Name: electron-pair reactions in acid-base chemistry
Uploaded: 2020-09-24T15:06:51.000Z
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Lewis Acids and Bases

Модель Бренстеда-Лоури определяет кислоты и основания в терминах протонов, где кислоты являются донорами протонов, а основания акцепторами протонов. Напротив, модель Льюиса определяет кислоты и основания в терминах электронных пар, где кислоты Льюиса являются акцепторами электронных пар, а основания Льюиса являются донорами электронных пар. В кислоте Бренстеда, такой как уксусная кислота, водород также может действовать как кислота Льюиса, потому что у него есть пустая орбиталь, чтобы принимать электроны, отданные основанием, таким как вода, действующая как основание Льюиса.Преимущество модели Льюиса состоит в том, что она позволяет ученым классифицировать большее количество соединений как кислоты, в том числе те, которые не имеют ионизируемых протонов. Например, трифторид бора не может быть классифицирован как кислота по модели Бренстеда-Лоури, потому что он не содержит водорода. Однако трифторид бора обладает неполным октетом с пустой орбиталью, которая может принимать электронную пару от основания Льюиса, такого как аммиак, и, следовательно, может действовать как кислота Льюиса.Результирующий продукт, образованный такими кислотно-основными реакциями Льюиса, называется кислотно-основным аддуктом Льюиса. Некоторые молекулы, такие как углекислый газ, могут перегруппировывать свои электроны и действовать как кислота Льюиса. Например, в реакции между водой и диоксидом углерода электронная пара перемещается от пи-связи углерод-кислород к концевому кислороду диоксида углерода.Получающаяся в результате пустая орбиталь на атоме углерода позволяет ему принимать электронную пару от молекулы воды и действовать как кислота Льюиса. Поскольку молекула воды отдает электронную пару, она действует как основание Льюиса. При дальнейшей перегруппировке протон переносится от кислорода воды к концевому кислороду диоксида углерода, что приводит к образованию аддукта угольной кислоты.Катионы небольших металлов, такие как Al могут повторно захватывать электронные пары и действовать как кислоты Льюиса. Например, Al принимает неподеленные пары электронов из воды и образует ионы гексаакваалюминия. Здесь молекулы воды отдают электронные пары и действуют как основание Льюиса.

Кислоты и основания Льюиса

В 1923 г. г. Н. Льюис предложил обобщенное определение поведения кислотно-основная, в котором кислоты и основания идентифицируются по их способности принимать или жертвовать пару электронов и формировать координатную ковалентную связь. 
 Связующее связывание координат (или денативная связь) происходит, когда один из атомов в связи обеспечивает оба связующего электрона. Например, связка ковалента координат возникает, когда молекула воды соединяется с ионом водорода, образуя ион гидрония. Связка ковалента координат также возникает, когда молекула аммиака соединяется с ионом водорода, образуя ион аммония. Здесь показаны оба этих уравнения. 
 <img alt="Image1" src="/files/ftp_upload/11407/11407_Image_1.svg" style="height: 175px;" />
 Реакции, связанные с формированием координатных ковалентных связей, классифицируются как химия Lewis кислотно-основная. Вид, жертвующий электронной парой, которая составляет связь, — Lewis основание, вид, принимающий электронную пару — это кислота Льюиса, а продуктом реакции является аддак Льюиса кислотно-основная. Как показано на двух примерах выше, реакции Брёнстед-Лоури кислотно-основная представляют собой подкатегорию кислых реакций Льюиса, в частности, тех, в которых кислотный вид является H+. Ниже приведено несколько примеров, связанных с другими кислотами и основами Льюиса. 
 Атом бора в трифториде бора, BF3, имеет всего шесть электронов в своей валентной оболочке. Будучи коротким предпочтительным октетом, BF3 является очень хорошей кислотой Льюиса и реагирует со многими основаниями Льюиса; ионом фтора является Lewis основание в этой реакции, жертвуя одной из своих одиночных пар: 
 <img alt="Image2" src="/files/ftp_upload/11407/11407_Image_2.svg" style="height: 150px;" />
 В следующей реакции каждая из двух молекул аммиака, оснований Льюиса, передает пару электронов на серебряный ион, кислоту Льюиса: 
 
 <img alt="Image3" src="/files/ftp_upload/11407/11407_Image_3.svg" style="height: 150px;" />
 Оксиды неметаллов действуют как кислоты Льюиса и реагируют с ионами оксида, основаниями Льюиса, образуя оксианионы: 
 <img alt="Image4" src="/files/ftp_upload/11407/11407_Image_4.svg" style="height: 150px;" />
 Многие реакции Lewis кислотно-основная являются реакциями смещения, в которых один Льюис основание вытесняет другой Льюис основание из аддука кислотно-основная, или в которых одна Льюис-кислота вытесняет другую Льюис-кислоту: 
 <img alt="Image5" src="/files/ftp_upload/11407/11407_Image_5.svg" style="height: 350px;" />
 Другой тип химии Lewis кислотно-основная включает образование сложного иона (или координационного комплекса), состоящего из центрального атома, обычно катиона переходных металлов, окруженного ионами или молекулами, называемыми лигандами. Эти лиганды могут быть нейтральными молекулами, такими как H2O или NH3, или ионами, такими как CN– или OH–. Часто лиганды действуют как основы Льюиса, жертвуя пару электронов на центральный атом. 
 Этот текст адаптирован из<a href="https://openstax.org/books/chemistry-2e/pages/15-2-lewis-acids-and-bases"> Openstax, Химия 2е изд., раздел 15.2: Кислоты и основания Льюиса</a>.

В 1923 году Г. Н. Льюис предложил обобщенное определение кислотно-щелочного поведения, в котором кислоты и основания идентифицируются по их способности воспринимать или ...

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Acids and Bases

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KEY TERMS AND CONCEPTS

ОСНОВНЫЕ ТЕРМИНЫ И ПОНЯТИЯ

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 acid-base adduct. The electron-pair donor is the Lewis base, and the electron-pair acceptor is the Lewis acid.

Brønsted-Lowry acid-base reactions fit inside Lewis acid chemistry. In those reactions, the acid species is H+. Other Lewis acids show the same electron-pair pattern in different ways. Boron trifluoride, BF3, is a strong Lewis acid because boron has only six electrons in its valence shell and still needs an octet. A fluoride ion can donate a lone pair to BF3 and form an adduct.

Lewis acids can also bind more than one base. In one example, two ammonia molecules each donate a pair of electrons to a silver ion, which acts as the Lewis acid. Nonmetal oxides also behave as Lewis acids and react with oxide ions, which are Lewis bases, to form oxyanions. Many Lewis acid-base reactions are displacement reactions. In these reactions, one Lewis base can replace another base in an adduct, or one Lewis acid can replace another acid.

Lewis acid-base chemistry also helps form complex ions, or coordination complexes. These complexes have a central atom, often a transition metal cation, surrounded by ligands. Ligands can be neutral molecules such as H2O or NH3, or ions such as CN- or OH-. In many cases, the ligands act as Lewis bases and donate electron pairs to the central atom.

Electron-Pair Reactions in Acid-Base Chemistry

electron-pair-reactions-in-acid-base-chemistry

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...

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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

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...

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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...

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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

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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

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 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 Lewis acids and bases use electron pairs to form coordinate covalent bonds, adducts, complex ions, and examples like BF3, Ag+, and hydronium.

Electron-Pair Reactions in Acid-Base Chemistry | JoVE K12

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Electron-Pair Reactions in Acid-Base Chemistry

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