Name: how bonding affects acid strength
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 5 s
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North Carolina K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.6.3

Acid Strength and Molecular Structure

Salzsäure ist eine starke Säure, während Flusssäure eine schwache Säure ist. Aber was bestimmt ihre Stärke?
Die Stärke von binären Säuren mit nur zwei Elementen wird durch die Bindungsenergie und die Bindungspolarität bestimmt.
Eine Säure mit einer höheren Bindungsenergie hat eine stärkere Bindung und ist daher eine schwächere Säure. Vergleicht man Säuren in einer Gruppe, so ist die Bindung in einer schwachen Säure, wie Flusssäure, schwerer zu brechen, so dass es weniger wahrscheinlich ist, dass die Säure Protonen abgibt.
Im Gegensatz dazu hat eine Säure mit niedrigerer Bindungsenergie eine schwächere Bindung und ist daher eine stärkere Säure. Zum Beispiel bricht die Bindung in einer starken Säure wie Salzsäure leichter und gibt Protonen leichter ab als Flusssäure.
Eine Bindung, wie die in Salzsäure, ist polar, wenn ein Atom elektronegativer ist als das andere Atom.
Eine Verbindung kann als Säure wirken, wenn das an Wasserstoff gebundene Atom eine höhere Elektronegativität als Wasserstoff aufweist. Das Atom hat eine teilweise negative Ladung, so dass der Wasserstoff eine teilweise positive Ladung hat, so dass er als Proton freigesetzt werden kann.
Eine Säure mit höherer Bindungspolarität hat eine schwächere Bindung und ist daher eine stärkere Säure.
Vergleicht man Verbindungen über einen Zeitraum, so ist Salzsäure stärker als Schwefelwasserstoff, da Chlor elektronegativer ist als Schwefel und daher leichter Protonen freisetzt.
Wenn Wasserstoff die gleiche oder eine größere Elektronegativität als das andere Atom hat, kann dieses Molekül keine Protonen abgeben und daher nicht als Säure wirken.
Oxysäuren sind Säuren, bei denen ein OH an ein drittes Atom gebunden ist, das elektronegativer ist als Wasserstoff. Die Stärke einer Oxysäure hängt von der Elektronegativität und der Anzahl der Sauerstoffe ab, die an das dritte Atom gebunden sind.
Je höher die Elektronegativität des Atoms ist, desto mehr polarisiert es und schwächt die Bindung zwischen Sauerstoff und Wasserstoff.
Wird das Zentralatom an weitere Sauerstoffatome gebunden, erhöht sich die Polarität der Bindung zwischen dem Sauerstoff und dem Wasserstoff weiter.
Zum Beispiel ist Perchlorsäure mit drei zusätzlichen Sauerstoffatomen stärker als Chlorsäure mit zwei zusätzlichen Sauerstoffatomen. Chlorsäure wiederum ist stärker als Chlorsäure, die nur einen zusätzlichen Sauerstoff und eine hypochlorige Säure ohne zusätzliche Sauerstoffatome enthält.
Carbonsäuren sind schwache Säuren, die eine Carboxylgruppe enthalten. Das zweite Sauerstoffatom macht die Sauerstoff-Wasserstoff-Bindung polarer und ermöglicht es dem Molekül, ein Proton abzugeben. Essigsäure und Ameisensäure sind Beispiele für Carbonsäuren.

säurestärke und molekülstruktur

Binäre Säuren und Basen Ohne eine nivellierende Wirkung nimmt die Säurestärke binärer Verbindungen von Wasserstoff mit Nichtmetallen (A) zu, wenn die Stärke der H-A-Bindung in einer Gruppe im Periodensystem abnimmt. Für Gruppe 17 ist die Reihenfolge des zunehmenden Säuregehalts HF < HCl < HBr < HI. Ebenso ist für Gruppe 16 die Reihenfolge der zunehmenden Säurestärke H2O < H2S < H2Se < H2Te. In einer Zeile des Periodensystems nimmt die Säurestärke binärer Wasserstoffverbindungen mit zunehmender Elektronegativität des Nichtmetallatoms zu, da die Polarität der H-A-Bindung zunimmt. Somit ist die Reihenfolge der zunehmenden Säure (zur Entfernung eines Protons) in der zweiten Reihe CH4 < NH3 < H2O < HF; In der dritten Reihe gilt: SiH4 < PH3 < H2S < HCl. Ternäre Säuren und Basen Ternäre Verbindungen, die aus Wasserstoff, Sauerstoff und einem dritten Element („E“) bestehen, können wie im Bild unten dargestellt strukturiert sein. In diesen Verbindungen ist das zentrale E-Atom an ein oder mehrere O-Atome gebunden, und mindestens eines der O-Atome ist zusätzlich an ein H-Atom gebunden, entsprechend der allgemeinen Summenformel OmE(OH)n. Abhängig von den Eigenschaften des zentralen E-Atoms können diese Verbindungen sauer, basisch oder amphoter sein. Beispiele für solche Verbindungen sind Schwefelsäure, O2S(OH)2, schweflige Säure, OS(OH)2, Salpetersäure, O2NOH, Perchlorsäure, O3ClOH, Aluminiumhydroxid, Al(OH)3, Calciumhydroxid, Ca(OH)2 und Kaliumhydroxid, KOH. <img alt="Eq1" src="/files/ftp_upload/11406/11406_Figure_1.svg" style="height: 90px;"> Wenn das Zentralatom E eine niedrige Elektronegativität aufweist, ist seine Anziehungskraft für Elektronen gering. Das Zentralatom neigt kaum dazu, eine starke kovalente Bindung mit dem Sauerstoffatom einzugehen, und die Bindung a zwischen dem Element und Sauerstoff wird leichter aufgebrochen als die Bindung b zwischen Sauerstoff und Wasserstoff. Daher ist Bindung a ionisch, Hydroxidionen werden an die Lösung abgegeben und das Material verhält sich wie eine Base – dies ist bei Ca(OH)2 und KOH der Fall. Charakteristisch für die eher metallischen Elemente ist eine geringere Elektronegativität; Daher bilden die metallischen Elemente ionische Hydroxide, die per Definition basische Verbindungen sind. Wenn das Atom E hingegen eine relativ hohe Elektronegativität aufweist, zieht es die Elektronen, die es mit dem Sauerstoffatom teilt, stark an und bildet so eine relativ starke kovalente Bindung. Die Sauerstoff-Wasserstoff-Bindung, Bindung b, wird dadurch geschwächt, da Elektronen in Richtung E verschoben werden. Bindung b ist polar und gibt leicht Wasserstoffionen an die Lösung ab, sodass sich das Material wie eine Säure verhält. Hohe Elektronegativitäten sind charakteristisch für die eher nichtmetallischen Elemente. Daher bilden nichtmetallische Elemente kovalente Verbindungen mit sauren −OH-Gruppen, die als Oxysäuren bezeichnet werden. Eine Erhöhung der Oxidationszahl des Zentralatoms E erhöht auch den Säuregehalt einer Oxysäure, da dies die Anziehungskraft von E für die Elektronen, die es mit Sauerstoff teilt, erhöht und dadurch die O-H-Bindung schwächt. Schwefelsäure, H2SO4, oder O2S(OH)2 (mit einer Schwefeloxidationszahl von +6), ist saurer als schweflige Säure, H2SO3, oder OS(OH)2 (mit einer Schwefeloxidationszahl von +4). Ebenso ist Salpetersäure, HNO3, oder O2NOH (N-Oxidationszahl = +5), saurer als salpetrige Säure, HNO2, oder ONOH (N-Oxidationszahl = +3). In jedem dieser Paare ist die Oxidationszahl des Zentralatoms für die stärkere Säure größer. Carbonsäuren Carbonsäuren enthalten eine Carboxylgruppe. Carbonsäuren sind schwache Säuren, das heißt, sie sind in Wasser nicht zu 100 % ionisiert. Carbonsäure wirkt als schwache Säure, da wie bei Oxysäuren der zweite, an das Kohlenstoffatom gebundene Sauerstoff die Polarität der O-H-Bindung erhöht und sie schwächer macht. Darüber hinaus wird die Carboxylgruppe nach dem Verlust des Protons in das Carboxylat-Ion umgewandelt, das Resonanz zeigt. Die unterschiedlichen Resonanzstrukturen stabilisieren das Carboxylat-Ion, da seine negative Ladung über mehrere Atome verteilt ist.

Säurestärke und Molekülstruktur

Binäre Säuren und Basen
In Ermangelung eines Nivellierungseffekts nimmt die Säurestärke von binären Verbindungen des Wasserstoffs mit Nichtmetallen (A) zu, wenn die H-A ...

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Acids and Bases

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KEY TERMS AND CONCEPTS

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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...

ph scale

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

relative-strengths-of-conjugate-acid-base-pairs

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

Strong acids and strong bases change the pH of a solution by dissociating completely in water. A strong acid produces hydronium ions, which are H3O+ ions, and a strong base produces hydroxide ions, OH-. Because the dissociation is complete, the ion concentration can be found from the starting concentration and the molar ratio. For example, 0.20 M hydrobromic acid dissociates fully in water. It produces 0.20 M hydronium ions and 0.20 M bromide ions. A 0.015 M KOH solution also dissociates...

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...

weak base solutions

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...

mixtures of acids

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

Salt Hydrolysis and pH of Solutions

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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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

acidbase-strengths-and-dissociation-constants

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

Stepwise Ionization of Polyprotic Acids

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...

lewis acids and bases

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...

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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New York State Science Learning Standards

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How Bonding Affects Acid Strength

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