Chemistry
Lewis Acid and Bases
Lewis acids and bases are a concept in chemistry that expands the traditional definition of acids and bases. According to the Lewis theory, a Lewis acid is a substance that can accept a pair of electrons, while a Lewis base is a substance that can donate a pair of electrons. This broader definition allows for a more comprehensive understanding of chemical reactions and interactions.
Written by Perlego with AI-assistance
Related key terms
1 of 5
10 Key excerpts on "Lewis Acid and Bases"
eBook - PDF- Ramesh Chandra, Snigdha Singh, Aarushi Singh(Authors)
- 2020(Publication Date)
- Arcler Press(Publisher)
In 1923, G. N. Lewis from UC Berkeley prescribes a substitute theory to describe acids and bases. This theory gave a general description of acids and bases on the basis of bonding and structure. In reference with the Lewis definition of acids and bases, chemists can now ascertain a broader variety of acid-base reactions. Lewis’ theory generally makes use of electrons rather than proton transfer and explicitly specified that an acid is a species which takes an electron pair while on the other hand, a base donates an electron pair. Figure 2.2. Reaction between Lewis Acid and Lewis base. Acids, Bases, and Functional Group Exchange Reactions 37 In the above figure, a Lewis Base (B) donates it electrons to a Lewis Acid (A) resulting in a coordinate covalently bonded compound, also known as an adduct . The reaction of a Lewis acid and a Lewis base will produce a coordinate covalent bond, as shown in figure above. A coordinate covalent bond is just a type of covalent bond in which one reactant gives it electron pair to another reactant. In this case, the Lewis base donates its electrons to the Lewis acid. When they do react this way, the resulting product is called an addition compound, or more cmonly an addict. • Lewis Acid: a species that accepts an electron pair (i.e., an electrophile) and will haveacant orbital’s. • Lewis Base: a species that donates an electron pair (i.e., a nucleophile) and will have lone-pair electrons. 2.3.1. Lewis Acids Lewis acids possess the ability to accept an electron pair. Lewis Acids are Electrophilic, which means they are attracted towards electron. Having bonding with a base, it allows the acid to use its lowest unoccupied molecular orbital orUMO (Figure 2.2). 1. There are several species that can work as a Lewis acid. All cations are Lewis acids as they can easily accept electrons. (e.gFe 2+ , Cu 2+ , Fe 3+ ) 2. An ion, atom, or molecule having an incomplete octet of electrons can work as a Lewis aci(e.g., BF 3 , AlF 3 ).
eBook - PDF- Ageetha Vanamudan(Author)
- 2023(Publication Date)
- Delve Publishing(Publisher)
The conjugate base of weak acids is less likely to supply a proton than the conjugate base of strong Bronsted- Lowry acids. Lewis’ Acids and Bases Theory Lewis defined a Lewis acid as a species with an empty orbital that can accept an electron pair. A Lewis base can act as an electron donor since it has just one pair of electrons. As a result, acids and bases no longer include a hydrogen atom. Lewis acids and bases have two properties: electrophilicity and nucleophilicity. Lewis acids include Cu2+, BF3, and Fe3+, to name a few. F-, NH3, and C2H4 are the three Lewis bases (ethylene). A coordinated covalent bond is formed when an electron pair from a Lewis base accepts an electron pair from a Lewis acid. The final product of this reaction is a Lewis adduct. Using this concept, many chemicals may be classed as acids or bases. This isn’t the ideal site for individuals who want to learn more about acid-base strengths. Acid-base interactions do not result in covalent coordination. 13.7 PH VALUES FOR ACIDS AND BASES The pH scale can be used to determine a substance’s acidity or basicity (pH stands for “potential of hydrogen”). The pH scale is the most popular and accurate technique to evaluate whether a substance is acidic or basic. There are 14 pH scale points in all, with 0 indicating the most acidic solution and 14 representing the most basic solution. A substance’s acidity or alkalinity can be determined using litmus paper (Krylov & Gill, 2013). Litmus paper Acids and Bases 225 can distinguish between acids and bases. Today’s market offers both red and blue litmus paper. When put in an alkaline or basic solution, blue litmus paper becomes scarlet, indicating the presence of acids or bases. Figure 13.8: Acid base test. Source: By Bordercolliez - Own work, CC0, https://commons.wikimedia.org/w/ index.php?curid=16884915 It is critical to understand the distinctions between acids and bases.
eBook - PDF- David R. Klein(Author)
- 2021(Publication Date)
- Wiley(Publisher)
An understanding of the relevant reactions will lead to a greater appreci- ation of food chemistry. In this chapter, our study of acids and bases will serve as an introduction to the role of electrons in ionic reactions. An ionic reac- tion is a reaction in which ions participate as reactants, intermediates, or products. These reactions represent 95% of the reactions covered in this textbook. In order to prepare ourselves for the study of ionic reactions, it is critical to be able to identify acids and bases. We will learn how to draw acid-base reactions and to compare the acidity or basicity of compounds. These tools will enable us to predict when acid-base reactions are likely to occur and to choose the appropriate reagent to carry out any specific acid-base reaction. 3 94 CHAPTER 3 Acids and Bases DO YOU REMEMBER? Before you go on, be sure you understand the following topics. If necessary, review the suggested sections to prepare for this chapter. • Identifying Formal Charges (Sections 1.4 and 2.4) • Drawing and Interpreting Bond-Line Structures (Sections 1.6 and 2.2) • Identifying Lone Pairs (Section 2.5) • Drawing Resonance Structures (Section 2.10) Take the DO YOU REMEMBER? QUIZ in the online course to check your understanding. 3.1 Introduction to Brønsted-Lowry Acids and Bases This chapter will focus primarily on Brønsted-Lowry acids and bases. There is also a brief section dealing with Lewis acids and bases, a topic that will be revisited in Chapter 6 and subsequent chapters. The definition of Brønsted-Lowry acids and bases is based on the transfer of a proton (H + ). An acid is defined as a proton donor, while a base is defined as a proton acceptor. As an example, con- sider the following acid-base reaction: O Cl Cl Acid (proton donor) O Base (proton acceptor) H H H H H H + + - + In the reaction above, HCl functions as an acid because it donates a proton to H 2 O, while H 2 O functions as a base because it accepts the proton from HCl.
eBook - ePubInorganic Chemistry for Geochemistry and Environmental Sciences
Fundamentals and Applications
- George W. Luther, III(Authors)
- 2016(Publication Date)
- Wiley(Publisher)
Chapter 7 Acids and Bases 7.1 Introduction Acids and bases have been known since before Roman times, for their ability to transform a chemical substance into other chemical forms. For example, sulfuric acid (, oil of vitriol; vitriolic acid) could be produced by the aqueous oxidation of pyrite, which is a common mineral used for many purposes [1]. dehydrates sugar as well as metal hydrates with sometimes dramatic color changes, and also dissolves metals via redox reactions. It is normally the chemical that is produced and sold in greatest quantity in the world each year. The first base was called lye, which was obtained by leaching ashes with water producing potassium hydroxide solution. Lye is a common name for bases, which are used in a variety of purposes including wood degradation into paper or fibers and soap production. 7.2 Arrhenius and Bronsted–Lowry Definitions Common acids such as hydrochloric acid (HCl) and nitric acid were not formally prepared until about the 16th century, so formal definitions of acids and bases similar to bonding theories are relatively new. In 1884, Svante Arrhenius defined an acid as a chemical species that when dissolved in water produced the hydrogen ion,. Although useful, the definition is limited as it does not encompass a large variety of reactions and only considers water as solvent. In 1923, Brønsted and Lowry defined an acid and base reaction as one that involves a hydrogen ion transfer between two reactants. The acid is the hydrogen ion donor (the Brønsted acid) and the base is the hydrogen ion acceptor (the Brønsted base). This definition applies to all solvents and the gas phase. For example, the reaction between HCl and can occur in water as solvent or in the gas phase (an atmospheric reaction) and results in complete transfer to form the hydronium ion,
eBook - PDF- Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
- 2022(Publication Date)
- Wiley(Publisher)
A Lewis acid must be able to accept an electron pair to form a new bond. Most Lewis acids are one of the following: a molecule with vacant valence orbitals, a molecule with delocalised bonds involving oxygen, or a metal cation. A hard Lewis base has electron pairs of low polarisability and high electronegativity. A soft Lewis base has a large donor atom of high polarisability and low electronegativity. A hard Lewis acid has an acceptor atom with low polarisability. A soft Lewis acid has a relatively high polarisability. The hard–soft acid–base (HSAB) principle states that hard Lewis acids tend to combine with hard Lewis bases, while soft Lewis acids tend to combine with soft Lewis bases. This is reflected in nature, where hard metal ions tend to be found bonded to hard bases, and vice versa. KEY CONCEPTS AND EQUATIONS Concept Section Description/equation Autoprotolysis constant for water [H 3 O + ][OH - ] = K w 11.2 This equation is used to calculate [H 3 O + ] if [OH - ] is known, and vice versa. [H 3 O + ][OH - ] = K w pH and pOH 11.2 These provide a convenient method for representing [H 3 O + ] and [OH - ] in aqueous solution. Relationship between pH and pOH: pH + pOH = 14.00 at 25 °C 11.2 This relationship can be used to calculate pH if pOH is known, and vice versa. pH + pOH = 14.00 (at 25 °C) K a K b = K w 11.4 This equation is used to calculate K a given K b , or vice versa. Periodic trends in strengths of acids 11.5 Periodic trends can be used to predict the relative acidities of acids. The principles involved also let you compare the acidities of compounds containing different electronegative elements. pH = pK a + log [A - ] [HA] 11.6 This equation can be used to calculate the pH of a buffer solution. Hard–soft acid–base principle 11.8 Hard Lewis acids tend to combine with hard Lewis bases. Soft Lewis acids tend to combine with soft Lewis bases.
eBook - PDFChemistry
Structure and Dynamics
- James N. Spencer, George M. Bodner, Lyman H. Rickard(Authors)
- 2011(Publication Date)
- Wiley(Publisher)
To understand the implications of this definition, look at how the prototypical base, the OH ion, accepts a proton. The only way to accept an H ion is to form a covalent bond to it. In order to form a covalent bond to an H ion that has no valence electrons, the base must provide both of the electrons needed to form the covalent bond. Thus only com- pounds that have pairs of nonbonding valence electrons can act as H ion accep- tors, or Brønsted bases. The following compounds, for example, can all act as Brønsted bases because they all contain nonbonding pairs of electrons. The Brønsted model therefore includes within the category of bases any ion or molecule that contains one or more pairs of nonbonding valence electrons that can accept a proton. Many molecules and ions satisfy the definition of a Brønsted O O O A N H NH 3 H H O Q O O H H 2 O H O B S S S O S S O S S S CO 3 2− 2− O C D G H O OH O O O HO OH O O O H 2 PO 4 - (aq) + H 2 O(l) uv H 3 O + (aq) + HPO 4 2 - (aq) NH 4 + (aq) + OH - (aq) uv NH 3 (aq) + H 2 O(l) HCl(aq) + NH 3 (aq) ¡ Cl - (aq) + NH 4 + (aq) 11.3 THE BRØNSTED–LOWRY DEFINITION OF ACIDS AND BASES 471 base. However, the following substances are not Brønsted bases because they have no nonbonding valence electrons. A O O A C H H CH 4 H H H 2 O H H A O O A N H NH 4 H H H 472 CHAPTER 11 / ACIDS AND BASES E x e r c i s e 1 1 . 1 Identify the reactant that behaves as a Brønsted acid and the reactant that behaves as a Brønsted base in each of the following reactions. (a) (b) (c) Solution (a) acid: HF base: OH (b) acid: CH 3 CO 2 H base: H 2 O (c) acid: HNO 3 base: C 6 H 5 NH 2 C 6 H 5 NH 2 (aq) + HNO 3 (aq) uv C 6 H 5 NH 3 + (aq) + NO 3 - (aq) CH 3 CO 2 H(aq) + H 2 O(l) uv CH 3 CO 2 - (aq) + H 3 O + (aq) HF(aq) + OH - (aq) uv H 2 O(l) + F - (aq) 11.4 Conjugate Acid–Base Pairs An important consequence of the Brønsted theory is the recognition that acids and bases are linked or coupled to form conjugate acid–base pairs.
eBook - PDFChemistry
An Industry-Based Introduction with CD-ROM
- John Kenkel, Paul B. Kelter, David S. Hage(Authors)
- 2000(Publication Date)
- CRC Press(Publisher)
This definition of an acid is similar to that in the Arrhenius Theory (a hydrogen ion is a proton) and, thus, all Arrhenius acids are also Bronsted-Lowry acids and vice versa. However, the Bronsted-Lowry definition of a base is broader. There are substances other than Arrhenius bases that accept hydrogen ions. In other words, the hydroxide group is a hydrogen ion acceptor, but there are other groups that also do this. What exactly is meant by a hydrogen ion acceptor? A hydrogen ion acceptor is a chemical species to which a hydrogen ion can attach and form a covalent bond. The most important examples, in addition to the Arrhenius bases (hydroxides), are ammonia, NH 3 , and the carbonates (see Table 12.4). Examples of reactions in which hydrogen ions are accepted are the following. (12.7) (12.8) (12.9) (12.10) (12.11) Notice that when hydroxides accept hydrogen ions, water is always the product. Water is not a product in the cases of ammonia and the carbonates and bicarbonates. However, in the case of carbonates and bicarbonates, water often does ultimately form. The carbonic acid, H 2 CO 3 , in the above reactions is unstable and will spontaneously decompose into water and carbon dioxide. This reaction is often accom-panied by vigorous gas (CO 2 ) evolution and will be discussed in more detail in Section 12.6. The list of Bronsted-Lowry bases is actually much longer than the list that is given here. There are many anions like carbonate and bicarbonate that will form a covalent bond with hydrogen ions and form a weak acid like the carbonic acid above. Thus, the anion of any weak acid qualifies as a Bronsted-Lowry base. This list would include sulfide and cyanide, as well as acetate and anions of other carboxylic acids. In Section 12.2 we mentioned that water can be thought of as an acid because in a sample of pure distilled water, water molecules release hydrogen ions to a small extent.
- Peter V. Hobbs(Author)
- 2000(Publication Date)
- Cambridge University Press(Publisher)
Thus, alkalis are soluble strong bases, such as NaOH and KOH. Alkali metal means any metal of the group Na, K, Li, Rb, and Cs. In general, the oxides of metallic elements are basic, and the oxides of nonmetallic elements are acidic. 2 The proton has a radius of about 10 15 m. Other ions that have electrons associated with them have radii of about 10~ 10 m. 3 Those familiar with the structure of atoms would expect from these definitions that acids are materials in which the outer electron orbitals are not completely filled with electrons, and bases are materials with electrons available for sharing. For example, from the Lewis viewpoint, H + is an acid because it has an empty orbital that can accept a pair of elec-trons, and OH is a base because it has pairs of electrons available for sharing. 4 The symbol pH was introduced by a Danish chemist, S. S0rensen; p stands for the Danish word for power and H for hydrogen. With a change in sign, pH is the power of ten of the hydrogen ion concentration in moles per liter. 5 Another way of viewing this is that hydrolysis is the reverse of acid dissociation [cf. Reac-tion (5.20)]. Thus, the weaker the acid [e.g., H2QH3O2 in Reaction (5.20)] the more dif-ficult it is to remove a proton from it, and the easier it is for its anion or conjugate base [e.g., C 2 H 3 O2(aq) in Reaction (5.20)] to attach a proton from water (i.e., to hydrolyze). 6 In calculations such as this, where several approximations are made, the solutions should be checked by substituting the derived values back into the original equations to see if reasonable equalities are obtained. This is left as an exercise for the reader for this case.
eBook - PDFChemistry
The Molecular Nature of Matter
- Neil D. Jespersen, Alison Hyslop(Authors)
- 2021(Publication Date)
- Wiley(Publisher)
• compare the strengths of Brønsted–Lowry acids and bases, how they are classified as strong or weak, and the strengths of conjugate acid–base pairs. • using the periodic table, describe the trends in the strengths of binary acids and oxoacids. • define Lewis acids and bases and compare them to the Arrhenius and Brønsted–Lowry definitions. • using the periodic table, describe which elements are most likely to form acids or bases. • describe how the production of advanced ceramics depends on acid–base chemistry, including the sol-gel process. Acids and Bases: A Molecular Look CHAPTER 15 Yuri Smityuk / Getty Images 760 CHAPTER 15 Acids and Bases: A Molecular Look FIGURE 15.1 The reaction of gaseous HCl with gaseous NH 3 . As each gas escapes from its concentrated aqueous solution and mingles with the other, a cloud of microcrystals of NH 4 Cl forms above the bottles. 1 When the single electron is removed from a hydrogen atom, what remains is just the nucleus of the atom, which is a proton. Therefore, a hydrogen ion, H + , consists of a proton, and the terms proton and hydrogen ion are often used interchangeably. Andy Washnik CONNECTING TO THE LEARNING OBJECTIVES From the beginning of this book, we have been suggesting that you outline the chapter that you are studying and that you work through the questions in the back of the chapter to practice solving chemistry problems. We have also suggested that you use a notebook dedicated to chemistry to keep your notes. Here are a few additional suggestions to help you learn the material in this chapter. 1. Take notes in class and when you read the chapter. Notes are messy, full of questions, have diagrams and pictures, include your reactions to the material, and draw connections between concepts. Sometimes they even include doodles. 2. Choose two concepts from the chapter and prepare a lesson to teach the material to someone else.
eBook - PDFIntroductory Chemistry
An Active Learning Approach
- Mark Cracolice, Edward Peters, Mark Cracolice(Authors)
- 2020(Publication Date)
- Cengage Learning EMEA(Publisher)
1. An acid–base reaction is a transfer of protons; a redox reaction is a transfer of electrons. 2. In both cases, the reactants are given special names to indicate their roles in the transfer process. An acid is a proton source; a base is a proton remover. A reducing agent is an electron source; an oxidizing agent is an electron remover. 3. Just as certain species can either provide or remove protons (e.g., HCO 3 2 and H 2 O) and thereby behave as an acid in one reaction and a base in another, certain species can either remove or provide electrons, acting as an oxidizing agent in one reaction and a reducing agent in another. An example is the Fe 21 ion, which can oxidize Zn atoms to Zn 21 in the reaction Fe 21 saqd 1 Znssd S Zn 21 saqd 1 Fessd Fe 21 can also reduce Cl 2 molecules to Cl 2 ions in another reaction: 2 Fe 21 saqd 1 Cl 2 sgd S 2 Cl 2 saqd 1 2 Fe 31 saqd 4. Just as acids and bases may be classified as stronger or weaker depending on how readily they remove or provide protons, the relative strengths of oxidizing and reducing agents may be compared according to their tendencies to attract or release electrons. 5. Just as most acid–base reactions in solution reach a state of equilibrium, most aqueous redox reactions also reach equilibrium. Just as the favored side of an acid–base equilibrium equation can be predicted from acid–base strength, the favored side of a redox equilibrium equation also can be predicted from oxidizing agent–reducing agent strength. 17.9 The Water Equilibrium Goal 12 Given the hydrogen ion or hydroxide ion concentration of water or a water solution, calculate the other value. In the remaining sections of this chapter, you will be multiplying and dividing expo- nentials, taking the square root of an exponential, and working with logarithms. We will furnish brief comments on these operations as we come to them. For more detailed instructions, see Appendix I, Parts A and C.
Index pages curate the most relevant extracts from our library of academic textbooks. They’ve been created using an in-house natural language model (NLM), each adding context and meaning to key research topics.
