Hydrogen is a chemical element with the symbol H and atomic number 1. It is the most abundant element in the universe and makes up about 75% of its elemental mass.
Hydrogen is a highly reactive and versatile element that can form a wide range of compounds, including simple diatomic molecules (H2), water (H2O), and organic molecules such as methane (CH4) and ethanol (C2H5OH). We will touch the chemistry of hydrogen, including its appearance, properties, reactivity, and that of the molecules and compounds.
Descriptive Chemistry of Hydrogen
Hydrogen is the first element on the periodic table with an atomic number of 1 and symbol H.
Hydrogen is a highly reactive non-metallic element in the periodic table. It has atomic weight of 1.008 because it does not have any neutrons. It is the only element with this attribute.
It is the most abundant element in the universe, constituting about 75% of its elemental mass. Hydrogen is essential for life as it is a constituent of water and most organic compounds.
Hydrogen is highly reactive, and it has numerous applications in various fields, such as chemical processing, fuel cells, and aerospace. Its chemical properties are unique and essential in many chemical reactions, which are governed by its electronic configuration.
The focus here is on discussing the occurrence, appearance the chemistry of hydrogen, including its electronic configuration, isotopes, bonding, reactions, and applications.
Also, on the periodic table, hydrogen is located in group 1 and period 1 of the periodic table.
Hydrogen has only one electron and can easily lose or gain electrons to form ions with a charge of +1 or -1. The reason it can be placed to either group one or group seven on the Periodic Table.
However, it most commonly forms a covalent bond by sharing its electron with another non-metallic element. Some examples are the alkanes, alkenes, boranes, amines etc.
The commonest bond that hydrogen forms is a single covalent bond with another hydrogen atom, resulting in the formation of H2, which is a colorless, odorless gas.
H2(g) + O2(g) → 2H2O(l)
The above reaction is exothermic, releasing a large amount of heat and light energy. It is a common reaction of hydrogen with oxygen in the presence of a spark or flame.
Oxidation States of Hydrogen
Hydrogen has three oxidation states: +1, -1, and 0.
The neutral hydrogen is zero, while the two other common oxidation states, +1 for H+ are referred to proton, and -1, which refers to H–, a hydride ion.
In the -1-oxidation state, hydrogen forms compounds called hydrides, which are formed by the reaction of hydrogen with metals. Examples of hydrides include sodium hydride (NaH), lithium hydride (LiH), NH3, and calcium hydride (CaH2).
In the +1-oxidation state, hydrogen bond with non-metals to form compounds are called acid,. Hydrogen ion is also called protonic ion (H+), the replaceable hydrogen ion that characterizes acidic substances in aqueous solution. It forms hydroxonium ion (H3O+) Examples include HF, HCl, HNO3, H2CO3.
In the 0-oxidation state, hydrogen exists as a free element as found in H2 molecule.
Hydrogen and the Position on the Periodic Table
Hydrogen’s position on the periodic table is unique and does not clearly define its properties. While located in Group 1 with the alkali metals, hydrogen does not exhibit the same metallic properties as the other elements in this group. However, it shares some characteristics with the halogens, particularly chlorine, which is located in Group 17 of the periodic table.
The electronic configuration of hydrogen, consisting of a single electron in its valence shell, makes it similar to the alkali metals in terms of its reactivity. Like the alkali metals, hydrogen readily loses its electron to form a cation with a charge of +1. This makes it an excellent reducing agent in chemical reactions. However, hydrogen does not form a metal cation like the alkali metals, and it is not a good conductor of electricity.
Hydrogen also shares some characteristics with the halogens, particularly chlorine, which is a highly electronegative element. Hydrogen and chlorine can react to form hydrogen chloride, which is a corrosive gas used in industrial processes such as the production of PVC plastic. The reaction is highly exothermic, and the resulting gas can cause severe respiratory problems if inhaled.
H2(g) + Cl2(g) → 2HCl(g)
Hydrogen can also react with other elements, such as carbon or sulfur, to form a variety of compounds. For example, when hydrogen reacts with carbon, it can form methane, which is a major component of natural gas. The reaction is typically catalyzed by metals such as nickel or palladium.
CH4(g) + 2H2(g) → C2H6(g)
Hydrogen can also react with sulfur to form hydrogen sulfide, which is a toxic gas with a distinctive rotten egg smell. The reaction is typically catalyzed by metals such as iron or zinc.
H2(g) + S(g) → H2S(g)
Electronic Configuration
The electronic configuration of hydrogen is simple and unique, with only one electron in its outermost shell. Its electronic configuration is 1s1, which means that it has one electron in the 1s orbital.
Please note: The electronic configuration of neutral hydrogen (H) is 1s1, while H+ is 1s0 and H– is 1s2.
Occurrence of Hydrogen
Hydrogen is the most abundant element in the universe, but it is not found in its free form on Earth. It is found in combined form with other elements, such as oxygen in water (H2O), and carbon in hydrocarbons.
Hydrogen is also found in the Earth’s atmosphere, where it constitutes about 0.00005% by volume.
It is produced by the action of sunlight on water vapor in the upper atmosphere, and it is also produced by the decay of radioactive elements in the Earth’s crust.
Preparation of Hydrogen
Hydrogen gas can be prepared by several methods.
One common method is the reaction of a metal with an acid. When a reactive metal, such as zinc or iron, is reacted with an acid, such as hydrochloric acid (HCl) or sulfuric acid (H2SO4), hydrogen gas is produced.
Zn + 2HCl → ZnCl2 + H2
Another method of preparing hydrogen gas is through the electrolysis of water.
In this process, an electric current is passed through water to separate the hydrogen and oxygen gases.
2H2O → 2H2 + O2
This process is commonly used in industry to produce large quantities of hydrogen gas. It requires a source of electricity, such as a power plant, to provide the energy needed for the electrolysis reaction.
A third method of preparing hydrogen gas is through the steam reforming of natural gas.
In this process, natural gas (primarily methane) is reacted with steam using nickel catalyst to produce hydrogen gas and carbon monoxide.
CH4 + H2O → CO + 3H2
This process is commonly used in the production of hydrogen gas for industrial applications, such as in the manufacture of ammonia for fertilizers.
Also, a method, water gas shift reaction where carbon monoxide and water react over iron and copper catalysts to produce carbon dioxide and hydrogen gas.
CO + H2O à CO2 + H2
Appearance of Hydrogen
Hydrogen is a colorless, odorless, and tasteless gas.
It is the lightest gas (this attributable to its molecular weight) and is therefore used in balloons and airships.
It is also highly flammable, and it burns with a pale blue flame.
Hydrogen gas is not very soluble in water, and it is less dense than air.
It can be liquefied at very low temperatures and compressed to very high pressures.
Physical Properties of Hydrogen
Hydrogen has a melting point of -259.14°C and a boiling point of -252.87°C.
It has a density of 0.0899 g/cm3 at 0°C and 1 atm pressure, which is less than one-fourteenth that of air. Hydrogen has a low viscosity and a low surface tension, which make it a poor conductor of heat and electricity.
Its thermal conductivity is about one-fourth that of copper, and its electrical conductivity is about one-seventh that of copper.
Physical Properties of Hydrogen
| Property | Description |
|---|---|
| Appearance | Colorless, odorless, tasteless gas |
| Boiling Point | -252.87°C |
| Melting Point | -259.14°C |
| Density | 0.0899 g/cm³ |
| Solubility | Not very soluble in water |
| Compressibility | Can be compressed to high pressures |
| Flammability | Highly flammable |
Isotopes of Hydrogen
Hydrogen has three isotopes: protium (1H), deuterium (2H or 2D), and tritium 3H or 3T).
Protium
Protium is the most abundant isotope, making up more than 99% of natural hydrogen.
Protium has only one proton and one electron in its nucleus and electron cloud, respectively.
It has no neutron and a mass number of 1. Protium has the same atomic number as hydrogen and behaves chemically like hydrogen.
Protium is stable and does not undergo radioactive decay.
It is essential for life and is present in all living organisms in the form of water, organic compounds, and other biomolecules.
Deuterium
Deuterium, also known as heavy hydrogen, is the second most abundant isotope of hydrogen and accounts for less than 1% of all hydrogen atoms.
Deuterium has one proton, one electron, and one neutron in its nucleus and electron cloud, respectively.
It has a mass number of 2 and is also known as heavy hydrogen.
Deuterium is stable and does not undergo radioactive decay.
Deuterium has a higher boiling and melting point than protium due to its higher mass.
It also has a higher density and is less soluble in water than protium.
Deuterium is used as a tracer to study chemical reactions, metabolic pathways, and other biological processes.
It is also used in nuclear fusion research to produce energy.
Tritium
Tritium accounts for less than 0.0001% of all hydrogen atoms. Tritium has two neutrons in its nucleus and is radioactive, undergoing β- decay to form helium-3 (3He).
It is rarest isotope, heavier than deuterium because of the number of the nucleus.
It has a mass number of 3 and is also known as super-heavy hydrogen. Tritium is unstable and undergoes radioactive decay with a half-life of about 12 years.
Tritium is used in research and industry as a tracer and in the production of self-luminous paints, light sources, and other devices that require low-energy radiation.
The isotopes of hydrogen have similar chemical properties, but their physical properties, such as density and melting point, differ significantly due to their atomic mass.
| Isotope | Symbol | Atomic Number | Mass Number | Natural Abundance (%) |
|---|---|---|---|---|
| Protium | H | 1 | 1 | 99.985 |
| Deuterium | D | 1 | 2 | 0.015 |
| Tritium | T | 1 | 3 | <0.0001 |
Atomic Properties of Hydrogen
The atomic structure of hydrogen consists of one electron and one proton. The proton is located in the center of the atom, called the nucleus, while the electron orbits around the nucleus.
Electronegativity: Hydrogen has an electronegativity of 2.20 on the Pauling scale, which is intermediate between the values for metals and non-metals.
Ionization Energy: The ionization energy of hydrogen is relatively low, at 1312 kJ/mol. This means that it takes relatively little energy to remove the electron from a hydrogen atom.
Electron Affinity: Hydrogen has a low electron affinity, meaning that it does not readily gain electrons to form negative ions.
Bonding of Hydrogen
Bonding of hydrogen is particularly interesting due to its unique electronic configuration, having only one electron in its outermost shell, the 1s orbital. This leads to different types of bonding and bond angles depending on the other elements involved.
Neutral hydrogen can make a single covalent bond involving the overlap of the 1s orbital with some other orbital to produce a sigma bond, which can occur in any direction. The electronic configuration of hydrogen is significant in understanding its bonding behavior and reactivity. Hydrogen has a high ionization energy, which makes it difficult to remove its only electron. Consequently, hydrogen prefers to share its electron to complete its outermost shell, resulting in a covalent bond.
Types of Bonding of Hydrogen (H2)
Hydrogen can form different types of bonds depending on the electronegativity of the other elements involved. The electronegativity of an element is the measure of its ability to attract electrons towards itself. Hydrogen can form ionic, covalent, and metallic bonds.
Ionic Bonding
Ionic bonding occurs when there is a complete transfer of electrons from one atom to another, resulting in the formation of positively and negatively charged ions that are held together by electrostatic attraction. Hydrogen can form ionic bonds with highly electronegative elements such as fluorine and oxygen.
For example, in hydrogen fluoride (HF), hydrogen forms an ionic bond with fluorine.
H2 + F2 → 2HF
Hydrogen loses its electron to form a positively charged ion, H+, while fluorine gains the electron to form a negatively charged ion, F–. The bond length in HF is 0.92 Å, and the bond angle is 105.9°.
The bond angle is due to the fact that fluorine has a small atomic radius, causing the electrons in the molecule to be tightly held, resulting in a slightly compressed angle.
Covalent Bonding
Covalent bonding occurs when atoms share electrons to form a stable electron configuration. Hydrogen forms covalent bonds with other non-metallic elements such as carbon, nitrogen, and oxygen. Covalent bonding can be polar or nonpolar depending on the electronegativity difference between the atoms.
In a nonpolar covalent bond, the electrons are shared equally between the two atoms.
An example of nonpolar covalent bonding of hydrogen is in the diatomic molecule, H2. The bond length in H2 is 0.74 Å, and the bond angle is 180°.
In a polar covalent bond, the electrons are not shared equally between the two atoms, resulting in partial charges on the atoms.
An example of polar covalent bonding of hydrogen is in water (H2O). In water, the oxygen atom is more electronegative than the hydrogen atoms, resulting in partial negative charges on the oxygen atom and partial positive charges on the hydrogen atoms. The bond length in water is 0.96 Å, and the bond angle is 104.5°.
Metallic Bonding
Metallic bonding occurs in metals when the outermost electrons of metal atoms delocalize and form a “sea” of electrons that are shared by all the atoms. Hydrogen can form metallic bonds with other metals such as palladium and platinum.
For example, in the metal hydride, palladium hydride (PdH), hydrogen forms a metallic bond with palladium. The bond length in PdH is 2.06 Å, and the bond angle is not applicable in metallic bonding.
Reactions of Hydrogen
Hydrogen is highly reactive, and it undergoes various chemical reactions such as combustion, reduction, and oxidation. The chemical reactivity of hydrogen can be attributed also to its electronic configuration, which enables it to form a single covalent bond with other atoms or molecules.
Combustion of Hydrogen
Hydrogen combustion is an exothermic reaction that releases a large amount of heat and light asides the product (water). The balanced equation for the combustion of hydrogen is:
2H2(g) + O2(g) → 2H2O(l) + energy
The reaction requires a spark or flame to initiate the reaction, after which it becomes self-sustaining. Hydrogen combustion is used in fuel cells to generate electricity, and it is also used as a fuel in rockets.
Reduction of Hydrogen
Hydrogen is an excellent reducing agent and is commonly used in reduction reactions. Reduction is the gain of electrons, and hydrogen can donate its electron to other atoms or molecules while it is reduced in a reduction reaction.
For example, the reaction of hydrogen gas with a metal ion, such as iron (III) ion, to form hydrogen gas and the reduced metal ion, iron (II):
2 Fe3+ + H2 → 2 Fe2+ + 2 H+
In this reaction, hydrogen gas is oxidized by the iron (III) ion to form water and the iron (II) ion, which has a lower oxidation state than the iron (III) ion.
Another example of the reduction of hydrogen is the reaction of hydrogen gas with a non-metal element, such as sulfur, to form hydrogen sulfide:
H2 + S → H2S
In this reaction, hydrogen gas is oxidized by sulfur to form hydrogen sulfide, which is a compound with a lower oxidation state than hydrogen gas.
The reduction of hydrogen can also occur in organic chemistry, where hydrogen atoms on a molecule are replaced with a reducing agent, such as lithium aluminum hydride (LiAlH4), to form a reduced product:
RCH2OH + LiAlH4 → RCH2OH2 + Al(OH)3 + LiH
In this reaction, the hydrogen atom on the alcohol functional group is replaced by a hydride ion (H–) from the LiAlH4, resulting in a reduced alcohol product.
Oxidation of Hydrogen
Hydrogen can also undergo oxidation reactions, where it loses electrons.
The oxidation of hydrogen occurs when hydrogen reacts with an oxidizing agent, such as oxygen or chlorine, to form water or hydrogen chloride respectively. The reaction can be represented by the following balanced chemical equation:
2H2 + O2 → 2H2O
In this reaction, two molecules of hydrogen react with one molecule of oxygen to produce two molecules of water. Each hydrogen atom loses one electron to the oxygen atom, resulting in oxidation of the hydrogen.
Similarly, the reaction of hydrogen with chlorine results in the formation of hydrogen chloride:
H2 + Cl2 → 2HCl
In this reaction, one molecule of hydrogen reacts with one molecule of chlorine to produce two molecules of hydrogen chloride. Each hydrogen atom again loses one electron to the chlorine atom, resulting in oxidation of the hydrogen.
Reaction with Metals
Hydrogen can also react with metals to form metal hydrides. For example, when hydrogen reacts with sodium, it forms sodium hydride:
2Na(s) + H2(g) → 2NaH(s)
Sodium hydride is a white solid that is used as a reducing agent in organic chemistry. It can also react with water to form sodium hydroxide and hydrogen gas:
NaH(s) + H2O(l) → NaOH(aq) + H2(g)
Reaction with Non-Metal
Hydrogen can also react with non-metals to form covalent compounds. Hydrogen react with non-metals such as oxygen, sulfur, and nitrogen to form a variety of compounds. These reactions can be represented by the following balanced chemical equations:
Reaction of hydrogen with oxygen had been mentioned in a previous section, to produce water, is a highly exothermic reaction and the basis for combustion reactions involving hydrogen.
Also, the reaction of hydrogen with sulfur:
H2 + S → H2S
In this reaction, one molecule of hydrogen reacts with one molecule of sulfur to produce one molecule of hydrogen sulfide. This reaction is also exothermic and can be used to produce hydrogen sulfide gas in the laboratory.
The reaction of hydrogen with nitrogen give ammonia:
3H2 + N2 → 2NH3
This reaction is key in the production of ammonia, which is an important industrial chemical.
Applications of Hydrogen
Hydrogen has numerous applications in various fields, such as chemical processing, fuel cells, and aerospace. Its unique chemical properties make it an excellent fuel and reducing agent.
Chemical Processing
Hydrogen is used in the production of numerous chemicals, including ammonia, methanol, and hydrogen peroxide. It is also used in the refining of petroleum and the production of high-quality steel.
Fuel Cells
Hydrogen fuel cells are devices that convert the chemical energy of hydrogen into electrical energy. The reaction involves the oxidation of hydrogen to form water, with the release of electrons, which are used to generate electricity. Fuel cells are used in transportation, such as cars and buses, and in stationary power applications, such as backup power systems for buildings and telecommunication towers.
Aerospace
Hydrogen is used as a rocket propellant because of its high specific impulse, which means that it can generate a large amount of thrust per unit mass of propellant. The Space Shuttle main engine and the Saturn V rocket used liquid hydrogen as a propellant.
Compounds of Hydrogen
Binary Compounds of Hydrogen
The binary compounds of hydrogen are called hydrides. These are compounds that contain hydrogen bonded to a metal or a non-metal. There are three types of hydrides: ionic, covalent, and metallic.
Ionic Hydrides
Ionic hydrides are a type of inorganic hydride that forms when an alkali metal or alkaline earth metal reacts with hydrogen. These hydrides are composed of ionic bonds between the metal cation and the hydride anion. Generally, they are white crystalline solids with high melting and boiling points and are often used as reducing agents in organic synthesis.
There are two types of ionic hydrides: saline hydrides and saline-earth hydrides. Saline hydrides are formed when an alkali metal reacts with hydrogen, while saline-earth hydrides are formed when an alkaline earth metal reacts with hydrogen.
Saline hydrides, such as sodium hydride (NaH) and lithium hydride (LiH), are highly reactive and often used as reducing agents in organic synthesis.
Generally, they react with water to give hydroxide and water:
MH + H2O → MOH + H2, where M is the metal cation.
Examples:
1. 2NaH + 2 H2O → 2 NaOH + H2
2. 2KH + 2H2O → 2KOH + H2
3. LiH + H2O → LiOH + H2
Saline-earth hydrides, such as magnesium hydride (MgH2) and calcium hydride (CaH2), similarly react with water but less than saline hydrides.
EH2 + 2H2O à E(OH)2 + 2H2 where E = Mg, Ca
They are usually white crystalline solids with high melting and boiling points, and they are highly reactive with water and other protic solvents. They are also strong reducing agents, and they can react with many organic compounds.
MH + H2O → MOH + H2, where M is the metal cation.
Examples:
1. 2NaH + 2 H2O → 2 NaOH + H2
2. 2KH + 2H2O → 2KOH + H2
3. LiH + H2O → LiOH + H2
4. EH2 + 2H2O à E(OH)2 + 2H2 where E = Mg, Ca
Covalent Hydrides
Covalent hydrides are formed by the direct combination of hydrogen with non-metals. These hydrides are typically gases or liquids at room temperature and have low boiling points.
Covalent hydrides are classified into two groups: molecular hydrides and network covalent hydrides.
Molecular covalent hydrides are typically gases or liquids at room temperature and pressure and have low boiling points. These hydrides are held together by covalent bonds between the hydrogen and non-metal atoms. Examples of molecular covalent hydrides include methane (CH4), ammonia (NH3), and water (H2O).
The properties of molecular covalent hydrides are primarily determined by the type and strength of the intermolecular forces present. These intermolecular forces are relatively weak compared to the covalent bonds within the molecules, and as a result, molecular covalent hydrides typically have low melting and boiling points. They are also generally insoluble in water and have low electrical conductivity.
The reactions of molecular covalent hydrides are determined by the types of functional groups present in the molecules. For example, water can undergo acid-base reactions with other molecules or ions. The balanced chemical equation for the reaction between hydrochloric acid and water is:
HCl + H2O → H3O+ + Cl–
Ammonia can act as a weak base and undergo acid-base reactions with stronger acids, such as hydrochloric acid. The balanced chemical equation for the reaction between ammonia and hydrochloric acid is:
NH3 + HCl → NH4Cl
Network covalent hydrides are characterized by their high melting and boiling points, as well as their insolubility in water. These hydrides are held together by a network of covalent bonds between the hydrogen and nonmetal atoms, resulting in a giant, three-dimensional structure.
Examples of network covalent hydrides include silicon hydride (SiH4), and boron hydride (BH3).
The properties of network covalent hydrides are largely determined by the strength and organization of the covalent bonds in the network. The covalent bonds are very strong, which results in high melting and boiling points. In addition, the network structure means that network covalent hydrides are typically insoluble in water and have low electrical conductivity.
The reactions of network covalent hydrides are determined by the types of functional groups present in the molecules. For example, boron hydride can act as a Lewis acid, accepting electron pairs from other molecules or ions.
BH3 + NH3 → H3BNH3
Metal Hydrides
Metal hydrides are formed by the absorption of hydrogen by metals. These hydrides are usually solids and are used as hydrogen storage materials. Metallic hydrides can be classified into two groups: interstitial hydrides and metal complex hydrides.
Interstitial Metal Hydrides
Interstitial metal hydrides are formed when hydrogen atoms occupy the interstitial sites in a metal lattice. These types of metal hydrides are typically covalently bonded, with the metal atom donating electron density to the hydrogen atom. The properties of interstitial metal hydrides depend on the nature of the metal and the hydrogen atom.
For example, titanium hydride (TiH2) is formed as showed in this chemical equation:
Ti + H2 → TiH2
TiH2 has a high melting point and is used as a deoxidizing agent in the steel industry.
The reaction between interstitial metal hydrides and water can produce hydrogen gas and the corresponding metal hydroxide. This same titanium hydride reacts water and can be represented by the following balanced chemical equation:
TiH2 + 2H2O → Ti(OH)4 + 2H2
Complex Metal Hydrides
Complex metal hydrides are formed when hydrogen atoms are bonded to a metal ion through covalent and ionic interactions. These types of metal hydrides are typically used as hydrogen storage materials due to their high hydrogen content. Examples of complex metal hydrides include sodium borohydride (NaBH4) and lithium aluminum hydride (LiAlH4).
The formation of sodium borohydride can be represented by:
NaH + BH3 → NaBH4
Complex metal hydrides can also react with water, producing hydrogen gas and the corresponding metal hydroxide. For example, the reaction between sodium borohydride and water releases hydrogen gas.
NaBH4 + 2H2O → NaBO2 + 4H2
Complex metal hydrides can also react with acids, producing hydrogen gas and the corresponding metal salt. For example, the reaction between lithium aluminum hydride reacts hydrochloric acid:
LiAlH4 + 4HCl → LiCl + AlCl3 + 4H2
| Type of Hydride | Properties | Reactions | Preparation Method |
|---|---|---|---|
| Ionic Hydrides | High melting and boiling points, brittle, non-conductive | Release hydrogen gas upon reaction with water, strong reducing agents | Reaction between an alkali metal and hydrogen gas, or reaction between a metal hydride and an alkali metal |
| Covalent Hydrides | Low melting and boiling points, generally insoluble in water, good conductors of electricity in the liquid or molten state | Reactions depend on the nature of the covalent bond, some covalent hydrides react with water to produce acidic solutions, while others are inert | Prepared by direct combination of the elements, or by reaction of a metal hydride with a nonmetallic element or compound |
| Metal Hydrides | Variable melting and boiling points, typically insoluble in water, can be highly reactive with water or acids | Can react with acids to release hydrogen gas, or with water to form metal hydroxides and hydrogen gas | Prepared by reaction between a metal and hydrogen gas, or by reaction between a metal salt and a reducing agent in the presence of hydrogen gas |
Ternary Compounds of Hydrogen
These compounds contain hydrogen as well as two other elements. These compounds can be classified into three main categories: salts, acids, and other complex compounds.
Salts of Ternary Compounds of Hydrogen
These salts of ternary compounds of hydrogen are formed when hydrogen in an acid is replaced by a metal or other positively charged ion. Some common examples of salts of ternary compounds of hydrogen include sodium bicarbonate (NaHCO3), and sodium hydrogen phosphate (Na2HPO4).
These salts are typically prepared by reacting an acid with a metal or other positively charged ion. For example, sodium bicarbonate is produced by reacting carbon dioxide and sodium hydroxide:
CO2 + NaOH → NaHCO3
Generally, salts of these ternary compounds are solid, crystalline compounds that are soluble in water. They have characteristic melting and boiling points, and they can conduct electricity when dissolved in water.
They react with acids to form other salts, or with bases to form hydroxides. They can also undergo redox reactions with other substances.
Acids of Ternary Compounds of Hydrogen
Some acids are ternary compounds of hydrogen and are formed when hydrogen is combined with a non-metal or other negatively charged ion. Examples of these acids include sulfuric acid (H2SO4), nitric acid (HNO3), and phosphoric acid (H3PO4).
Preparation: Acids are prepared from the reaction between water and acidic oxide. For example, sulfuric acid is produced by reacting sulfur trioxide with water:
SO3 + H2O → H2SO4
Acids of ternary compounds of hydrogen can react with bases to form salts, and they can also undergo redox reactions with other substances.
Other Complex Compounds of Ternary Compounds of Hydrogen
These are formed when hydrogen is combined with two other elements to form a complex compound. Some common examples of other complex compounds include ammonium sulfate ((NH4)2SO4), and sodium bicarbonate (NaHCO3).
Ammonium sulfate is produced by combining ammonia and sulfuric acid:
2NH3 + H2SO4 → (NH4)2SO4
These compounds can be solid or liquid, and they may have a range of solubilities and melting points.
They can undergo redox reactions, acid-base reactions, and other types of chemical reactions.
