Name: how bonding affects acid strength
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 5 s
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Acid Strength and Molecular Structure

L'acido cloridrico è un acido forte, mentre l'acido fluoridrico è un acido debole. Ma cosa determina la loro forza? La forza degli acidi binari, con solo due elementi, è determinata dall'energia di legame e dalla polarità di legame.

Un acido con un'energia di legame più alta ha un legame più forte e, quindi, sarà un acido più debole. Paragonando gli acidi in un gruppo, il legame in un acido debole, come l'acido fluoridrico, è più difficile da rompere, quindi è meno probabile che l'acido doni protoni. Al contrario, un acido con un'energia di legame inferiore ha un legame più debole e, quindi, sarà un acido più forte.

Per esempio, il legame in un acido forte, come l'acido cloridrico, si rompe più facilmente e dona protoni più prontamente dell'acido fluoridrico. Un legame, come quello dell'acido cloridrico, è polare quando un atomo è più elettronegativo dell'altro. Un composto può agire come un acido quando l'atomo attaccato all'idrogeno ha un'elettronegatività maggiore dell'idrogeno.

L'atomo avrà una carica negativa parziale, consentendo all'idrogeno di avere una carica positiva parziale cosicché possa essere rilasciato come protone. Un acido con una polarità di legame più alta ha un legame più debole e, quindi, sarà un acido più forte. Confrontando i composti in un periodo, l'acido cloridrico è più forte dell'idrogeno solforato poiché il cloro è più elettronegativo dello zolfo e, quindi, rilascia protoni più facilmente.

Se l'idrogeno ha un'elettronegatività uguale o superiore dell'altro atomo, quella molecola non può donare protoni e, quindi, non può agire come un acido. Gli ossiacidi sono acidi in cui un OH è attaccato a un terzo atomo che è più elettronegativo dell'idrogeno. La forza di un ossiacido dipende dall'elettronegatività e dal numero di molecole di ossigeno attaccati al terzo atomo.

Maggiore è l'elettronegatività dell'atomo, più si polarizza, indebolendo il legame tra l'ossigeno e l'idrogeno. Se l'atomo centrale è attaccato ad atomi di ossigeno aggiuntivi, aumenta ulteriormente la polarità del legame fra l'ossigeno e l'idrogeno. Per esempio, l'acido perclorico con tre atomi di ossigeno aggiuntivi è più forte dell'acido clorico con due atomi di ossigeno aggiuntivi.

L'acido clorico, da parte sua, è più forte dell'acido cloridrico, che ha solo un ossigeno aggiuntivo, e acido ipocloroso, senza atomi di ossigeno aggiuntivi. Gli acidi carbossilici sono acidi deboli che contengono un gruppo carbossilico. Il secondo atomo di ossigeno rende il legame ossigeno-idrogeno più polare e pertanto consente alla molecola di donare un protone.

L'acido acetico e l'acido formico sono esempi di acidi carbossilici.

forza acida e struttura molecolare

Acidi e basi binari

In assenza di qualsiasi effetto livellante, la forza acida dei composti binari dell’idrogeno con i non metalli (A) aumenta man mano che la forza del legame H-A diminuisce lungo un gruppo nella tavola periodica. Per il gruppo 17, l'ordine crescente di acidità è HF < HCl < HBr < HI. Allo stesso modo, per il gruppo 16, l'ordine crescente di forza acida è H_2O < H_2S < H_2Se < H2Te. Attraverso una riga nella tavola periodica, la forza acida dei composti binari dell'idrogeno aumenta con l'aumentare dell'elettronegatività dell'atomo non metallico perché aumenta la polarità del legame H-A. Pertanto, l'ordine di acidità crescente (per la rimozione di un protone) attraverso la seconda fila è CH_4 < NH_3 < H_2O < HF; nella terza riga è SiH_4 < PH_3 < H_2S < HCl.

Acidi e basi ternari

I composti ternari formati da idrogeno, ossigeno e qualche terzo elemento ("E") possono essere strutturati come illustrato nell'immagine seguente. In questi composti, l'atomo E centrale è legato a uno o più atomi di O, e almeno uno degli atomi di O è anche legato a un atomo di H, corrispondente alla formula molecolare generale O_mE(OH)_n. Questi composti possono essere acidi, basici o anfoteri a seconda delle proprietà dell'atomo E centrale. Esempi di tali composti includono acido solforico, O_2S(OH)_2, acido solforoso, OS(OH)_2, acido nitrico, O_2NOH, acido perclorico, O_3ClOH, idrossido di alluminio, Al(OH)_3, idrossido di calcio, Ca(OH)_2 e idrossido di potassio, KOH.

<img alt="Eq1" src="/files/ftp_upload/11406/11406_Figure_1.svg" style="height: 90px;" />

Se l'atomo centrale, E, ha una bassa elettronegatività, la sua attrazione per gli elettroni è bassa. Esiste una scarsa tendenza da parte dell'atomo centrale a formare un forte legame covalente con l'atomo di ossigeno, e il legame a tra l'elemento e l'ossigeno si rompe più facilmente del legame b tra ossigeno e idrogeno. Quindi il legame a è ionico, gli ioni idrossido vengono rilasciati nella soluzione e il materiale si comporta come una base: questo è il caso di Ca(OH)_2 e KOH. Una minore elettronegatività è caratteristica degli elementi più metallici; quindi gli elementi metallici formano idrossidi ionici che sono, per definizione, composti basici.

Se invece l'atomo E ha un'elettronegatività relativamente elevata, attrae fortemente gli elettroni che condivide con l'atomo di ossigeno, formando un legame covalente relativamente forte. Il legame ossigeno-idrogeno, legame b, viene quindi indebolito perché gli elettroni vengono spostati verso E. Il legame b è polare e rilascia facilmente ioni idrogeno nella soluzione, quindi il materiale si comporta come un acido. Elevate elettronegatività sono caratteristiche degli elementi più non metallici. Pertanto, gli elementi non metallici formano composti covalenti contenenti gruppi −OH acidi chiamati ossiacidi.

L’aumento del numero di ossidazione dell’atomo centrale E aumenta anche l’acidità di un ossiacido perché ciò aumenta l’attrazione di E per gli elettroni che condivide con l’ossigeno e quindi indebolisce il legame OH. L'acido solforico, H_2SO_4 o O_2S(OH)_2 (con un numero di ossidazione dello zolfo pari a +6), è più acido dell'acido solforoso, H_2SO_3 o OS(OH)2 (con un numero di ossidazione dello zolfo pari a +4). Allo stesso modo, l'acido nitrico, HNO_3 o O_2NOH (numero di ossidazione N = +5), è più acido dell'acido nitroso, HNO_2 o ONOH (numero di ossidazione N = +3). In ciascuna di queste coppie, il numero di ossidazione dell'atomo centrale è maggiore per l'acido più forte.

Acidi carbossilici

Gli acidi carbossilici contengono un gruppo carbossilico. Gli acidi carbossilici sono acidi deboli, il che significa che non sono ionizzati al 100% in acqua.

L'acido carbossilico agisce come un acido debole perché, come nel caso degli ossiacidi, il secondo ossigeno attaccato all'atomo di carbonio aumenta la polarità del legame OH e lo rende più debole. Inoltre, dopo la perdita del protone, il gruppo carbossilico viene convertito nello ione carbossilato, che presenta risonanza. Le diverse strutture di risonanza stabilizzano lo ione carbossilato poiché la sua carica negativa è delocalizzata su più atomi.



Questo testo è adattato da<a href="https://openstax.org/books/chemistry-2e/pages/14-3-relative-strengths-of-acids-and-bases"> Openstax, Chemistry 2e, Section 14.3: Relative Strengths of Acids and Bases</a> and <a href="https://openstax.org/books/chemistry-2e/pages/20-3-aldehydes-ketones-carboxylic-acids-and-esters">Openstax, Chemistry 2e, Section 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters</a>.

Forza acida e struttura molecolare

Acidi e basi binarie
In assenza di qualsiasi effetto livellante, la forza acida dei composti binari dell'idrogeno con i non metalli (A) aumenta con l'aumento ...

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how bonding affects acid strength

Acid strength depends on bonding, electronegativity, and oxidation number. These factors help explain why some hydrogen compounds act as strong acids, while others behave as weak acids or bases.

For binary acids, which are hydrogen compounds with one other nonmetal, acidity increases as the H-A bond gets weaker down a group. That is why the order for group 17 is HF &lt; HCl &lt; HBr &lt; HI, and for group 16 it is H2O &lt; H2S &lt; H2Se &lt; H2Te. Across a period, acid strength increases as the nonmetal becomes more electronegative because the H-A bond becomes more polar. The second row follows CH4 &lt; NH3 &lt; H2O &lt; HF, and the third row follows SiH4 &lt; PH3 &lt; H2S &lt; HCl.

Ternary acids and bases contain hydrogen, oxygen, and a third element E. Their general form is OmE(OH)n, and they may be acidic, basic, or amphoteric. Examples include sulfuric acid, sulfurous acid, nitric acid, perchloric acid, aluminum hydroxide, calcium hydroxide, and potassium hydroxide.

If E has low electronegativity, it does not hold electrons strongly. In that case, the bond between E and oxygen is more likely to break than the O-H bond, so hydroxide ions are released and the compound acts as a base. This is typical of more metallic elements, which form ionic hydroxides such as Ca(OH)2 and KOH.

If E has high electronegativity, it pulls electron density toward itself. That weakens the O-H bond and makes it easier to release H+ ions, so the compound behaves as an acid. These nonmetal compounds are called oxyacids. A higher oxidation number on E also increases oxyacid strength because it pulls harder on shared electrons. For example, H2SO4 is more acidic than H2SO3, and HNO3 is more acidic than HNO2.

Carboxylic acids also show weak-acid behavior. They contain a carboxyl group, and they are not fully ionized in water. The second oxygen attached to carbon makes the O-H bond more polar and weaker. After the proton leaves, the carboxylate ion is stabilized by resonance, which spreads the negative charge over several atoms.

How Bonding Affects Acid Strength

bronsted-lowry acids and bases

Proton Transfer in Acids and Bases

Acids and bases can be understood through proton transfer. Early scientists described acids by their sour taste, their effect on dyes, and their reaction with limestone, which releases carbon dioxide. Robert Boyle also noted that acid solutions can lose these properties after contact with alkali, or base, solutions. The modern view developed step by step. Humphry Davy showed that hydrogen is the essential part of acids. Joseph Louis Gay-Lussac then described acids as substances that neutralize...

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Water as an Amphiprotic Substance

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pH and pOH in Water

pH and pOH show how hydronium and hydroxide ions behave in water and other aqueous solutions. These ion concentrations are linked by the ion product of water, K w , so when one goes up, the other goes down. Because ion concentrations can differ by huge amounts, a logarithmic scale is used to describe them. pH is defined from the hydronium ion concentration, [H3O+], and pOH is the matching p-function for hydroxide ion concentration. These two values are also connected through the K w...

relative strengths of conjugate acid-base pairs

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How Ka and Kb Show Acid-Base Strength

Brønsted-Lowry acid-base chemistry is based on proton transfer. That idea helps show how the strength of an acid is linked to the strength of its conjugate base, and vice versa. Acid and base strength is measured with ionization constants, Ka and Kb, which tell how far an ionization reaction goes. For the conjugate pair HA and A−, the acid ionizes in water and the base reaction can be written for the partner species. When the two reactions are added, they give the autoionization of water. The...

strong acid and base solutions

Strong Acids, Strong Bases, and pH

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weak acid solutions

Weak Acid Equilibrium in Water

Weak acid equilibrium in water explains why many acids only partially react when dissolved. Most weak acid molecules stay in their original form, and only a small amount of hydronium ions forms. Weak acids are common in nature. They help create the tangy taste of citrus fruits, the stinging sensation from insect bites, and the smells linked to body odor. Acetic acid is a familiar weak acid. It is the main ingredient in vinegar. Its equation uses a double arrow to show that the reaction with...

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Weak Base pH and Kb Calculations

Weak base pH and Kb calculations show how partially reacting bases change water chemistry. Weak bases produce hydroxide ions when they dissolve by reacting with water molecules. Because the reaction only happens partly, these compounds are classified as weak bases. Weak bases are common in nature and in technology. Ammonia is one major example. Global production of ammonia is typically well over 100 metric tons each year. It is used as an agricultural fertilizer, a starting material for...

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The pH of acid mixtures depends on the acids present, their dissociation constants, and their starting concentrations. When a solution contains two different acids, the pH can be found by choosing the method that fits their relative strengths. This makes acid mixtures a useful case for applying acid-base rules in chemistry. In a mixture of a strong acid and a weak acid, the strong acid dissociates completely and supplies almost all of the hydronium ions in the solution. The weak acid...

ions as acids and bases

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Salt hydrolysis explains why some salt solutions are acidic, basic, or neutral. Salts are ionic compounds made of cations and anions. In water, either ion may react with water through an acid or base ionization reaction. The pH of the solution depends on the relative acid-base strength of the ions in the salt. Ammonium chloride is a good example of an acidic salt. When it dissolves, it separates into ammonium ions and chloride ions. The ammonium ion is the conjugate acid of ammonia, so it can...

determining the ph of salt solutions

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acid/base strengths and dissociation constants

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Acid and Base Strength in Water

Acid and base strength in water depends on how much a substance ionizes after it dissolves. If ionization is nearly complete, the substance is called strong. If only a small amount ionizes, it is called weak. Weak acids and weak bases are much more common than strong ones. The most common strong acids are HClO4, HCl, HBr, HI, HNO3, and H2SO4. The most common strong bases are LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2. These strong acids and bases ionize essentially completely in water. <br...

polyprotic acids

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acid strength and molecular structure

Acid strength depends on bonding, electronegativity, and oxidation number. These factors help explain why some hydrogen compounds act as strong acids, while others behave as weak acids or bases. For binary acids, which are hydrogen compounds with one other nonmetal, acidity increases as the H-A bond gets weaker down a group. That is why the order for group 17 is HF &lt; HCl &lt; HBr &lt; HI, and for group 16 it is H2O &lt; H2S &lt; H2Se &lt; H2Te. Across a period, acid strength increases as...

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key terms and concepts

Learn how bond strength, electronegativity, oxidation number, and resonance help explain acid strength in binary acids, oxyacids, and carboxylic acids.

How Bonding Affects Acid Strength | JoVE K12

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How Bonding Affects Acid Strength

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