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

L'acide chlorhydrique est un acide fort, tandis que l'acide fluorhydrique est un acide faible. Mais qu'est-ce qui détermine leur force? La force des acides binaires, avec seulement deux éléments, est déterminée par l'énergie et la polarité de la liaison.Un acide à une énergie de liaison plus élevée a une liaison plus forte et, par conséquent, sera un acide plus faible. Comparant les acides dans un groupe, la liaison dans un acide faible, comme l'acide fluorhydrique, est plus difficile à briser, alors, l'acide est moins susceptible de donner des protons. En revanche, un acide à une énergie de liaison inférieure a une liaison plus faible et, par conséquent, sera un acide plus fort.Par exemple, la liaison dans un acide fort, comme l'acide chlorhydrique, se brise plus facilement et donne des protons plus facilement que l'acide fluorhydrique. Un lien, comme celui de l'acide chlorhydrique, est polaire lorsqu'un atome est plus électronégatif que l'autre atome. Un composé peut agir comme un acide lorsque l'atome attaché à l'hydrogène a une électronégativité plus élevée que l'hydrogène.L'atome aura une charge partiellement négative, permettant à l'hydrogène d'avoir une charge positive partielle de sorte qu'il peut être libéré sous forme de proton. Un acide avec une polarité de liaison plus élevée a une liaison plus faible et, par conséquent, sera un acide plus fort. Comparer des composés sur une période, l'acide chlorhydrique est plus fort que le sulfure d'hydrogène car le chlore est plus électronégatif que le soufre et, par conséquent, libère des protons plus facilement.Si l'hydrogène a une électronégativité égale ou supérieure que l'autre atome, cette molécule ne peut pas donner de protons et, par conséquent, ne peut pas agir comme un acide. Les oxyacides sont des acides auxquels un OH est attaché à un troisième atome plus électronégatif que l'hydrogène. La force d'un oxyacide dépend sur l'électronégativité et le nombre d'oxygène attaché au troisième atome.Plus l'électronégativité de l'atome est élevée, plus il se polarise, affaiblissant le lien entre l'oxygène et l'hydrogène. Si l'atome central est lié à des atomes d'oxygène supplémentaires, il augmente encore la polarité de la liaison entre l'oxygène et l'hydrogène. Par exemple, l'acide perchlorique avec trois atomes d'oxygène supplémentaire est plus fort que l'acide chlorique avec deux atomes d'oxygène supplémentaire.L'acide chlorique, à son tour, est plus fort que l'acide chloreux, qui n'a qu'un seul oxygène supplémentaire et acide hypochloreux, sans atomes d'oxygène supplémentaires. Les acides carboxyliques sont des acides faibles qui contiennent un groupe carboxyle. Le deuxième atome d'oxygène rend la liaison oxygène-hydrogène plus polaire, et permet ainsi à la molécule de donner un proton.L'acide acétique et l'acide formique sont des exemples d'acides carboxyliques.

force des acides et structure moléculaire

Acides et bases binaires 
 En l'absence de tout effet de nivellement, la force des acides pour les composés binaires de l'hydrogène avec des non-métaux (A) augmente à mesure que la force de la liaison H-A diminue en descendant dans un groupe du tableau périodique. Pour le groupe 17, l'ordre de l'acidité croissante est HF < HCl < HBr < HI. De même, pour le groupe 16, l'ordre de la force croissante des acides est H2O < H2S < H2Se < H2Te. Le long d'une rangée du tableau périodique, la force acide des composés d'hydrogène binaires augmente avec l'augmentation de l'électronégativité de l'atome non métallique parce que la polarité de la liaison H-A augmente. Ainsi, l'ordre de l'acidité croissante (pour l'élimination d'un proton) à travers la deuxième rangée est CH4 < NH3 < H2O < HF ; à travers la troisième rangée, c'est SiH4 < PH3 < H2S < HCl. 
 Acides et bases ternaires 
 Les composés ternaires constitués d'hydrogène, d'oxygène et d'un troisième élément (“ E ”) peuvent être structurés comme illustré dans l'image ci-dessous. Dans ces composés, l'atome E central est lié à un ou plusieurs atomes O, et au moins un des atomes O est également lié à un atome H, ce qui correspond à la formule moléculaire générale OmE(OH)n. Ces composés peuvent être acides, basiques ou amphotères selon les propriétés de l'atome E central. Des exemples de tels composés incluent l'acide sulfurique, O2S(OH)2, l'acide sulfureux, OS(OH)2, l'acide nitrique, O2NOH, l'acide perchlorique, O3ClOH, l'hydroxyde d'aluminium, Al(OH)3, l'hydroxyde de calcium, Ca(OH)2, et l'hydroxyde de potassium, KOH. 
 <img alt="Eq1" src="/files/ftp_upload/11406/11406_Figure_1.svg" style="height: 90px;" />
 Si l'atome central, E, possède une faible électronégativité, son attraction envers les électrons est faible. L'atome central a peu tendance à former une forte liaison covalente avec l'atome d'oxygène, et la liaison a entre l'élément et l'oxygène est plus facilement rompue que la liaison b entre l'oxygène et l'hydrogène. Par conséquent, la liaison a est ionique, les ions d'hydroxyde sont libérés dans la solution et le matériau se comporte comme une base, c'est le cas avec Ca(OH)2 et KOH. Une électronégativité plus faible est caractéristique des éléments plus métalliques ; par conséquent, les éléments métalliques forment des hydroxydes ioniques qui sont, par définition, des composés basiques. 
 Si, par contre, l'atome E a une électronégativité relativement élevée, il attire fortement les électrons qu'il partage avec l'atome d'oxygène, formant une liaison covalente relativement forte. La liaison oxygène-hydrogène, la liaison b, est ainsi affaiblie parce que les électrons sont déplacés vers E. La liaison b est polarisée et libère facilement des ions hydrogène dans la solution, de sorte que le matériau se comporte comme un acide. Les électronégativités élevées sont caractéristiques des éléments non métalliques. Ainsi, les éléments non métalliques forment des composés covalents contenant des groupes −OH acides appelés oxyacides. 
 L'augmentation du nombre d'oxydation de l'atome central E augmente également l'acidité d'un oxyacide car cela augmente l'attraction de E envers les électrons qu'il partage avec l'oxygène et affaiblit ainsi la liaison O-H. L'acide sulfurique, H2SO4 ou O2S(OH)2 (avec un nombre d'oxydation du soufre de +6), est plus acide que l'acide sulfureux, H2SO3 ou OS(OH)2 (avec un nombre d'oxydation du soufre de +4). De même, l'acide nitrique, HNO3 ou O2NOH (nombre d'oxydation de N = +5), est plus acide que l'acide nitreux, HNO2 ou ONOH (nombre d'oxydation de N = +3). Dans chacune de ces paires, le nombre d'oxydation de l'atome central est plus grand pour l'acide plus fort. 
 Acides carboxyliques 
 Les acides carboxyliques contiennent un groupe carboxyle. Les acides carboxyliques sont des acides faibles, ce qui signifie qu'ils ne sont pas ionisés à 100 % dans l'eau. 
 L'acide carboxylique agit comme un acide faible parce que, comme dans le cas des oxyacides, le deuxième oxygène fixé à l'atome de carbone augmente la polarité de la liaison O-H et la rend plus faible. De plus, après la perte du proton, le groupe carboxyle est converti en ion carboxylate, qui présente une résonance. Les différentes structures de résonance stabilisent l'ion carboxylate car sa charge négative est délocalisée sur plusieurs atomes. 
 Ce texte est adapté de<a href="https://openstax.org/books/chemistry-2e/pages/14-3-relative-strengths-of-acids-and-bases"> Openstax, Chimie 2e, Section 14.3 : Forces relatives des acides et bases</a> et<a href="https://openstax.org/books/chemistry-2e/pages/20-3-aldehydes-ketones-carboxylic-acids-and-esters"> Openstax, Chimie 2e, Section 20.3 Aldéhydes, cétones, acides carboxyliques, et esters</a>.

Force des acides et structure moléculaire

Acides et bases binaires
En l’absence de tout effet de nivellement, la résistance à l’acide des composés binaires de l’hydrogène avec des non-métaux (A) augmente à mesure que le H-A ...

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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

Water can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base. When hydrogen fluoride dissolves in water, it donates protons to water. This is acid ionization. The products are hydronium ions and fluoride ions. Water can also accept protons from another substance. When pyridine dissolves in water, it accepts a proton from water molecules. This is base ionization. The products are hydroxide ions and pyridinium ions. These reactions show that water can play either role. A substance that...

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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...

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Strong Acids, Strong Bases, and pH

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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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pH of Acid Mixtures

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...

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Salt Hydrolysis and pH in Solutions

Salt hydrolysis and pH in solutions depend on the ions released when a salt dissolves. Some ions react with water, while others do not. The balance between the cation and anion helps show whether the solution becomes acidic, basic, or neutral. Salts with pH-neutral anions and cations that produce hydronium ions give acidic solutions with pH below 7. Ammonium nitrate is one example. In NH4NO3, NO3− does not react with water, but NH4+ forms hydronium ions, so the solution is acidic. Other salts...

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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Polyprotic acids lose hydrogen ions in steps. The number of ionizable protons in a molecule determines whether an acid is monoprotic, diprotic, or triprotic. This video also shows that polyprotic bases can accept more than one proton. Monoprotic acids have one ionizable hydrogen atom per molecule. Examples include HCl, HNO3, and HCN. Their reactions with water form hydronium ions and the matching anion. Acetic acid, CH3CO2H, is also monoprotic because only the hydrogen in the carboxyl group...

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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Electron-Pair Reactions in Acid-Base Chemistry

Electron-pair reactions explain Lewis acid-base chemistry. In 1923, G. N. Lewis described acids and bases by how they accept or donate a pair of electrons. When that electron pair forms a coordinate covalent bond, one atom supplies both bonding electrons. This bond is also called a dative bond. A simple example is the reaction of water with a hydrogen ion to make a hydronium ion. Ammonia can also combine with a hydrogen ion to form an ammonium ion. In both cases, the reaction creates a Lewis...

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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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