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

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Chemistry

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Weak Acid Solutions

Um ácido fraco, como o ácido cianídrico, é um ácido Brønsted pois doa um próton à molécula de água e produz o íon hidrônico. Um ácido fraco dissocia-se parcialmente na água de acordo com a sua constante de dissociação ácida, Ka, que é 4, 9 10-¹⁰ para o ácido cianídrico. Para o ácido cianídrico, o Ka é igual à concentração hidrônica vezes a concentração de íons de cianeto divididos pela concentração do ácido cianídrico.A constante de dissociação ácida, Ka, pode ser utilizada para determinar a concentração de íons hidrônicos numa solução ácida fraca e consequentemente, o pH da solução. A concentração de íons hidrônicos e o pH de uma solução mols de 0, 15 de ácido cianídrico pode ser calculada utilizando a sua expressão de equilíbrio e uma tabela ICE. As concentrações de ácido cianídrico, hidrónio, e cianeto inicialmente e em equilíbrio podem ser expressas numa tabela que mostra as concentrações Inicial, de Alteração, e de Equilíbrio de cada uma das moléculas.Para alcançar o equilíbrio, a concentração inicial dos reagentes diminui como a concentração inicial dos produtos aumenta de acordo com as suas proporções molses. Esta alteração na concentração dos reagentes e produtos é assinalada por x. A substituição das concentrações de equilíbrio na expressão para os rendimentos Ka x vezes x dividido por 0, 15 menos x.Em muitos ácidos fracos, x, a quantidade de dissociação, é provável que seja muito pequena em comparação com a concentração inicial de 0, 15 mols. 0, 15 menos x pode ser aproximadamente 0, 15. Quando a equação é resolvida, x é igual a 8, 6 10-⁶ mols.A aproximação, 0.15 menos x igual a 0, 15, só é válido se x for menos do que 5%de 0, 15 mols. Aqui, x é 0, 0057%de 0, 15 mols e daí esta aproximação é válida. Por conseguinte, a concentração de hidrónio é de 8, 6 10-⁶ mols.Para determinar o pH, tome o registo negativo da concentração de íons hidrônicos. A resolução desta questão mostra o pH do ácido cianídrico 0, 15 M A solução é 5.07. O pH de uma solução pode ser utilizado para determinar o Ka de um ácido fraco.Por exemplo, o ácido acético dissocia-se parcialmente em íons hidrônicos e íons de acetato quando dissolvido em água. O Ka para o ácido acético pode ser expresso como a concentração de íons hidrônicos vezes a concentração de íons de acetato dividida pela concentração de ácido acético. Se o pH de uma solução de ácido acético 0, 20 mols é de 2, 72, a sua concentração de hidrónio pode ser calculada, sendo 1, 9 10-³ mols.A tabela ICE pode ser construída a partir das concentrações inicial e de equilíbrio do ácido acético, íons hidrônicos e íons de acetato. Utilizando números significativos, 0, 20 menos 1, 9 10-³ é essencialmente igual a 0, 20. Ao substituir os valores de Equilíbrio na expressão Ka, Ka equivale a 1, 8 10-⁵.

soluções de Ácidos fracos

Poucos compostos agem como ácidos fortes. Um número muito maior de compostos comporta-se como ácidos fracos e reage apenas parcialmente com a água, deixando uma grande maioria de moléculas dissolvidas na sua forma original e gerando uma quantidade relativamente pequena de iões hidrónio. Os ácidos fracos são comumente encontrados na natureza, sendo as substâncias parcialmente responsáveis pelo sabor acídico de frutas cítricas, a sensação de ardor de picadas de insetos, e os cheiros desagradáveis associados ao odor corporal. Um exemplo familiar de um ácido fraco é o ácido acético, o ingrediente principal no vinagre:
<img alt="Eq1" src="/files/ftp_upload/11400/11400_Equation_1.png" style="height: 50px;" />
O uso de uma seta dupla na equação acima denota o aspecto de reação parcial deste processo. Quando dissolvido em água em condições típicas, apenas cerca de 1% das moléculas de ácido acético estão presentes na forma ionizada, CH3COO−.
Cálculo das Concentrações de Iões Hidrónio e do pH de uma Solução de Ácido Fraco
O ácido fórmico, HCO2H, é um irritante que causa a reação do corpo a algumas picadas e mordidas de formigas. Qual é a concentração de iões hidrónio e o pH de uma solução de ácido fórmico a 0,534-M?
<img alt="Eq2" src="/files/ftp_upload/11400/11400_Equation_2.png" style="height: 50px;" />
A tabela ICE para este sistema é
<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;">Concentração Inicial (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;">Alteração (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;">Concentração em Equilíbrio (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>
Substituindo os termos de concentração de equilíbrio na expressão de Ka obtemos
<img alt="Eq3" src="/files/ftp_upload/11400/11400_Equation_3.png" style="height: 90px;" />
A concentração inicial relativamente grande e a constante de equilíbrio pequena permitem a suposição de simplificação de que x será muito menor que 0,534, e assim a equação torna-se
<img alt="Eq4" src="/files/ftp_upload/11400/11400_Equation_4.png" style="height: 90px;" />
Resolvendo a equação para x obtemos
<img alt="Eq5" src="/files/ftp_upload/11400/11400_Equation_5.png" style="height: 90px;" />
Para verificar a suposição de que x é pequeno em comparação com 0,534, a sua magnitude relativa pode ser estimada:
<img alt="Eq6" src="/files/ftp_upload/11400/11400_Equation_6.png" style="height: 90px;" />
Uma vez que x é inferior a 5% da concentração inicial, o pressuposto é válido. Conforme definido na tabela ICE, x é igual à concentração de equilíbrio de iões hidrónio:
<img alt="Eq7" src="/files/ftp_upload/11400/11400_Equation_7.png" style="height: 50px;" />
Por fim, o pH é calculado para ser
<img alt="Eq8" src="/files/ftp_upload/11400/11400_Equation_8.png" style="height: 50px;" />
Determinação de Ka a partir do pH
O pH de uma solução de ácido nitroso, HNO2, a 0,0516 M é de 2,34. Qual é a sua Ka?
<img alt="Eq9" src="/files/ftp_upload/11400/11400_Equation_9.png" style="height: 50px;" />
A concentração de ácido nitroso fornecida é uma concentração formal, que não tem em conta quaisquer equilíbrios químicos que possam ser estabelecidos em solução. Essas concentrações são tratadas como valores “iniciais” para os cálculos de equilíbrio utilizando a abordagem da tabela ICE. Observe que o valor inicial de iões hidrónio está listado como aproximadamente zero, pois uma pequena concentração de H3O+ está presente (1 × 10−7 M) devido à autoionização da água. Em muitos casos, esta concentração é muito menor do que aquela gerada pela ionização do ácido (ou base) em questão e pode ser negligenciada.
O pH fornecido é uma medida logarítmica da concentração de iões hidrónio resultante da ionização ácida do ácido nitroso, representando assim um valor de “equilíbrio” para a tabela ICE:
<img alt="Eq1" src="/files/ftp_upload/11400/11400_Equation_10.png" style="height: 50px;" />
A tabela ICE para este sistema é então
<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;">Concentração Inicial (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;">Alteração (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;">Concentração em Equilíbrio (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>
Por fim, calcula-se o valor da constante de equilíbrio utilizando os dados da tabela:
<img alt="Eq11" src="/files/ftp_upload/11400/11400_Equation_11.png" style="height: 90px;" />
Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/4-2-classifying-chemical-reactions">Openstax, Chemistry 2e, Section 4.2: Classifying Chemical Reactions</a> e <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>

Soluções de Ácidos Fracos

Poucos compostos atuam como ácidos fortes. Um número muito maior de compostos se comporta como ácidos fracos e reage apenas parcialmente com a água, deixando uma grande maioria de ...

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Acids and Bases

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KEY TERMS AND CONCEPTS

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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 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...

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pH and pOH in Water

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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

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...

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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...

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

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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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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...

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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...

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

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JoVE Core - Química dá vida a conceitos gerais de química usando vídeos de animação fáceis de entender. Além disso, nos vídeos Cientistas em ação, demonstramos experimentos de pesquisa clássicos e originais realizados atualmente por cientistas em laboratórios em todo o mundo.

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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 ...

comparing-intermolecular-forces-melting-point-boiling-point

Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, ...

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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