User:Sad Rist/Chemical functions

In chemistry, the group of some compound substances that possess similar chemical properties, known as functional properties, is called a chemical function. When a particular compound has characteristics such as acidity or basicity, solubility in water, reactivity according to a specific chemical function, it is said to belong to that chemical function. Chemical functions are divided according to the classical division of chemistry.

There are four types of inorganic functions: oxide, acid, base, and salt. The classification criterion for a substance into one of these functions is the type of ions formed when it is dissolved in water.

Based on the non-existent nature of chemical compounds, functions can primarily be divided into inorganic functions, which are the functions of compounds that do not have a carbon chain, which is the main characteristic of these compounds. They are further divided into acids, bases, salts, and oxides, and organic functions, which are related to organic compounds.

Oxides
Oxides are binary compounds in which the element oxygen has an oxidation number of -2 and is the most electronegative element in the formula, making it the most powerful element.

Classification of oxides
Note: Amphoteric oxides behave as basic oxides in the presence of acids and as acidic oxides in the presence of bases.

Note 2: Mixed oxides are the "sum" of the oxides formed by an element, in other words, it is a cloud with all the types of oxides of that element:


 * FeO + Fe2O3 → Fe3O4

Naming of oxides
Oxides are binary compounds of oxygen with another element.

Oxides are divided into two groups according to the nature of the element that combines with oxygen. If the element is metallic, the resulting oxide is called a basic oxide or simply an oxide.

When the element that combines with oxygen is non-metallic, the compound formed is an acidic oxide or anhydride. This classification is based on the behavior of oxides with water. Basic oxides react with water to produce basic solutions or hydroxides, which neutralize acids, for example:

CaO(s) + H2O(l) ———› Ca(OH)2(s)

Many acidic oxides dissolve in water to form acidic solutions, for example:

SO3(g) + H2O(l) ———› H2SO4(l)

The formed sulfuric acid and similar solutions neutralize bases or hydroxides. Some metals located in the periodic table from group IV onwards often form basic oxides or acidic oxides, which is why they are called indifferent or amphoteric oxides. In the presence of acids, they behave like basic oxides, and in the presence of bases, they behave like acidic oxides.

Ex: ZnO(s) + 2H+(g) ———› 2Zn+(s) + H2O(l) ZnO(s) + 2OH-(aq) + H2O(l) ———› Zn(OH)2-4 (s)

Observed that in the first reaction, the oxide acts as a base (hydroxide), and in the second reaction, it acts as an acid.

In general terms, the nomenclature of basic oxides uses the generic term "oxide" followed by the name of the metal present in the compound's formation.

Example: Na2O is sodium oxide.

When a metal forms more than one oxide, for nomenclature, the lower oxidation state (valence) is considered, and the metal's name ends in -ous, and for the higher oxidation state, it ends in -ic.

Iron also forms two oxides: Fe2O is ferrous oxide (These oxides with water produce hydroxides) Fe2O3 is ferric oxide (These oxides with water produce hydroxides)

The oxide Fe3O4, ferrous ferric oxide, is considered to be formed by the combination of the two previous oxides; it is a saline oxide.

The nomenclature of acidic oxides or anhydrides is similar to that of basic oxides, using the generic term "anhydride" followed by the specific name of the non-metal.

CO2 is carbon dioxide SO2 is sulfur dioxide

Here, too, there are non-metals that form more than one anhydride. Their nomenclature follows the same rules mentioned earlier. Additionally, if there are three oxidation states for the same number of anhydrides of the same non-metal, the prefix "hypo-" is added to the -ous ending of the specific word for the lower oxidation state.

Br2O is hypobromous anhydride

Br2O3 is bromous anhydride

Br2O5 is bromic anhydride

SO is hyposulfurous anhydrid

SO2 is sulfurous anhydride

SO3 is sulfuric anhydride

For elements that have four oxidation states (continue editing)

Acids
According to Arrhenius, an acid is any substance that releases an H+ ion in water or, more precisely, substances that dissociate in an aqueous medium, releasing the cation H+. The current theory of Brønsted-Lowry defines an acid as a substance capable of accepting a pair of electrons. Furthermore, Arrhenius's theory has also been updated:

An acid is any substance that releases an H3O+ ion.

Classification of acids

a) According to the presence of oxygen

Hydrides: They do not contain oxygen in their formula. Examples: HI, HCl, HF

Oxyacids: They contain oxygen in their formula. Examples: H2CO3, H2SO3, H2SO4, HNO2

b) According to the degree of ionic dissociation

Note: The calculation of α in acids is the same as that in bases.


 * α (in percentage) = 100 x number of dissociated molecules/total number of dissolved molecules


 * α > 50% → strong
 * α < 5% → weak

Hydrides:

Strong: HCl < HBr < HI Medium: HF (can be considered weak) Weak: the rest Oxyacids:

Strong: when x > 1 (H2SO4) Medium: when x = 1 (HClO2) Weak: when x < 1 (HClO) x = number of oxygen - number of hydrogen

Nomenclature of acids
a) Hydrides


 * acid + element + hydric

Examples:


 * HI is hydriodic acid


 * HCl is hydrochloric acid


 * H2S is sulfhydric acid

b) Oxyacids

Since they can be obtained through the hydration of acidic oxides, the nomenclature follows the same system.


 * acid + prefix + element + suffix

Note: The lower the oxygen, the lower the oxidation number of the central element, and the more oxygen, the higher the oxidation number, as shown in the examples below.

Examples:


 * HClO is hypochlorous acid (nox Cl = +1)


 * HClO2 is chlorous acid (nox Cl = +3)


 * HClO3 is chloric acid (nox Cl = +5)


 * HClO4 is perchloric acid (nox Cl = +7)

Bases
Bases, according to Arrhenius, are compounds that dissociate in water, releasing the hydroxide anion (OH-) and a cation other than H+. The current theory of Lewis defines a base as a substance capable of donating a pair of electrons.

Classification of bases
A) According to the degree of dissociation

The same calculation used for acids is used for bases.

Strong: α = 100% → Bases formed by alkali metals and alkaline earth metals. When the degree of ionization is practically 100%. This is the case for hydroxides of alkali metals and alkaline earth metals, which are inherently ionic.

Weak: α < 5% → Their degree of ionization is generally less than 5%. This includes ammonium hydroxide and hydroxides of metals in general, excluding alkali metals and alkaline earth metals, which are molecular by nature.

Nomenclature of bases
A) When the cation has a fixed oxidation number


 * hydroxide + cation

Examples:


 * NaOH is sodium hydroxide


 * KOH is potassium hydroxide

B) When the cation does not have a fixed oxidation number


 * hydroxide + cation + suffix O

or
 * hydroxide + cation + oxidation number in Roman numerals

Examples:


 * Fe(OH)2 is iron(II) hydroxide or ferrous hydroxide


 * Fe(OH)3 is iron(III) hydroxide or ferric hydroxide

Salts
Salts are compounds that dissociate in water, releasing at least one cation other than H+ and at least one anion other than OH-. They are defined, in a limited sense, as binary compounds resulting from the reaction of an acid and a base.

Note: When they are dissolved in water, their dissociated ions become mobile and become conductors of electricity.

Classification of salts
a) According to the presence of oxygen

Halide salts: They do not contain oxygen. Examples: NaI, KBr

Oxisalts: They contain oxygen. Examples: CaCO3, MgSO4

b) According to the presence of H+ or OH-

Normal salts: They are formed by complete neutralization between an acid and a base. They do not contain H+ or OH-. Example: HCl + NaOH → NaCl + H2O

Acid salts or hydroxysalts: They are formed in a neutralization reaction when the acid and base are not in stoichiometric proportion. In such cases, there is partial neutralization, leaving excess H+ or OH-. Example (acid salt): H2CO3 + NaOH → NaHCO3 + H2O Example (hydroxysalt): Mg(OH)2 + HCl → Mg(OH)Cl + H2O

Mixed salts: The salt contains more than one cation or more than one anion. It is formed from the neutralization of an acid by more than one base or a base by more than one acid. Example: Al(OH)3 + HCl + H2SO4 → AlClSO4 + 3H2O

Nomenclature of salts
a) For halide salts


 * metal + -ide

Example: NaCl is sodium chloride

b) For oxisalts

We use an extension of the table of acidic oxides and oxyacids because the nomenclature of oxisalts also depends on the oxidation state.

Exceptions: Boron (B) +3, Carbon (C) +4, and Silicon (Si) +4 only have the -ic suffix in the form of acids. When they form salts, the -ate suffix is always used.

Examples:

KNO2 (nox N = +3) is potassium nitrite

NaClO (nox Cl = +1) is sodium hypochlorite

KMnO4 (nox Mn = +7) is potassium permanganate

Note: When there is hydrogen in the salt formula, the prefix "bi-" is added to the name of the cation.

NaHCO3 is sodium bicarbonate