There are hundreds examples of acids and bases that can be found in all branches of chemistry, but as a whole are separated into two large families: inorganic and organic. Inorganic acids are usually known as mineral acids, characterized by being especially strong compared to organic ones..
Acids and bases are understood as substances that have sour or saponaceous flavors, respectively. Both are corrosive, although the word 'caustic' is often used for strong bases. Bottom line: they burn and corrode the skin if touched. Its characteristics in dissolving media have guided a series of definitions throughout history..
The image below shows the generic behavior of acids and bases when they are added or dissolved in a glass of water. Acids produce solutions with pH values below 7 due to hydronium ions, H3OR+; while the bases produce solutions with a pH above 7 due to the hydroxyl (or hydroxyl) ions, OH-.
If we add hydrochloric acid, HCl (red drop), to the glass, there will be H ions3OR+ and Cl- hydrated. On the other hand, if we repeat the experiment with sodium hydroxide, NaOH (purple drop), we will have OH ions- and Na+.
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The increasingly studied and understood characteristics of acids and bases established more than one definition for these chemical compounds. Among these definitions we have Arrhenius's, Bronsted-Lowry's, and finally Lewis's. Before citing the examples it is necessary to be clear about this.
Acids and bases, according to Arrhenius, are those that, when dissolved in water, produce H ions3OR+ or OH-, respectively. That is, the image already represents this definition. However, by itself it neglects some acids or bases too weak to produce such ions. This is where the Bronsted-Lowry definition comes in..
Bronsted-Lowry acids are those that can donate H ions+, and the bases are those that accept these H+. If an acid very easily donates its H+, means it is a strong acid. The same happens with the bases, but accepting H+.
Thus, we have strong or weak acids and bases, and their forces are measured in different solvents; especially in water, from which the known pH units are established (from 0 to 14).
Therefore, a strong acid HA will fully donate your H+ to water in a reaction of the type:
HA + HtwoO => A- + H3OR+
Where to- is the conjugate base of HA. Hence, the H3OR+ present in the glass with acid solution.
Meanwhile, a weak base B will deprotonate the water to gain its respective H+:
B + HtwoOR <=> HB + OH-
Where HB is the conjugated acid of B. This is the case of ammonia, NH3:
NH3 + HtwoOR <=> NH4+ + Oh-
A very strong base can directly donate OH ions- without the need to react with water; just like NaOH.
Finally, Lewis acids are those that gain or accept electrons, and Lewis bases are those that donate or lose electrons..
For example, the Bronsted-Lowry NH base3 It is also a Lewis base, since the nitrogen atom accepts an H+ by donating his pair of free electrons (H3N: H+). That is why the three definitions do not disagree with each other, but rather intertwine and help study acidity and basicity in a broader spectrum of chemical compounds..
Having clarified the definitions, a series of acids with their respective formulas and names will be mentioned below:
-HF: hydrofluoric acid
-HBr: hydrobromic acid
-HI: hydroiodic acid
-HtwoS: hydrogen sulfide
-HtwoSe: selenhydric acid
-HtwoTea: tellurhydric acid
These are the binary acids, also called hydracids, to which the aforementioned hydrochloric acid, HCl belongs.
-HNO3: nitric acid
-HNOtwo: nitrous acid
-HNO: hyponitrous acid
-HtwoCO3: carbonic acid
-HtwoCOtwo: carbonaceous acid, which is actually better known by the name of formic acid, HCOOH, the simplest organic acid of all
-H3PO4: phosphoric acid
-H3PO3 or Htwo[HPO3]: phosphorous acid, with an H-P bond
-H3POtwo or H [HtwoPOtwo]: hypophosphorous acid, with two H-P bonds
-HtwoSW4: sulfuric acid
-HtwoSW3: sulfurous acid
-HtwoStwoOR7: disulfuric acid
-HIO4: periodic acid
-HIO3: iodic acid
-HIOtwo: iodine acid
-HIO: hypoiodine acid
-HtwoCrO4: chromic acid
-HMnO4: manganic acid
-CH3COOH: acetic acid (vinegar)
-CH3SW3H: methanesulfonic acid
All of these acids, except formic and the last two, are known as oxacids or ternary acids..
-AlCl3: aluminum chloride
-FeCl3: ferric chloride
-BF3: boron trifluoride
-Metallic cations dissolved in water
-Carbocations
-H (CHBelevenCleleven): superacid carborane
- FSO3H: fluorosulfonic acid
- HSbF6: fluoroantimonic acid
- FSO3H SbF5: magic acid
The last four examples make up the fearsome super acids; compounds capable of disintegrating almost any material just by touching it. AlCl3 is an example of Lewis acid, since the metal center of aluminum is able to accept electrons due to its electronic deficiency (it does not complete its valence octet).
Among the inorganic bases we have metal hydroxides, such as sodium hydroxide, and some molecular hydrides, such as ammonia already mentioned. Here are other examples of bases:
-KOH: potassium hydroxide
-LiOH: lithium hydroxide
-RbOH: rubidium hydroxide
-CsOH: cesium hydroxide
-FrOH: francium hydroxide
-Be (OH)two: beryllium hydroxide
-Mg (OH)two: magnesium hydroxide
-Ca (OH)two: calcium hydroxide
-Sr (OH)two: strontium hydroxide
-Ba (OH)two: barium hydroxide
-Ra (OH)two: radium hydroxide
-Fe (OH)two: ferrous hydroxide
-Fe (OH)3: ferric hydroxide
-Al (OH)3: aluminum hydroxide
-Pb (OH)4: lead hydroxide
-Zn (OH)two: zinc hydroxide
-Cd (OH)two: cadmium hydroxide
-Cu (OH)two: cupric hydroxide
-Ti (OH)4: titanic hydroxide
-PH3: phosphine
-AsH3: arsine
-NaNHtwo: sodium amide
- C5H5N: pyridine
-(CH3) N: trimethylamine
- C6H5NHtwo: phenylamine or aniline
-NaH: sodium hydride
-KH: potassium hydride
-Carbanions
-Li3N: lithium nitride
-Alkoxides
-[(CH3)twoCH]twoNLi: lithium diisopropylamide
-Diethynylbenzene anion: C6H4C4two- (the strongest base known so far)
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