Name: weak acid equilibrium in water
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 4 min 2 s
Channel: JoVE

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North Carolina K-12 Science Standard Course of Study

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Chemistry

Matter and its Interactions

PS.Chm.6.3

Weak Acid Solutions

Eine schwache Säure, wie Blausäure, ist eine Brønsted-Säure, da sie ein Proton an das Wassermolekül abgibt und das Hydronium-Ion produziert. Eine schwache Säure dissoziiert teilweise in Wasser entsprechend ihrer Säuredissoziationskonstante Ka, die für Blausäure 4,9 × 10−10 beträgt.
Bei Blausäure ist K a gleich der Konzentration von Hydronium multipliziert mit der Konzentration von Cyanidionen dividiert durch die Konzentration von Blausäure.
Die Säuredissoziationskonstante Ka kann verwendet werden, um die Hydroniumionenkonzentration in einer schwach sauren Lösung und damit den pH-Wert der Lösung zu bestimmen.
Die Konzentration von Hydroniumionen und der pH-Wert einer 0,15 molaren Blausäurelösung können anhand ihrer Gleichgewichtsexpression und einer ICE-Tabelle berechnet werden.
Die Konzentrationen von Blausäure, Hydronium und Cyanid zu Beginn und im Gleichgewicht können in einer Tabelle ausgedrückt werden, die die Anfangs-, Änderungs- und Gleichgewichtskonzentrationen jedes der Moleküle zeigt.
Um ein Gleichgewicht zu erreichen, nimmt die Anfangskonzentration der Reaktanten ab, wenn die Anfangskonzentration der Produkte entsprechend ihren molaren Verhältnissen zunimmt. Diese Änderung der Konzentration der Reaktanten und Produkte wird mit x bezeichnet.
Das Einsetzen von Gleichgewichtskonzentrationen im Ausdruck für Ka ergibt x mal x geteilt durch 0,15 minus x. 
In vielen schwachen Säuren ist x, die Dissoziationsmenge, im Vergleich zur Anfangskonzentration von 0,15 molar wahrscheinlich sehr gering. 0,15 minus x kann mit etwa 0,15 angenommen werden.
Wenn die Gleichung gelöst ist, ist x gleich 8,6 × 10−6 molaren.
Die Näherung 0,15 minus x gleich 0,15 ist nur gültig, wenn x kleiner als 5 % von 0,15 molar ist. Hier ist x 0,0057% von 0,15 molar und daher ist diese Näherung gültig.
Daher beträgt die Konzentration von Hydronium 8,6 × 10-6 molaren. Um den pH-Wert zu bestimmen, nehmen Sie das negative Protokoll der Hydroniumionenkonzentration. Die Lösung dieses Problems zeigt, dass der pH-Wert der 0,15 M Blausäurelösung 5,07 beträgt.
Der pH-Wert einer Lösung kann verwendet werden, um den Ka einer schwachen Säure zu bestimmen.
So dissoziiert Essigsäure teilweise in Hydronium-Ionen und Acetat-Ionen, wenn sie in Wasser gelöst wird. Das Ka für Essigsäure kann ausgedrückt werden als die Hydroniumionenkonzentration multipliziert mit der Acetationenkonzentration dividiert durch die Konzentration der Essigsäure.
Wenn der pH-Wert einer 0,20 molaren Essigsäurelösung 2,72 beträgt, kann ihre Hydroniumkonzentration berechnet werden, die 1,9 × 10−3 molar beträgt.
Die ICE-Tabelle kann aus den Anfangs- und Gleichgewichtskonzentrationen der Essigsäure-, Hydronium- und Acetationen erstellt werden. 
Unter Verwendung signifikanter Zahlen ist 0,20 minus 1,9 × 10−3 im Wesentlichen gleich 0,20. Durch Einsetzen der Gleichgewichtswerte in den Ausdruck Ka ergibt sich für Ka 1,8 × 10−5.

schwache säurelösungen

Nur wenige Verbindungen wirken als starke Säuren. Eine weitaus größere Anzahl von Verbindungen verhalten sich wie schwache Säuren und reagieren nur teilweise mit Wasser, wodurch ein Großteil der gelösten Moleküle in ihrer ursprünglichen Form zurückbleibt und eine relativ kleine Menge an Hydroniumionen erzeugt. Schwache Säuren kommen in der Natur häufig vor. Sie sind mitverantwortlich für den würzigen Geschmack von Zitrusfrüchten, das stechende Gefühl bei Insektenstichen und die unangenehmen Gerüche, die mit Körpergeruch einhergehen. Ein bekanntes Beispiel für eine schwache Säure ist Essigsäure, der Hauptbestandteil von Essig: <img alt="Eq1" src="/files/ftp_upload/11400/11400_Equation_1.png" style="height: 50px;"> Die Verwendung eines Doppelpfeils in der obigen Gleichung bezeichnet den Teilreaktionsaspekt dieses Prozesses. Wenn es unter typischen Bedingungen in Wasser gelöst wird, liegen nur etwa 1 % der Essigsäuremoleküle in der ionisierten Form CH3COO− vor. Berechnung der Hydroniumionenkonzentrationen und des pH-Werts einer schwach sauren Lösung Ameisensäure, HCO2H, ist ein Reizstoff, der die Reaktion des Körpers auf einige Ameisenbisse und -stiche hervorruft. Wie hoch ist die Hydroniumionenkonzentration und der pH-Wert einer 0,534 M Ameisensäurelösung? <img alt="Eq2" src="/files/ftp_upload/11400/11400_Equation_2.png" style="height: 50px;"> Die ICE-Tabelle für dieses System ist <table style="margin-left: auto; margin-right: auto;"> <tbody> <tr> <td style="color: black; border: 1px solid black; text-align: center;"></td> <td style="color: black; border: 1px solid black; text-align: center;"> HCO2H (aq) </td> <td style="color: black; border: 1px solid black; text-align: center;">H3O+ (aq)</td> <td style="color: black; border: 1px solid black; text-align: center;">HCO2− (aq)</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center;">Initial Concentration (M)</td> <td style="color: black; border: 1px solid black; text-align: center;">0.534</td> <td style="color: black; border: 1px solid black; text-align: center;">~0</td> <td style="color: black; border: 1px solid black; text-align: center;">0</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center;">Change (M)</td> <td style="color: black; border: 1px solid black; text-align: center;">−x</td> <td style="color: black; border: 1px solid black; text-align: center;">+x</td> <td style="color: black; border: 1px solid black; text-align: center;">+x</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center;">Equilibrium Concentration (M)</td> <td style="color: black; border: 1px solid black; text-align: center;">0.534 − x</td> <td style="color: black; border: 1px solid black; text-align: center;">x</td> <td style="color: black; border: 1px solid black; text-align: center;">x</td> </tr> </tbody> </table> Das Einsetzen der Gleichgewichtskonzentrationsterme in den Ka-Ausdruck ergibt <img alt="Eq3" src="/files/ftp_upload/11400/11400_Equation_3.png" style="height: 90px;"> Die relativ große Anfangskonzentration und die kleine Gleichgewichtskonstante erlauben die vereinfachende Annahme, dass x viel kleiner als 0,534 sein wird, und so ergibt sich die Gleichung <img alt="Eq4" src="/files/ftp_upload/11400/11400_Equation_4.png" style="height: 90px;"> Das Auflösen der Gleichung nach x ergibt <img alt="Eq5" src="/files/ftp_upload/11400/11400_Equation_5.png" style="height: 90px;"> Um die Annahme zu überprüfen, dass x im Vergleich zu 0,534 klein ist, kann seine relative Größe geschätzt werden: <img alt="Eq6" src="/files/ftp_upload/11400/11400_Equation_6.png" style="height: 90px;"> Da x weniger als 5 % der Anfangskonzentration beträgt, ist die Annahme gültig. Wie in der ICE-Tabelle definiert, ist x gleich der Gleichgewichtskonzentration von Hydroniumionen: <img alt="Eq7" src="/files/ftp_upload/11400/11400_Equation_7.png" style="height: 50px;"> Abschließend wird der pH-Wert berechnet <img alt="Eq8" src="/files/ftp_upload/11400/11400_Equation_8.png" style="height: 50px;"> Bestimmung von Ka aus pH Der pH-Wert einer 0,0516 M Lösung salpetriger Säure, HNO2, beträgt 2,34. Was ist sein Ka? <img alt="Eq9" src="/files/ftp_upload/11400/11400_Equation_9.png" style="height: 50px;"> Bei der angegebenen salpetrigen Säurekonzentration handelt es sich um eine formale Konzentration, die keine chemischen Gleichgewichte berücksichtigt, die sich in der Lösung einstellen können. Solche Konzentrationen werden als „Ausgangswerte“ für Gleichgewichtsberechnungen unter Verwendung des ICE-Tabellenansatzes behandelt. Beachten Sie, dass der Anfangswert des Hydroniumions mit ungefähr Null angegeben ist, da aufgrund der Autoionisierung von Wasser eine geringe Konzentration an H3O+ vorhanden ist (1 × 10−7 M). In vielen Fällen ist diese Konzentration viel geringer als die, die durch die Ionisierung der betreffenden Säure (oder Base) erzeugt wird, und kann vernachlässigt werden. Der angegebene pH-Wert ist ein logarithmisches Maß für die Hydroniumionenkonzentration, die sich aus der Säureionisierung der salpetrigen Säure ergibt, und stellt somit einen „Gleichgewichtswert“ für die ICE-Tabelle dar: <img alt="Eq1" src="/files/ftp_upload/11400/11400_Equation_10.png" style="height: 50px;"> Die ICE-Tabelle für dieses System lautet dann <table style="margin-left: auto; margin-right: auto;"> <tbody> <tr> <td style="color: black; border: 1px solid black; text-align: center;"></td> <td style="color: black; border: 1px solid black; text-align: center;"> HNO2 (aq)  </td> <td style="color: black; border: 1px solid black; text-align: center;">H3O+ (aq)</td> <td style="color: black; border: 1px solid black; text-align: center;">NO2− (aq)</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center;">Initial Concentration (M)</td> <td style="color: black; border: 1px solid black; text-align: center;">0.0516</td> <td style="color: black; border: 1px solid black; text-align: center;">~0</td> <td style="color: black; border: 1px solid black; text-align: center;">0</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center;">Change (M)</td> <td style="color: black; border: 1px solid black; text-align: center;">−0.0046</td> <td style="color: black; border: 1px solid black; text-align: center;">+0.0046</td> <td style="color: black; border: 1px solid black; text-align: center;">+0.0046</td> </tr> <tr> <td style="color: black; border: 1px solid black; text-align: center;">Equilibrium Concentration (M)</td> <td style="color: black; border: 1px solid black; text-align: center;">0.0470</td> <td style="color: black; border: 1px solid black; text-align: center;">0.0046</td> <td style="color: black; border: 1px solid black; text-align: center;">0.0046</td> </tr> </tbody> </table> Berechnen Sie abschließend den Wert der Gleichgewichtskonstante anhand der Daten in der Tabelle: <img alt="Eq11" src="/files/ftp_upload/11400/11400_Equation_11.png" style="height: 90px;"> Dieser Text ist adaptiert von <a href="https://openstax.org/books/chemistry-2e/pages/4-2-classifying-chemical-reactions">Openstax, Chemistry 2e, Section 4.2: Classifying Chemical Reactions</a> and <a href="https://openstax.org/books/chemistry-2e/pages/14-3-relative-strengths-of-acids-and-bases">Openstax, Chemistry 2e,14.3 Relative Strengths of Acids and Bases.</a>

Schwache Säurelösungen

Nur wenige Verbindungen wirken als starke Säuren. Eine weitaus größere Anzahl von Verbindungen verhält sich wie schwache Säuren und reagiert nur teilweise mit Wasser, so dass ein großer Teil der Verbindungen ...

Education

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KERNBEGRIFFE UND KONZEPTE

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 water is only partial. Under typical conditions, only about 1% of acetic acid molecules are present as the ionized form, CH3COO−.

Weak acid problems often use an ICE table to find hydronium ion concentration and pH. Formic acid, HCO2H, is one example because it can cause the irritation from some ant bites and stings. For a 0.534-M formic acid solution, the ICE table starts with formic acid, H3O+, and HCO2−, then tracks the change as x forms and the equilibrium amounts are written in terms of x. The Ka expression is set up from those equilibrium concentrations.

Because the initial formic acid concentration is much larger than the amount that reacts, x is assumed to be much less than 0.534. That lets the expression simplify and makes the calculation easier. After solving for x, the result is checked against the 5% rule. Since x is less than 5% of the initial concentration, the assumption is valid. The value of x equals the equilibrium hydronium ion concentration, and the pH is then calculated from that concentration.

The same ICE table method can also be used in reverse to find Ka from pH. For nitrous acid, HNO2, the given pH is 2.34 for a 0.0516-M solution. That concentration is a formal concentration, which means it does not include any equilibrium changes. In these setups, the initial hydronium ion is shown as approximately zero because water already contains a very small amount, about 1 × 10−7 M, from autoionization. In many cases, that amount is much smaller than the acid-produced hydronium ion and can be ignored.

The pH gives the equilibrium hydronium ion concentration for the nitrous acid solution. The ICE table then shows HNO2, H3O+, and NO2− changing to equilibrium values. Using those values in the equilibrium expression gives the Ka of the acid. This approach shows how pH and Ka connect to the behavior of weak acids in water.

Weak Acid Equilibrium in Water

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...

water: a bronsted-lowry acid and base

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 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

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

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

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...

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

Explore how weak acids partially ionize in water, use ICE tables to find hydronium ion concentration and pH, and calculate Ka from pH.

Weak Acid Equilibrium in Water | JoVE K12

chemistry

JoVE Core Chemie erweckt allgemeine Konzepte der Chemie durch kompakte und leicht verständliche animierte Videolektionen zum Leben. Außerdem zeigen wir in den Videos Wissenschaftler in Aktion bei klassischen und neuartigen Forschungsexperimenten, wie sie heutzutage in Laboren weltweit durchgeführt wurden.

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New York State Science Learning Standards

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HS-PS1-11.

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An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

bond-polarity-dipole-moment-and-percent-ionic-character

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

lewis-structures-of-molecular-compounds-and-polyatomic-ions

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Resonance

The Lewis structure of a nitrite anion (NO2&minus;) may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O ...

Formal Charges

In some cases, there are seemingly more than one valid Lewis structures for molecules and polyatomic ions. The concept of formal charges can be used to ...

Exceptions to the Octet Rule

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:


	...

Bond Energies and Bond Lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

PS.Chm.3.1

molecular-comparison-of-gases-liquids-and-solids

Molecular Comparison of Gases, Liquids, and Solids

Particles in a solid are tightly packed together (fixed shape) and often arranged in a regular pattern; in a liquid, they are close together with no ...

Intermolecular vs Intramolecular Forces

Intermolecular forces (IMF) are electrostatic attractions arising from charge-charge interactions between molecules. The strength of the intermolecular ...

Intermolecular Forces

Atoms and molecules interact through bonds (or forces): intramolecular and intermolecular. The forces are electrostatic as they arise from interactions ...

Construct an explanation to summarize the influences intermolecular forces have on the properties of chemical compounds.

PS.Chm.3.2

Analyze the structure of atoms and isotopes.

Weak Acid Equilibrium in Water

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