaced element need not always be a metal, however. Consider a common type of single-displacement reaction, the displacement of hydrogen from water or from acids by metals.

The very active metals react with water. For example, calcium reacts with water to form calcium hydroxide and hydrogen gas. Calcium metal has an oxidation number of 0, whereas Ca ²⁺ in Ca(OH) ₂ has an oxidation number of +2, so calcium is oxidized. Hydrogen's oxidation number changes from +1 to 0, so it is reduced.

Ca ( s ) + 2H ₂ O ( l ) → Ca(OH) ₂ ( aq ) + H ₂ ( g )

Some metals, such as magnesium, do not react with cold water, but react slowly with steam:

Mg ( s ) + 2H ₂ O ( g ) → Mg(OH) ₂ ( aq ) + H ₂ ( g )

Still less active metals, such as iron, do not react with water at all, but react with acids.

Fe ( s ) + 2HCl ( aq ) → FeCl ₂ ( aq ) + H ₂ ( g )

Metals that are even less active, such as copper, generally do not react with acids.

To determine which metals react with water or with acids, we can use an activity series (see Figure 1), a list of metals in order of decreasing activity. Elements at the top of the series react with cold water. Elements above hydrogen in the series react with acids; elements below hydrogen do not react to release hydrogen gas.

The displacement of hydrogen from water or acids is just one type of single-displacement reaction. Other elements can also be displaced from their compounds. For example, copper metal reduces aqueous solutions of ionic silver compounds, such as silver nitrate, to deposit silver metal. The copper is oxidized.

Cu ( s ) + 2AgNO ₃ ( aq ) → Cu(NO ₃ ) ₂ ( aq ) + 2Ag ( s )

The activity series can be used to predict which single-displacement reactions will take place. The elemental metal produced is always lower in the activity series than the displacing element. Thus, iron could be displaced from FeCl ₂ by zinc metal but not by tin.

Figure 1. Activity series.

ACTIVITY SERIES
Li
K	These metals will displace hydrogen gas from water
Ba
Ca
Na
Mg
Al
Zn	These metals will displace hydrogen gas from acids
Fe
Cd
Ni
Sn
Pb
H
Cu
Hg	These metals will not displace hydrogen gas from water or acids
Ag
Au

Double-Displacement Reactions

Aqueous barium chloride reacts with sulfuric acid to form solid barium sulfate and hydrochloric acid:

BaCl ₂ ( aq ) + H ₂ SO ₄ ( aq ) → BaSO ₄ ( s ) + 2HCl ( aq )

Sodium sulfide reacts with hydrochloric acid to form sodium chloride and hydrogen sulfide gas:

Na ₂ S ( aq ) + 2HCl ( aq ) → 2NaCl ( aq ) + H ₂ S ( g )

Potassium hydroxide reacts with nitric acid to form water and potassium nitrate:

KOH ( aq ) + HNO ₃ ( aq ) → H ₂ O ( l ) + KNO ₃ ( aq )

These double-displacement reactions have two major features in common. First, two compounds exchange ions or elements to form new compounds. Second, one of the products is either a compound that will separate from the reaction mixture in some way (commonly as a solid or gas) or a stable covalent compound, often water.

Double-displacement reactions can be further classified as precipitation, gas formation, and acid–base neutralization reactions.

Precipitation Reactions

Precipitation reactions are those in which the reactants exchange ions to form an insoluble salt—one which does not dissolve in water. Reaction occurs when two ions combine to form an insoluble solid or precipitate. We predict whether such a compound can be formed by consulting solubility rules (see Table 1). If a possible product is insoluble, a precipitation reaction should occur.

A mixture of aqueous solutions of barium chloride and sodium sulfate contains the following ions: Ba ²⁺ ( aq ), Cl ⁻ ( aq ), Na ⁺ ( aq ), and SO ₄ ²⁻ ( aq ). According to solubility rules, most sulfate, sodium, and chloride salts are soluble. However, barium sulfate is insoluble. Since a barium ion and sulfate ion could combine to form insoluble barium sulfate, a reaction occurs.

Table 1.

SOME SOLUBILITY RULES FOR INORGANIC SALTS IN WATER
Compound	Solubility
Na + , K + , NH 4 +	Most salts of sodium, potassium, and ammonium ions are soluble.
NO 3 −	All nitrates are soluble.
SO 4 2−	Most sulfates are soluble. Exceptions: BaSO 4 , SrSO 4 , PbSO 4 , CaSO 4 , Hg 2 SO 4 , and Ag 2 SO 4 .
Cl − , Br − , I − ,	Most chlorides, bromides, and iodides are soluble. Exceptions: AgX, Hg 2 X 2 , PbX 2 , and HgI 2 .
Ag +	Silver salts, except AgNO 3 , are insoluble.
O 2− , OH −	Oxides and hydroxides are insoluble. Exceptions: NaOH, KOH, NH 4 OH, Ba(OH) 2 , and Ca(OH) 2 (somewhat soluble).
S 2−	Sulfides are insoluble. Exceptions: salts of Na + , K + , NH 4 + and the alkaline earth metal ions.
CrO 4 2−	Most chromates are insoluble. Exceptions: salts of K + , Na + , NH 4 + , Mg 2+ , Ca 2+ , Al 3+ , and Ni 2+ .
CO 3 2− , PO4 3− , SO 3 2− , SiO 3 2−	Most carbonates, phosphates, sulfites, and silicates are insoluble. Exceptions: salts of K + , Na + , and NH 4 + .

BaCl ₂ ( aq ) + Na ₂ SO ₄ ( aq ) → BaSO ₄ ( s ) + 2NaCl ( aq )

Gas-Formation Reactions

A double-displacement reaction should also occur if an insoluble gas is formed. All gases are soluble in water to some extent, but only a few gases [HCl ( g ) and NH ₃ ( g )] are highly soluble. All other gases, generally binary covalent compounds, are sufficiently insoluble to provide a driving force if they are formed as a reaction product. For example, many sulfide salts will react with acids to form gaseous hydrogen sulfide:

ZnS ( s ) + 2HCl ( aq ) → ZnCl ₂ ( aq ) + H ₂ S ( g )

Insoluble gases are often formed by the breakdown of an unstable double-displacement reaction product. For example, carbonates react with acids to form carbonic acid (H ₂ CO ₃ ), an unstable substance that readily decomposes into water and carbon dioxide. Calcium carbonate reacts with hydrochloric acid to form calcium chloride and carbonic acid:

CaCO ₃ ( s ) + 2HCl ( aq ) → CaCl ₂ ( aq ) + H ₂ CO ₃ ( aq )

Carbonic acid decomposes into water and carbon dioxide:

H ₂ CO ₃ ( aq ) → H ₂ O ( l ) + CO ₂ ( g )

The net reaction is:

CaCO ₃ ( s ) + 2HCl (aq) → CaCl ₂ ( aq ) + H ₂ O ( l ) + CO ₂ ( g )

Sulfites react with acids in a similar manner to release sulfur dioxide.

Acid-Base Neutralization Reactions

A neutralization reaction is a double-displacement reaction of an acid and a base. Acids are compounds that can release hydrogen ions; bases are compounds that can neutralize acids by reacting with hydrogen ions. The most common bases are hydroxide and oxide compounds of the metals. Normally, an acid reacts with a base to form a salt and water. Neutralization reactions occur because of the formation of the very stable covalent water molecule, H ₂ O, from hydrogen and hydroxide ions.

HCl ( aq ) + NaOH ( aq ) → NaCl ( aq ) + H ₂ O ( l )

Recognizing the pattern of reactants (element or compound, and the number of each) allows us to assign a possible reaction to one of the described classes. Recognizing the class of reaction allows us to predict possible products with some reliability

Acids, Bases, and Salts

Acids
In everyday life we deal with many compounds that chemists classify as acids. For example, orange juice and grapefruit juice contain citric acid. These juices, and others, also contain ascorbic acid, a substance more commonly known as Vitamin C. Salads are often flavored with vinegar, which contains dilute acetic acid. Boric acid is a substance that is sometimes used to wash the eyes.
In any chemistry laboratory, we find acids such as hydrochloric acid, sulfuric acid, and nitric acid. These acids are called mineral acids because they can be prepared from naturally occurring compounds called minerals. Mineral acids are generally stronger than household acids, and should be handled with great care because they can burn skin and clothing.
Properties of Acids:
Acids taste sour. Citric acid is responsible for the sour taste of lemons, limes, grapefruits, and oranges. Acetic acid is responsible for the sour taste of vinegar.
Acids turn litmus (or indicator papers) red. Litmus is a vegetable dye that may be either red or blue, depending on the acidity. When a sample of an acid is placed on red litmus paper, the color of the litmus does not change. Red litmus has been previously treated with acid. Adding more acid does not change the red color. However, when the same acid is placed on blue litmus paper, the color turns from blue to red. (Blue litmus has been treated with a base).
Acids contain combined hydrogen. When a sample of zinc, a fairly reactive metal, is dropped into a test tube containing an acid such as hydrochloric acid, a reaction occurs. The bubbling in the tube indicates that a gas is released. When we test this gas by inserting a burning splint into the test tube, the gas burst into flame and produces a small popping sound. This is the characteristic test for hydrogen gas. In general, when certain acids react with metals, hydrogen gas is released. See following reactions:
Zn (s) + 2HCl (aq) 􀃆 H2 (g) + ZnCl2 (aq)
Zn (s) + H2SO4 (aq) 􀃆 H2 (g) + ZnSO4 (aq)
Acids release hydrogen in water solutions. When an acid dissolves in water, the acid ionizes, releasing both hydrogen ions and ions of a nonmetal or nonmetallic polyatomic ion. Thus, when hydrochloric acid is dissolved in water, the acid ionizes, forming hydrogen ions and chloride ions, as shown in the following equation:
HCl (aq) 􀃆 H+ (aq) + Cl1- (aq)
Other examples:
H2SO4 (aq) 􀃆 2 H+ (aq) + SO42- (aq)
H3PO4 (aq) 􀃆 3 H+ (aq) + PO43- (aq)
Special example:
HC2H3O2 (aq) 􀃅 -> H+ (aq) + C2H3O21- (aq)
Note the use of the double arrow in the ionization of acetic acid. We know that acetic acid is a weak acid. The smaller arrow pointing to the right indicates that the change to the right (ionization) is relatively small. This means that, in a solution of acetic acid, we have a large number of acetic acid molecules and few hydrogens and few acetate ions.
Thus acids are defined as substances that release hydrogen ions in solution. It is these H+ (aq) that are responsible for the properties of acids.
How Acids are Prepared:
Acids may be prepared when certain gases, related to acids, are dissolved in water and when certain salt compounds, again related to acids, are allowed to react with sulfuric acid.
CO2 (g) + H2O (l) 􀃆 H2CO3 (aq) carbonic acid
2 NaCl (s) + H2SO4 (aq) + heat 􀃆 2 HCl (g) + Na2SO4 (aq)
Uses of Acids:
Sulfuric acid is the chemical most widely used in industry. Sulfuric acid is also used to make other acids such as hydrochloric and nitric acid because the boiling point of sulfuric acid is higher that that of other acids. This allows the acid being produced to be distilled and collected separately from the starting material.
Sulfuric acid is also used to remove the surface oxide layers on metals (pickling) before the metals are coated with materials that prevent rusting. For example, before iron is coated with chromium (in chromium plating), the iron is dipped into dilute sulfuric acid to remove the iron oxide normally present on the surface of the iron. Another important use of sulfuric acid is the storage cell. In a lead storage cell, dilute sulfuric acid serves as the electrolyte through which ions move between the lead plates, acting as
the cathode, and the spongy lead dioxide, acting as the anode. Several such cells connected together make up the type of storage battery used in automobiles.
Nitric acid, another important industrial acid, is used in the manufacture of fertilizers, plastics, photographic film, and dyes. Nitric acid is also used in the preparation of such explosives as dynamite and TNT.
Hydrochloric acid, like sulfuric acid, is used to clean metals. Hydrochloric acid is also used to clean brick and tile; it is used in the manufacture of sugar and glue. Hydrochloric acid is produced in small quantities in the stomach where the acid aids digestion.
Bases:
Ammonium hydroxide, or ammonia water, is very irritating to the nose and the eyes. This substance, called a hydroxide, or a base, is often used in the home for cleaning because bases generally dissolve grease. Milk of magnesia (magnesium hydroxide), which is used as an antacid, is a base; lye (sodium hydroxide), which is used in the manufacture of soap, is another familiar example of base.
Bases are ionic compounds containing metal ions and hydroxide ions. For example, sodium hydroxide contains sodium ions and hydroxide ions. When sodium metal is placed in water, sodium hydroxide is formed and hydrogen gas is released. Since the formula for water can be written as HOH instead of H2O, the reaction involves single replacement:
2 Na (s) + 2 HOH (l) 􀃆 2 NaOH (aq) + H2 (g)
Properties of Bases: (in water solutions)
1. Bases taste bitter. A bitter taste is characteristic of all bases. It is the presence of a base that give unflavored milk of magnesia its bitter taste.
2. Bases feel slippery. If you rub a drop or two of household ammonia between your fingers, you experience the slippery feeling of a base. Wet soap is also slippery because of the presence of a base.
3. Bases turn red litmus blue. A common indicator, used to detect the presence of a base, is phenolphthalein which, when mixed with a base, turns pink.
4. Bases release hydroxide ions in water solutions. When dissolved in water, bases ionize releasing metal ions (or metallic polyatomic ions) and hydroxide ions. For example: when sodium hydroxide is dissolved in water, it ionizes as:
NaOH (s) + H2O (l) 􀃆 Na1+ (aq) + OH1- (aq)
NH4OH (aq) 􀃅 -> NH41+ (aq) + OH1- (aq) double arrow indicates ionization of weak base ammonium hydroxide
Thus bases are defined as substances that release hydroxide ions in solution. It is these OH1- (aq) ions that are responsible for the properties of bases.
How Bases are Prepared:
Bases may be prepared by the reaction of water and very reactive metals, related to the base, and by the reaction of water and certain oxides, again related to the base.
2 K (s) + 2 HOH (l) 􀃆 2 KOH (aq) + H2 (g)
CaO (s) + HOH (l) 􀃆 Ca(OH)2 (aq) plus considerable heat released
Uses of Bases:
Ammonium hydroxide, frequently called ammonia, is used in the preparation of important related compounds such as nitric acid and ammonium chloride. Ammonia is also used as a cleaning agent.
Sodium hydroxide is used in the manufacture of soap, rayon, and paper. Strong solutions of this base are very caustic; that is, they are extremely harmful to the skin.
Calcium hydroxide, commonly known as slaked lime, is used in the preparation of plaster and mortar. Water solutions of calcium hydroxide, called limewater, can be used in the lab as a test for the presence of carbon dioxide.
Salts:
Many chemical compounds may be classified as salts. The salt most familiar to all of us is table salt -- sodium chloride. Baking soda is the salt sodium bicarbonate. Magnesium sulfate, also called Epsom salts, is often found in the home.
In general, salts are ionic compounds that are composed of metallic ions and nonmetallic ions. For example, sodium chloride is composed of metallic sodium ions and nonmetallic chloride ions. Some salts are composed of metallic polyatomic ions and nonmetallic polyatomic ions (ammonium nitrate is composed of ammonium ions and nitrate ions).
Properties of Salts:
The salty taste of ocean water is due to the presence of salts such as sodium chloride and magnesium bromide. There are many different salts present in salt water:
salt in saltwater
formula
percentage in saltwater
sodium chloride
NaCl
(2.72 %)
magnesium chloride
MgCl2
(0.38%)
magnesium sulfate
MgSO4
(0.17 %)
calcium sulfate
CaSO4
(0.13 %)
potassium chloride
KCl
(0.09 %)
calcium carbonate
CaCO3
(0.01 %)
magnesium bromide
MgBr2
(0.01 %)
Salts dissociate in water. Salts consist of tightly bonded ions. In water, these bonds are weakened and the ions become mobile. This accounts for the fact that salt solutions are generally electrolytes. In water, for example, sodium chloride ionizes, or dissociates like this:
NaCl (s) 􀃆 Na1+ (aq) + Cl1- (aq)
Salts may react with water. Some salt solutions, when tested with litmus show acid, others base, and others still nothing. How does this happen?
1) When sodium carbonate dissolves in water, the salt liberates sodium ions and carbonate ions. At the same time, the water itself ionizes slightly to form hydrogen and hydroxide ions (remember that water is a weak electrolyte):
Na2CO3 (s) 􀃆 2 Na1+ (aq) + CO32- (aq) HOH 􀃆 H1+ (aq) + OH1- (aq)
2) Thus, the following particles may be present in a solution of sodium carbonate: water molecules, sodium ions, carbonate ions, hydrogen ions, and hydroxide ions. The ions of opposite charge attract one another and combine to form sodium hydroxide and carbonic acid:
2 Na1+ (aq) + CO32- (aq) + 2 H1+ (aq) + 2 OH1- (aq) 􀃆 2 NaOH (aq) + H2CO3 (aq)
The reaction of a salt and water to form an acid and a base is called a hydrolysis reaction. Since acids and bases react to form water and salt (neutralization reactions) hydrolysis reactions are the reverse of neutralization reactions. Thus, when sodium carbonate is dissolved in water, carbonic acid and sodium hydroxide are formed. Carbonic acid, H2CO3 is the acid present in soda water. Since carbonic acid decomposes on standing to form CO2 gas and H2O, it is called a weak acid. From conductivity experiments we know that sodium hydroxide, NaOH, is a strong base.
Na2CO3 (aq) + 2 HOH (l) 􀃆 H2CO3 (aq) + 2 NaOH (aq)
When the acid formed in a hydrolysis reaction is stronger that the base, the effect of such a solution on litmus is that of an acid. For example; a water solution of the weak base, ammonium hydroxide and water, produces ammonium hydroxide (a weak base) and hydrochloric acid ( strong acid). Ammonium hydroxide, on standing, decomposes to form gaseous ammonia, NH3.
When both the acid and the base in a hydrolysis reaction are equally strong (or equally weak), the effect of such a solution on litmus is neither that of an acid or a base. For example, when sodium chloride is dissolved in water and the solution is tested with litmus, neither color changes, indicating that the solution is neither acidic nor basic. Such a solution is said to be neutral and the sodium chloride has not undergone hydrolysis.
Uses of Salts:
Name of salt
Formula
Uses
ammonium chloride
NH4Cl
in soldering, as electrolyte in dry cells
sodium bicarbonate
NaHCO3
in baking powder, in manufacture of glass
sodium chloride
NaCl
for seasoning and preserving food, essential in life processes
calcium chloride
CaCl2
as a drying agent to absorb moisture, in freezing mixtures
silver bromide
AgBr
in making photographic film
potassium nitrate
KNO3
in manufacture of explosives; fertilizer
sodium nitrate
NaNO3
fertilizer; source of nitric acid
Preparation of Salts:
Salts may be prepared by three methods:
1) neutralization of acid and base -- When an acid and base react, they counteract each other, that is, they neutralize each other. Such a reaction, known as a neutralization reaction, results in the formation of water and a salt. For example, when sodium hydroxide and hydrochloric acid react, water and the salt sodium chloride are formed. This occurs because the hydrochloric acid and the sodium hydroxide first ionize, and then react. The compounds ionize releasing hydrogen, chloride, sodium, and hydroxide ions.
Since these are mobile in solution, hydrogen ions meet hydroxide ions and unite to form water. At the same time sodium ions and chloride ions remain as aqueous salt.
2) Direct combination -- When a metal reacts with a nonmetal, a salt is generally formed. For example, when the metal magnesium is burned in chlorine gas, the salt magnesium chloride is formed.
3) Metal oxide and acid -- when a metal oxide reacts with an acid, a salt is formed. For example, when calcium oxide reacts with nitric acid, the salt calcium nitrate is formed.
Acids, Bases, Salts, pH, Buffers, Indicators, Titration, Antacids
Acids Acids are substances which ionize in water solutions to produce hydrogen ion (H+ or free p+). Some ionize more than others. Acids taste sour and turn litmus red.
Bases Bases are substances which ionize to produce hydroxide ions in water solutions. They taste bitter, feel slippery to the touch, and turn litmus blue.
Salts Salts are crystalline compounds composed of the negative ions of an acid and the positive ions of a base.
Hydronium ion H+ combines with a water molecule to form H30+ any solution containing hydronium ions is acidic and the strength of an acid is based on the number of hydronium ions in solution.
Ionization Ionization is the formation of ions in solution.
Dissociation Dissociation is the separation of ions in solution.
Neutralization Neutralization is the reaction of an acid and a base to form a salt plus water.
Hydrolysis Hydrolysis is a reaction of a salt with water to form a weak acid or base.
Titration Titration is a technique for measuring the relative strength of a solution.
Endpoint Endpoint is the point in a titration where equal amounts of reactants are present.
Buffer A buffer is a solution which can receive moderate amounts of of either acid or base without significant changes in pH.
Indicator An indicator is used to indicate the pH of a solution.
pH or pOH pH is the negative log of the hydrogen ion concentration. pOH is the negative log of the hydroxide ion concentration.
A weak acid/base ionizes only slightly in solution whereas a strong acid/base completely ionizes (dissociates) into positive and negative ions in solution.
Demonstration Procedure:
I. heat 75 mL of 0.1 M HCI on hot plate
II. crush antacid tablet and add this to acid
III. add solution to graduated cylinder with enough water to have 100 mL
IV. place 25 mL of solution in Erlenmeyer flask
V. add phenolphthalein as indicator
VI. titrate with 0.1 M NaOH to endpoint
VII. test solution with calibrated pH meter (using buffer solutions)
Hydrogen Ions, pH, and Indicators
Background:
The hydrogen ion concentration [H+] in water solutions can vary over a wide range of values. These values may be over 1 Molar [H+] or less than 1 E – 14 Molar [H+]. To deal with such large possible variations a mathematical conversion of H+ to a logarithmic value has been used to express the H+ concentration as a pH.
pH = -Log10[H+]
As you know perhaps, a logarithm is the exponent of a base number which makes the exponential form equal to another natural number. For example [102 = 100], if the base number 10 has an exponent of 2, the expression is equal to 100. The exponent 2 is the logarithm of 100. The logarithm of 1000 would be 3 [103 = 1000]. In the base 10 system, a difference of one in two logarithmic values would mean a 10 fold difference in the values of the two natural numbers which they represent, i.e., the logarithms of 2 and 3 represent 100 and 1000. There are logarithm tables which will give decimal logarithmic values so that any natural number can be represented. [101.699 = 50; the logarithm of 50 is 1.699].
The variation of the value of the [H+] of 1 to 10-14 molar can be expressed as pH values of 0 to 14. Neutral water has a [H+] of 1 E –7 M and therefore has a pH of 7. Acidic solutions will have pH values less than 7 and basic solutions will have pH values greater than 7.
The pH of a solution can be used to describe the acidic or basic qualities of soil samples, cosmetics, food, beverages, etc.
The pH can be determined by electrical devices which have electrodes which are sensitive to the hydrogen ion concentration (pH meter) or by certain compounds which change color over certain pH ranges.
Indicators are weak organic acids or bases which change color (and structure) with a change in pH.
pH and Indicators
The concentration of hydrogen ion is a measure of the acidity and the basicity of a solution. The concentration of hydrogen ion may be expressed in terms of the molarity of the acid or base solution; however, it is frequently more convenient to express the concentration as a function of the hydrogen ion concentration, pH. The pH of a solution may be defined as the exponent of the hydrogen ion concentration. This definition may be stated mathematically as:
pH = - log [H+] where [H+] is the molar hydrogen ion concentration.
The pH scale for water systems ranges from a value to 0 to 14 only. See the table that follows for the strength of the acid or base in water at 25º C.
pH range
[H+] range
Strength
0 to 2
1 E 0 to 1 E - 2
Strong acid
3 to 4
1 E – 3 to 1 E – 4
Moderate acid
5 to 6
1 E – 5 to 1 E – 6
Weak acid
7
1 E – 7
Neutral
8 to 9
1 E – 8 to 1 E – 9
Weak base
10 to 11
1 E – 10 to 1 E – 11
Moderate base
12 to 14
1 E – 12 to 1 E - 14
Strong base
Indicators have been developed in order to assist in the determination of the pH of a solution. These indicators are weak organic acids or bases which have the property of changing color in solution when the hydrogen ion concentration reaches a definite value. An acid indicator may be represented by the equation: HIn = H+ + In-
The anion, In-, represents a complex organic group which has changed its structure due to the loss of a hydrogen ion. The loss of the hydrogen ion is accompanied by a change in color. Since an indicator reaction is an equilibrium reaction, the addition of hydrogen ions would force the above reaction to the left and a color indicating an acid solution would result. The addition of hydroxide ions would cause the reaction to go to the right and a color associated to a basic solution would result.
The pH ranges of some indicators are given in the following table. With this table you can estimate the pH of a solution. Suppose phenolphthalein in introduced into a solution and the color of the solution becomes red. This red color indicates that the pH of the solution is 10.0 or higher. If indigo carmine is added to a new sample of the same solution and a blue color results, the pH will be narrowed to a range of 10.0 to 11.4, since the lower limit of color change for indigo carmine is blue. By using an additional indicator or indicators and a new sample of the solution, the pH of the solution can be narrowed to a small range.
Indicator
pH Range
Color Range
Methyl violet
0.0 – 1.6
Yellow to blue
Thymol blue
1.2 – 2.8
Red to yellow
Methyl orange
3.2 – 4.4
Red to yellow
Congo red
3.0 – 5.0
Blue to red
Methyl red
4.8 – 6.0
Red to yellow
Phenol red
6.6 – 8.0
Red to blue
Litmus
4.7 – 8.2
Red to blue
Cresol red
7.4 – 8.6
Yellow to red
Phenolphthalein
8. 2 – 10.0
Colorless to red
Thymolphthalein
9.4 – 10.6
Colorless to blue
Alizarin yellow R
10.1 – 12.0
Yellow to red
Indigo carmine
11.4 – 14.0
Blue to yellow

Chemical bonding

Chemical compounds are formed by the joining of two or more atoms. A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. The bound state implies a net attractive force between the atoms ... a chemical bond. The two extreme cases of chemical bonds are: Covalent bond: bond in which one or more pairs of electrons are shared by two atoms. Ionic bond: bond in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions which attract each other. Other types of bonds include metallic bonds and hydrogen bonding. The attractive forces between molecules in a liquid can be characterized as van der Waals bonds. Sodium chloride Ionic Hydrogen molecule Covalent

Covalent Bonds Covalent chemical bonds involve the sharing of a pair of valence electrons by two atoms, in contrast to the transfer of electrons in ionic bonds. Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom. Hydrogen gas forms the simplest covalent bond in the diatomic hydrogen molecule. The halogens such as chlorine also exist as diatomic gases by forming covalent bonds. The nitrogen and oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic molecules. Covalent bonding can be visualized with the aid ofLewis diagrams. Comparison of ionic and covalent materials.

Polar Covalent Bonds Covalent bonds in which the sharing of the electron pair is unequal, with the electrons spending more time around the more nonmetallic atom, are called polar covalent bonds. In such a bond there is a charge separation with one atom being slightly more positive and the other more negative, i.e., the bond will produce adipole moment. The ability of an atom to attract electrons in the presense of another atom is a measurable property called electronegativity.

Ionic Bonds In chemical bonds, atoms can either transfer or share their valence electrons. In the extreme case where one or more atoms lose electrons and other atoms gain them in order to produce a noble gas electron configuration, the bond is called an ionic bond. Typical of ionic bonds are those in the alkali halides such as sodium chloride, NaCl. Ionic bonding can be visualized with the aid of Lewis diagrams. Comparison of ionic and covalent materials. Energy contributions to ionic bonds Table of ionic diatomic bonds

Metallic Bonds The properties of metals suggest that their atoms possess strong bonds, yet the ease of conduction of heat and electricity suggest that electrons can move freely in all directions in a metal. The general observations give rise to a picture of "positive ions in a sea of electrons" to describe metallic bonding.

Metal Properties The general properties of metals include malleability and ductility and most are strong and durable. They are good conductors of heat and electricity. Their strength indicates that the atoms are difficult to separate, but malleability and ductility suggest that the atoms are relatively easy to move in various directions. The electrical conductivity suggests that it is easy to move electrons in any direction in these materials. The thermal conductivity also involves the motion of electrons. All of these properties suggest the nature of the metallic bonds between atoms.

Hydrogen Bonding Hydrogen bonding differs from other uses of the word "bond" since it is a force of attraction between a hydrogen atom in one molecule and a small atom of highelectronegativity in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word bond. When hydrogen atoms are joined in a polar covalent bondwith a small atom of high electronegativity such as O, F or N, the partial positive charge on the hydrogen is highly concentrated because of its small size. If the hydrogen is close to another oxygen, fluorine or nitrogen in another molecule, then there is a force of attraction termed a dipole-dipole interaction. This attraction or "hydrogen bond" can have about 5% to 10% of the strength of a covalent bond. Hydrogen bonding has a very important effect on the properties of water and ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes. The "unzipping" of DNA is a breaking of hydrogen bonds which help hold the two strands of the double helix together

Atomic Structure

An atom is the smallest building block of matter. Atoms are made of neutrons, protons and electrons. The nucleus of an atom is extremely small in comparison to the atom. If an atom was the size of the Houston Astrodome, then its nucleus would be the size of a pea.

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Introduction to the Periodic Table

Scientists use the Periodic Table in order to find out important information about various elements. Created by Dmitri Mendeleev (1834-1907), the periodic table orders all known elements in accordance to their similarities. When Mendeleev began grouping elements, he noticed the Law of Chemical Periodicity. This law states, "the properties of the elements are periodic functions of atomic number." The periodic table is a chart that categorizes elements by "groups" and "periods." All elements are ordered by their atomic number. The atomic number is the number of protons per atom. In a neutral atom, the number of electrons equals the number of protons. The periodic table represents neutral atoms. The atomic number is typically located above the element symbol. Beneath the element symbol is the atomic mass. Atomic mass is measured in Atomic Mass Units where 1 amu = (1/12) mass of carbon measured in grams. The atomic mass number is equal to the number of protons plus neutrons, which provides the average weight of all isotopes of any given element. This number is typically found beneath the element symbol. Atoms with the same atomic number, but different mass numbers are called isotopes. Below is a diagram of a typical cells on the periodic table.
There are two main classifications in the periodic table, "groups" and "periods." Groups are the vertical columns that include elements with similar chemical and physical properties. Periods are the horizontal rows. Going from left to right on the periodic table, you will find metals, then metalloids, and finally nonmetals. The 4th, 5th, and 6th periods are called the transition metals. These elements are all metals and can be found pure in nature. They are known for their beauty and durability. The transition metals include two periods known as the lanthanides and the actinides, which are located at the very bottom of the periodic table. The chart below gives a brief description of each group in the periodic table.

Group 1A	Known as Alkali Metals Very reactive Never found free in nature React readily with water
Group 2A	Known as Alkaline earth elements All are metals Occur only in compounds React with oxygen in the general formula EO (where O is oxygen and E is Group 2A element)
Group 3A	Metalloids Includes Aluminum (the most abundant metal in the earth) Forms oxygen compounds with a X2O3 formula
Group 4A	Includes metals and nonmetals Go from nonmetals at the top of the column to metals at the bottom All oxygen form compounds with a XO2 formula
Group 5A	All elements form an oxygen or sulfur compound with E2O3 or E2S3 formulas
Group 6A	Includes oxygen, one of the most abundant elements. Generally, oxygen compound formulas within this group are EO2 and EO3
Group 7A	Elements combine violently with alkali metals to form salts Called halogens, which mean "salt forming" Are all highly reactive
Group 8A	Least reactive group All elements are gases Not very abundant on earth Given the name noble gas because they are not very reactive

Charges in the Atom

The charges in the atom are crucial in understanding how the atom works. An electron has a negative charge, a proton has a positive charge and a neutron has no charge. Electrons and protons have the same magnitude of charge. Like charges repel, so protons repel one another as do electrons. Opposite charges attract which causes the electrons to be attracted to the protons. As the electrons and protons grow farther apart, the forces they exert on each other decrease.

Atomic Models and the Quantum Numbers

There are different models of the structure of the atom. One of the first models was created by Niels Bohr, a Danish physicist. He proposed a model in which electrons circle the nucleus in "orbits" around the nucleus, much in the same way as planets orbit the sun. Each orbit represents an energy level which can be determined using equations generated by Planck and others discussed in more detail below. The Bohr model was later proven to be incorrect, but provides a useful model for building an explanation.The "accepted" model is the quantum model. In the quantum model, we state that the electron cannot be found precisely, but we can predict the probability, or likelihood, of an electron being at some location in the atom. You should be familiar with quantum numbers, a series of three numbers used to describe the location of some object (like an electron) in three-dimensional space:

1. n: the principal quantum number, an integer value (1, 2, 3...) that is used to describe the quantum level, or shell, in which an electron resides. The principal quantum number is the primary number used to determine the amount of energy in an atom. Using one of the first important equations in atomic structure (developed by Niels Bohr), we can calculate the amount of energy in an atom with an electron at some value of n:
   

   En = -

   Rhc n2

   where:
   R = Rydberg constant, a value of 1.097 X 10⁷ m^-1
   c = speed of light, 3.00 X 10⁸ m/s
   h = Planck's constant, 6.63 X 10 ^-34 J-s
   n = principal quantum number, no unitFor example, how much energy does one electron with a principal quantum number of n= 2 have?
   
   

   En = - Rhc n2

   or

   En = - (1.097x107 m-1 ∗ (6.63x10-34 J•s)∗(3.0x108 m•s-1) 22

   = 5.5x10-19 J

   You might ask, well, who cares? In addition to the importance of knowing how much energy is in an atom (a very important characteristic!), we can also derive, or calculate, other information from this energy value. For example, can we see this energy? The table below suggests that we can. For example, suppose that an electron starts at the n=3 level (we'll call this the excited state) and it falls down to n=1 (the ground state). We can calculate the change in energy using the equation:
   
   

   ΔE = hv = RH

   1 ni2

   -

   1 nf2

   Where:
   ΔE = change in energy (Joules)
   h = Planck's constant with a value of 6.63 x 10^-34 (J-s)
   ν is frequency (s^-1)
   R_H is the Rydberg constant with a value of 2.18 x 10^-18J.
   n_i is the initial quantum number
   n_f is the final quantum number
   Using the equation below, we can calculate the wavelength and the frequency of the energy. The wavelength and the frequency give us information about how we might "see" the energy:
   
   

   vλ = c

   Where:
   ν = the frequency of radiation (s^-1)
   λ = the wavelength (m)
   c = the speed of light with a value of 3.00 x 10⁸ m/s in a vacuum
   
   

   Speed of light =	3.00E+08
   Rydberg constant =	2.18E-18
   Planck's constant =	6.63E-34
   Excited state, n =	3	4	5
   Ground state, n =	2	2	2
   Excited state energy (J)	2.42222E-19	1.363E-19	8.72E-20
   Ground state energy (J)	5.45E-19	5.45E-19	5.45E-19
   ΔE =	-3.02778E-19	-4.09E-19	-4.58E-19
   ν =	4.56678E+14	6.165E+14	6.905E+14
   λ(nm) =	656.92	486.61	434.47

   
   
2. l ("el", not the number 1): the azimuthal quantum number, a number that specifies a sublevel, or subshell, in an orbital. The value of the azimuthal quantum number is always one less than the principal quantum number n. For example, if n=1, then "el"=0. If n=3, then l can have three values: 0,1, and 2. The values of l are typically not identified as "0, 1, 2, and 3" but are more commonly called by their historic names, "s, p, d, and f", respectively. Since the quantum numbers were discovered through the study of light and lines on an electromagnetic spectra, chemists identified the lines by their quality:sharp, principal, diffuse and fundamental. The table below shows the relationship:
   

   Value of l	Subshell designation
   0	s
   1	p
   2	d
   3	f

3. m: the magnetic quantum number. Each subshell is composed of one or more orbitals. In the study of light, it was discovered that additional lines appeared in the spectra produced when light was emitted in a magnetic field. The magnetic quantum number has values between -l and +l. When l =1, for example, m can have three values: -1, 0, and +1. Because you know from the chart above that the subshell designation for l =1 is "p", you now know that the p orbital has three components. In your study of chemistry, you will be presented with p_x, p_y, and p_z. Notice how the subscripts are related to a three-dimensional coordinate system, x, y, and z. The chart below shows a summary of the quantum numbers:
   

   Principal Quantum Number (n)	Azimuthal Quantum Number (l)	Subshell Designation	Magnetic Quantum Number (m)	Number of orbitals in subshell
   1	0	1s	0	1
   2	0 1	2s 2p	0 -1 0 +1	1 3
   3	0 1 2	3s 3p 3d	0 -1 0 +1 -2 -1 0 +1 +2	1 3 5
   4	0 1 2 3	4s 4p 4d 4f	0 -1 0 +1 -2 -1 0 +1 +2 -3 -2 -1 0 +1 +2 +3	1 3 5 7

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Chemists care about where electrons are in an atom or a molecule. In the early models, we believed that electrons move like billiard balls, and followed the rules of classical physics. The graphic below attempts to show that earlier models thought that we could identify the exact path, position, velocity, etc. of an electron or electrons in an atom:
A more accurate picture is that the electron(s) reside in a "cloud" that surrounds the nucleus of the atom. This concept is shown in the graphic below:

Chemists are interested in predicting the probability that the electron will be at some particular part of this cloud. The cloud is better known as an orbital, and comes in several different types, or shapes. Atomic orbitals are known as s, p, d, and f orbitals. Each type of atomic orbital has certain characteristics, such as shape. For example, as the graphic below shows, an s orbital is spherical in shape:

On this graph, the horizontal (x) axis represents the distance from the nucleus in units of a₀, or atomic units. The value of a₀ is 0.0529 nanometers (nm). The vertical (y) axis represents the probability density. What you should notice is that as the electron moves farther away from the nucleus, the probability of its being found at that distance decreases. In other words, the electron prefers to hang around close to the nucleus.
The three graphics below show some other orbitals. The first graph (top left) is of a "2s" orbital. Each "s" orbital can hold two electrons in its cloud. Notice how there is a relatively high probability of an electron being near the nucleus, then some space where the probability is close to zero, then the probability increases substantially at some distance from the nucleus. The graphic at the top right shows a "2p" atomic orbital. Orbitals that are "p" orbitals can hold up to six (6) electrons in their cloud. Notice its "dumbbell" or "figure of eight" shape. At the bottom left is a "3s" orbital. Again, notice its spherical shape. Finally, at the bottom right, is a "3p" orbital.

Determining Electron Configuration

One of the skills you will need to learn to succeed in freshman chemistry is being able to determine the electron configuration of an atom. An electron configuration is basically an account of how many electrons there are, and in what orbitals they reside under "normal" conditions. For example, the element hydrogen (H) has one electron. We know this because its atomic number is one (1), and the atomic number tells you the number of electrons. Where does this electron go? The one electron of hydrogen goes into the lowest energy state it possibly can, which means it will start at "level" one and goes into "s" orbitals first. We say that hydrogen has a "[1s¹]" electron configuration. Looking at the next element on the Periodic Table --helium, or He -- we see it has an atomic number of two, so two electrons. Since " s" orbitals can hold up to two electrons, helium has an electron configuration of "[1s²]".What about larger atoms? Let's look at carbon, with an atomic number of 6. Where do its 6 electrons go?

• First two: 1s²
• Next two: 2s²
• Last two: 2p²

We can therefore say that carbon has the electron configuration of "[1s²2s²2p²]".
The table below shows the subshells, the number of orbitals, and the maximum number of electrons allowed:

Subshell	Number of Orbitals	Maximum Number of Electrons
s	1	2
p	3	6
d	5	10
f	7	14

The Abridged (shortened) Periodic Table below shows the electron configurations of the elements. Notice for space reasons we sometimes leave off a portion of the electron configuration. For example, look at argon (Ar), element 18. The table below shows its electron configuration as "[3s²3p⁶]" (remembering that "p" orbitals can hold up to six (6) electrons). Its actual electron configuration is:
Ar = [1s²2s²2p⁶3s²3p⁶]Sometimes you will see the notation: "[Ne]3s²3p⁶", which means to include everything that is in neon (Ne, 10) plus the stuff in the "3"-level orbitals.

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Try It Out

1. What is the frequency of infrared radiation that has a wavelength of 1.25 x 10³ nm?Using the equation:
   where:
   v = frequency (in units of "per second", or s^-1
   c = speed of light at 3.0 x 10⁸ m/s
   = wavelength in units of nanometers (nm)
   Solving for frequency, we have the equation:
   
   To ensure that our units match, we must convert the wavelength from nanometers to meters:
   1.25 x 10³ nm * (1 x 10^-9 meters / 1 nm) = 1.25 x 10^-6 m
   Now, it's just a simple "plug and chug":
   v = 3.00 x 10⁸ m s^-1 / 1.25 x 10^-6 m
   = 2.40 x 10¹⁴ s^-1 or 2.40 x 10¹⁴ s^-1 hertz (Hz)

Organic Chemistry

An Introduction

To understand life as we know it, we must first understand a little bit of organic chemistry. Organic molecules contain both carbon and hydrogen. Though many organic chemicals also contain other elements, it is the carbon-hydrogen bond that defines them as organic. Organic chemistry defines life. Just as there are millions of different types of living organismson this planet, there are millions of different organic molecules, each with different chemical and physical properties. There are organic chemicals that make up your hair, your skin, your fingernails, and so on. The diversity of organic chemicals is due to the versatility of the carbon atom. Why is carbon such a special element? Let's look at its chemistry in a little more detail.

Carbon (C) appears in the second row of the periodic table and has four bonding electrons in its valence shell (see our Periodic Table module for more information). Similar to other non-metals, carbon needs eight electrons to satisfy its valence shell. Carbon therefore forms four bonds with otheratoms (each bond consisting of one of carbon's electrons and one of the bonding atom's electrons). Every valence electron participates in bonding, thus a carbon atom's bonds will be distributed evenly over the atom's surface. These bonds form a tetrahedron (a pyramid with a spike at the top), as illustrated below:

Carbon forms 4 bonds

Organic chemicals get their diversity from the many different ways carbon can bond to other atoms. The simplest organic chemicals, calledhydrocarbons, contain only carbon and hydrogen atoms; the simplest hydrocarbon (called methane) contains a single carbon atom bonded to four hydrogen atoms:

Methane - a carbon atom bonded to 4 hydrogen atoms

But carbon can bond to other carbon atoms in addition to hydrogen, as illustrated in the molecule ethane below:

Ethane - a carbon-carbon bond

In fact, the uniqueness of carbon comes from the fact that it can bond to itself in many different ways. Carbon atoms can form long chains:

Hexane - a 6-carbon chain

branched chains:

Isohexane - a branched-carbon chain

rings:

Cyclohexane - a ringed hydrocarbon

There appears to be almost no limit to the number of different structures that carbon can form.  To add to the complexity of organic chemistry, neighboring carbon atoms can form double and triple bonds in addition to single carbon-carbon bonds:

Single bonding	Double bonding	Triple bonding

Keep in mind that each carbon atom forms four bonds. As the number of bonds between any two carbon atoms increases, the number of hydrogen atoms in the molecule decreases (as can be seen in the figures above).

Simple hydrocarbons

The simplest hydrocarbons are those that contain only carbon and hydrogen. These simple hydrocarbons come in three varieties depending on the type of carbon-carbon bonds that occur in the molecule. Alkanes are the first class of simple hydrocarbons and contain only carbon-carbon single bonds. The alkanes are named by combining a prefix that describes the number of carbon atoms in the molecule with the root ending "ane". The names and prefixes for the first ten alkanes are given in the following table.

Carbon Atoms	Prefix	Alkane Name	Chemical Formula	Structural Formula
1	Meth	Methane	CH 4	CH4
2	Eth	Ethane	C2H6	CH3CH3
3	Prop	Propane	C3H8	CH3CH2CH3
4	But	Butane	C4H10	CH3CH2CH2CH3
5	Pent	Pentane	C5H12	CH3CH2CH2CH2CH3
6	Hex	Hexane	C6H14	...
7	Hept	Heptane	C7H16
8	Oct	Octane	C8H18
9	Non	Nonane	C9H20
10	Dec	Decane	C10H22

The chemical formula for any alkane is given by the expression C_nH_2n+2. The structural formula, shown for the first five alkanes in the table, shows each carbon atom and the elements that are attached to it. This structural formula is important when we begin to discuss more complex hydrocarbons. The simple alkanes share many properties in common. All enter intocombustion reactions with oxygen to produce carbon dioxide and water vapor. In other words, many alkanes are flammable. This makes them good fuels. For example, methane is the principle component of natural gas, and butane is common lighter fluid.

CH4 + 2O2  CO2 + 2H2O

The combustion of methane

The second class of simple hydrocarbons, the alkenes, consists ofmolecules that contain at least one double-bonded carbon pair. Alkenesfollow the same naming convention used for alkanes. A prefix (to describe the number of carbon atoms) is combined with the ending "ene" to denote an alkene. Ethene, for example is the two- carbon molecule that contains one double bond. The chemical formula for the simple alkenes follows the expression C_nH_2n. Because one of the carbon pairs is double bonded, simple alkenes have two fewer hydrogen atoms than alkanes.

Ethene

Alkynes are the third class of simple hydrocarbons and are molecules that contain at least one triple-bonded carbon pair. Like the alkanes and alkenes,alkynes are named by combining a prefix with the ending "yne" to denote the triple bond. The chemical formula for the simple alkynes follows the expression C_nH_2n-2.

Ethyne

Isomers

Because carbon can bond in so many different ways, a single molecule can have different bonding configurations. Consider the two molecules illustrated here:

C6H14	CH3CH2CH2CH2CH2CH3
C6H14	CH 3   I CH3 CH2 CH CH2 CH3

Both molecules have identical chemical formulas (shown in the left column); however, their structural formulas (and thus some chemical properties) are different. These two molecules are called isomers. Isomers are molecules that have the same chemical formula but different structural formulas.

Functional groups

In addition to carbon and hydrogen, hydrocarbons can also contain otherelements. In fact, many common groups of atoms can occur within organicmolecules, these groups of atoms are called functional groups. One good example is the hydroxyl functional group. The hydroxyl group consists of a single oxygen atom bound to a single hydrogen atom (-OH). The group of hydrocarbons that contain a hydroxyl functional group is called alcohols. The alcohols are named in a similar fashion to the simple hydrocarbons, a prefix is attached to a root ending (in this case "anol") that designates the alcohol. The existence of the functional group completely changes the chemical properties of the molecule. Ethane, the two-carbon alkane, is a gas at room temperature; ethanol, the two-carbon alcohol, is a liquid.

Ethanol

Ethanol, common drinking alcohol, is the active ingredient in "alcoholic" beverages such as beer and wine.