Name: how ka and kb show acid-base strength
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 29 s
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Relative Strengths of Conjugate Acid-Base Pairs

La relation entre les constantes de dissociation d'une paire acide-base conjuguée peut être exprimée quantitativement. Pour un acide faible, HA, sa constante de dissociation acide, Ka, est exprimée en concentration d'ions hydronium fois la concentration d'ions A divisée par la concentration de HA.La constante de dissociation de base, Kb, pour sa base conjuguée, l'ion A, est exprimée en concentration d'HA fois la concentration en ions hydroxyde divisée par la concentration de l'ion A.Si les expressions pour Ka et Kb sont multipliées, l'équation résultante est l'expression d'équilibre pour Kw.Cette équation montre que le Ka pour un acide faible et le Kb pour sa base conjuguée sont inversement proportionnelles. Un acide plus fort avec un Ka plus élevé a toujours une base conjuguée proportionnellement plus faible avec un Kb inférieur et vice versa.Comme la valeur de Kw est constante, cette équation peut être utilisée pour calculer Ka ou Kb d'une paire acide-base conjuguée si l'un d'entre eux est connu. Par exemple, si le Kb d'une base est de 1 10⁻⁶, alors le Ka de son acide conjugué peut être calculé comme étant de 1 10⁻⁸. La relation entre Ka, Kb et Kw peut également être exprimée en termes de leurs logarithmes négatifs, pKa, pKb et pKw.Pour l'équation, Ka fois Kb est égal à Kw, lorsque les logarithmes négatifs des deux côtés sont pris, l'équation résultante est pKa plus pKb qui est égal à pKw, et qui est de 14 à 25 degrés Celsius. le pKa et le pKb peuvent également être utilisés pour comparer les forces des acides et des bases faibles. Plus la valeur du pKa est basse, plus l'acide est fort.De la même manière, plus la valeur du pKb est basse, plus la base est forte. Par exemple, un acide avec un pKa de 2, 1 est plus fort qu'un acide avec un pKa de 4, 6.

forces relatives des paires acide-base conjuguées

La chimie des acide-base de Brønsted-Lowry est le transfert de protons ; ainsi, la logique suggère une relation entre les forces relatives des couples acide-base conjugués. La force d'un acide ou d'une base est quantifiée dans sa constante d'ionisation, Ka ou Kb, qui représente l'étendue de la réaction d'ionisation de l'acide ou de la base. Pour le couple conjugué acide-base HA/A−, les équations des équilibres d'ionisation et les expressions des constantes d'ionisation sont 
 <img alt="Eq1" src="/files/ftp_upload/11398/11398_Equation_1.png" style="height: 90px;" />
 <img alt="Eq2" src="/files/ftp_upload/11398/11398_Equation_2v2.png" style="height: 70px;" />
 L'addition de ces deux équations chimiques conduit à l'équation de l'auto-ionisation de l'eau : 
 <img alt="Eq3" src="/files/ftp_upload/11398/11398_Equation_3.png" style="height: 50px;" />
 Comme nous l'avons déjà mentionné, la constante d'équilibre pour la somme des réactions est égale au produit mathématique des constantes d'équilibre pour les réactions additionnées, et ainsi de suite 
 <img alt="Eq4" src="/files/ftp_upload/11398/11398_Equation_4.png" style="height: 90px;" />
 Cette équation indique la relation entre les constantes d'ionisation pour tout couple acide-base conjugué, à savoir, leur produit mathématique est égal au produit ionique de l'eau, KE. En réorganisant cette équation, une relation réciproque entre les forces d'un couple acide-base conjugué devient évidente : 
 <img alt="Eq5" src="/files/ftp_upload/11398/11398_Equation_5.png" style="height: 90px;" />
 La relation inversement proportionnelle entre Ka et Kb signifie que plus l'acide ou la base est fort, plus son partenaire conjugué est faible. 
 En prenant le log négatif des deux côtés de l'équation, Ka × Kb = KE donne 
 <img alt="Eq6" src="/files/ftp_upload/11398/11398_Equation_6.png" style="height: 50px;" />
 alors 
 <img alt="Eq7" src="/files/ftp_upload/11398/11398_Equation_7.png" style="height: 50px;" />
 Comme le pKE est de 14 à 25 °C, cette équation peut également être écrite comme 
 <img alt="Eq1" src="/files/ftp_upload/11398/11398_Equation_8.png" style="height: 50px;" />
 Le pKa et le pKb représentent également la force des acides et des bases, respectivement. Comme le pH et le pOH, plus la valeur de pKa ou pKb est élevée, plus l'acide ou la base est faible, respectivement.  
 <table style="margin-left: auto; margin-right: auto;">
 <tbody>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Base </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide perchlorique (HClO4)* </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion perchlorate (ClO4−)** </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide sulfurique (H2SO4)* </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydrogénosulfate (HSO4−)** </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Iodure d'hydrogène (HI)* </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion iodure (I−)** </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Bromure d'hydrogène (HBr)* </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion bromure (Br−)** </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Chlorure d'hydrogène (HCl)* </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion chlorure (Cl−)** </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide nitrique (HNO3)* </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion nitrate (NO3−)** </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydronium (H3O+) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Eau (H2O) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion sulfate d'hydrogène (HSO4−) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion sulfate (SO42−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide phosphorique (H3PO4) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion dihydrogénophosphate (H2PO4−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Fluorure d'hydrogène (HF) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion fluorure (F−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide nitreux (HNO2) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion nitrite (NO2−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide acétique (CH3CO2H) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion acétate (CH3CO2−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Acide carbonique (H2CO3) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydrogénocarbonate (HCO3−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Sulfure d'hydrogène (H2S) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydrogénosulfure (HS−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion ammonium (NH4+) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ammoniac (NH3) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Cyanure d'hydrogène (HCN) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion cyanure (CN−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydrogénocarbonate (HCO3−) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion carbonate (CO32−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Eau (H2O) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydroxyde (OH−) </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydrogénosulfure (HS−)†</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion sulfure (S2−)‡</td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Éthanol (C2H5OH)†</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion éthoxyde (C2H5O−)‡</td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Ammoniac (NH3)∔</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion amidure (NH2−)‡</td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Hydrogène (H2)†</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion hydrure (H−)‡</td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">Méthane (CH4)†</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Ion méthide (CH3−)‡</td>
 </tr>
 <tr>
 <td colspan="2" style="color: black; border: 1px solid black; ">*subissent une ionisation totale de l'acide dans l'eau </td>
 </tr>
 <tr>
 <td colspan="2" style="color: black; border: 1px solid black; ">†ne subit pas d'ionisation de l'acide dans l'eau </td>
 </tr>
 <tr>
 <td colspan="2" style="color: black; border: 1px solid black; ">**ne subit pas d'ionisation de la base dans l'eau </td>
 </tr>
 <tr>
 <td colspan="2" style="color: black; border: 1px solid black; ">‡ subissent une ionisation totale de la base dans l'eau </td>
 </tr>
 </tbody>
 </table>
 La liste des couples acide–base conjugués est organisée de façon à montrer la force relative de chaque espèce par rapport à l'eau. Dans la colonne des acides, les espèces énumérées au-dessous de l'eau sont des acides plus faibles que l'eau. Ces espèces ne subissent pas d'ionisation de l'acide dans l'eau ; ce ne sont pas des acides de Br&#248nsted-Lowry. Toutes les espèces énumérées au-dessus de l'eau sont des acides plus forts, transférant les protons à l'eau dans une certaine mesure lorsqu'elles sont dissoutes dans une solution aqueuse pour générer des ions hydronium. Les espèces au-dessus de l'eau mais au-dessous de l'ion hydronium sont des acides faibles, soumis à une ionisation de l'acide partielle, alors que celles au-dessus de l'ion hydronium sont des acides forts qui sont complètement ionisés en solution aqueuse. 
 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:</a><a href="https://openstax.org/books/chemistry-2e/pages/14-3-relative-strengths-of-acids-and-bases"> Forces relatives des acides et des bases</a>.

Forces relatives des paires acide-base conjuguées

La chimie acido-basique de Brønsted-Lowry est le transfert de protons ; Ainsi, la logique suggère une relation entre les forces relatives de l’acide conjugué ...

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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 equilibrium constant for the summed reaction is the product of the two individual constants, so Ka × Kb = Kw for any conjugate acid-base pair.

This relationship shows a reciprocal pattern. If one member of the pair is stronger, its conjugate partner is weaker. The stronger the acid, the weaker its conjugate base, and the stronger the base, the weaker its conjugate acid.

Taking the negative log of Ka × Kb = Kw gives the pKa and pKb form of the same idea. At 25 °C, pKw is 14, so the relationship can also be written as pKa + pKb = 14. Like pH and pOH, a larger pKa or pKb means a weaker acid or base.

The conjugate acid-base list is arranged to compare each species with water. Species below water in the acid column are weaker acids than water, so they do not undergo acid ionization in water and are not Brønsted-Lowry acids. Species above water transfer protons to water to some extent and form hydronium ions.

Among those acids, the ones above water but below hydronium ion are weak acids, so they ionize only partially in aqueous solution. The acids above hydronium ion are strong acids, and they are completely ionized in water. The list includes examples such as perchloric acid, sulfuric acid, hydrogen chloride, nitric acid, acetic acid, ammonia, ethanol, hydride, and methane with their conjugate partners.

How Ka and Kb Show Acid-Base 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 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 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

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...

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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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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...

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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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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...

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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...

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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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See how Ka, Kb, pKa, and pKb reveal the strength of conjugate acid-base pairs, including why stronger acids have weaker conjugate bases.

How Ka and Kb Show Acid-Base Strength | JoVE K12

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Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

PS.Chm.2.2

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

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

trends-in-lattice-energy-ion-size-and-charge

Trends in Lattice Energy: Ion Size and Charge

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

Analyze the structure of atoms and isotopes.

How Ka and Kb Show Acid-Base Strength

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