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Acid-Base Chemistry

Acid-Base Chemistry

Acid-base chemistry is a fundamental aspect of chemical science that plays a crucial role in our daily lives. Its applications range from industrial processes to biological systems. Understanding acid-base chemistry is not just essential for scientists, but also for everyday life, as it helps in making informed decisions about health, nutrition, and environmental issues.

Acids and Bases: Definitions

Acids and bases are two types of compounds that readily react with one another.

  • Acids are substances that donate protons (H⁺ ions) or accept electron pairs. Common examples include vinegar (acetic acid, CH₃COOH), citrus fruits (citric acid, C₆H₈O₇), and stomach acid (hydrochloric acid, HCl). The properties of acids include a sour taste, ability to turn blue litmus paper red, and corrosiveness.
  • Bases , on the other hand, are substances that accept protons or donate electron pairs. Examples include baking soda ( sodium bicarbonate , NaHCO₃), soap (sodium hydroxide, NaOH), and ammonia (NH₃). Bases are characterized by a bitter taste, slippery feel, and the ability to turn red litmus paper blue.

Acid-Base Chemistry and the pH Scale

The pH scale , ranging from 0 to 14, measures the acidity or alkalinity (basicity) of a solution. A pH less than 7 indicates acidity, while a pH greater than 7 indicates alkalinity. A pH of 7 is neutral. Pure water is an example of a substance with a neutral pH.

Acid-Base Chemistry Theories

Acid-Base Theories

The three main theories of acids and bases are the Arrhenius theory, Brønsted-Lowry theory, and Lewis theory. Each of these theories has its uses in chemistry.

  • Acids : Substances that increase the concentration of H⁺ ions in water.
  • Bases : Substances that increase the concentration of OH⁻ ions in water.
  • Acids : Proton donors.
  • Bases : Proton acceptors.
  • Acids : Electron pair acceptors.
  • Bases : Electron pair donors.

Strength of Acids and Bases

One way of classifying acids and bases is as strong or weak:

  • Strong Acids and Bases : These dissociate completely in water. Examples include hydrochloric acid (HCl) and sodium hydroxide (NaOH).
  • Weak Acids and Bases : These partially dissociate in water. Examples include acetic acid (CH₃COOH) and ammonia (NH₃).

Acid-Base Reactions and Neutralization

Acid-base reactions typically involve the transfer of protons from acids to bases. Neutralization is a specific type of acid-base reaction where an acid and a base react to form water and a salt , effectively neutralizing each other.

The outcome of an acid-base reaction depends on the strength of the acids and bases.

  • Strong Acid with Strong Base : This leads to complete neutralization, forming a neutral salt and water. Example: HCl (acid) + NaOH (base) → NaCl (salt) + H₂O (water).
  • Strong Acid with Weak Base : The resulting solution is slightly acidic, as the weak base cannot completely neutralize the strong acid. Example: HCl (acid) + NH₃ (base) → NH₄Cl (salt) + H₂O (water).
  • Weak Acid with Strong Base : The resulting solution is slightly basic, as the strong base completely neutralizes the weak acid. Example: CH₃COOH (acid) + NaOH (base) → CH₃COONa (salt) + H₂O (water).
  • Weak Acid with Weak Base : This leads to partial neutralization, with the pH of the resulting solution depending on the relative strengths of the acid and base. Example: CH₃COOH (acid) + NH₃ (base) → CH₃COONH₄ (salt) + H₂O (water).

Buffers in Acid-Base Chemistry

A buffer is a solution that resists changes in pH when small amounts of an acid or a base are added. This property is essential in various chemical, biological, and environmental contexts where maintaining a stable pH is critical.

Buffers typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid. This dual presence allows the buffer to neutralize added acids or bases. For instance, in a buffer composed of acetic acid (CH₃COOH) and its conjugate base, sodium acetate (CH₃COONa), the acetic acid neutralizes added bases while the sodium acetate neutralizes added acids.

The buffer capacity refers to the amount of acid or base a buffer solution can absorb without a significant change in pH. This capacity depends on the concentration of the buffer components and the closeness of the solution’s pH to the pKa (acid dissociation constant) of the buffer acid.

Frequently Asked Questions (FAQs) on Acid-Base Chemistry

What is the difference between a strong acid and a weak acid?

  • A strong acid completely dissociates into its ions in water, releasing all of its hydrogen ions. Examples include hydrochloric acid (HCl) and sulfuric acid (H₂SO₄). A weak acid only partially dissociates in water, leaving many of its hydrogen ions un-released. Acetic acid (CH₃COOH) and citric acid (C₆H₈O₇) are common examples.

Can you explain what a pH of 7 means?

  • A pH of 7 is neutral, meaning the solution is neither acidic nor basic. It indicates a balance between hydrogen ions (H⁺) and hydroxide ions (OH⁻) in water. Pure water at 25°C (77°F) has a pH of 7.

Why is pH important in daily life?

  • pH plays a crucial role in everyday life. It affects food taste, digestion, skin health, pool water safety, garden soil quality, and even the functioning of batteries and car fluids.

What are some common household acids and bases ?

  • Common household acids include vinegar (acetic acid), lemon juice (citric acid), and battery acid (sulfuric acid). Household bases include baking soda (sodium bicarbonate), bleach (sodium hypochlorite), and ammonia-based cleaners.

How do buffers work?

  • Buffers work by using a weak acid and its conjugate base (or a weak base and its conjugate acid) to resist changes in pH. When you add an acid or base to the buffer, the buffer components react to neutralize the added substance. This keeps the pH relatively stable.

What is an acid-base neutralization reaction?

  • An acid-base neutralization reaction occurs when an acid and a base react to form water and a salt. This reaction typically decreases the solution’s acidity or basicity.

How are acids and bases used in industries?

  • Acids and bases have wide industrial applications. Acids find use in metal processing, fertilizer production, and petroleum refining. Bases are important in soap and detergent manufacturing, textile processing, and papermaking.

What safety precautions should be taken when handling acids and bases?

  • When handling acids and bases, wear protective gear, work in a well-ventilated area, and be aware of proper storage and disposal methods. Avoid skin contact and inhalation of fumes. In case of spills, neutralize the acid or base and clean up the spill safely. Always have a first aid kit and emergency protocols in place.
  • Finston, H.L.; Rychtman, A.C. (1983).  A New View of Current Acid-Base Theories . New York: John Wiley & Sons.
  • Masterton, William; Hurley, Cecile; Neth, Edward (2011).  Chemistry: Principles and Reactions . Cengage Learning. ISBN 978-1-133-38694-0.
  • Paik, Seoung-Hey (2015). “Understanding the Relationship Among Arrhenius, Brønsted–Lowry, and Lewis Theories”.  Journal of Chemical Education . 92 (9): 1484–1489. doi: 10.1021/ed500891w
  • Petrucci, R. H.; Harwood, R. S.; Herring, F. G. (2002).  General Chemistry: Principles and Modern Applications  (8th ed.). Prentice Hall. ISBN 0-13-014329-4.

Related Posts

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Course info.

  • Prof. Donald Sadoway

Departments

  • Materials Science and Engineering

As Taught In

  • Chemical Engineering

Learning Resource Types

Introduction to solid state chemistry, 26. acids & bases.

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Session Overview

Prerequisites.

Before starting this session, you should be familiar with:

  • the Molecules and Bonding module ( Session 7 through Session 12 )
  • Solutions ( Session 25 )

Learning Objectives

After completing this session, you should be able to:

  • Compare the acid-base models of Arrhenius, Brønsted-Lowry, and Lewis, and know the salient features of each.
  • Write the general acid-base reaction.
  • Explain the chemical basis of acid strength and the pH measurement.
  • Describe the behaviors of conjugate acid-base pairs.
  • Explain how dissociation affects ionic compounds.
  • Solve specific acid-base reaction problems.

Lecture Video

  • Download video
  • Download transcript

Lecture Slides (PDF - 2.2MB)

Lecture Summary

This lecture introduces the chemical models and behaviors of acids and bases. Starting from the historical origins (“acid” derives from the Latin acidus , meaning “sour”), Prof. Sadoway discusses the evolving acid-base models of Lavoisier (1776), Arrhenius (1887), Brønsted and Lowry (1923), and Lewis (1923-1938).

The lecture proceeds to cover:

  • The general acid-base reaction
  • Conjugate acid-base pairs
  • The dissociation process of ionic compounds
  • Solving acid-base reaction problems

Problems (PDF)

Solutions (PDF)

Textbook Problems

For further study, supplemental readings.

Djerassi, C., and R. Hoffmann. Oxygen: A Play in Two Acts . New York, NY: Wiley-VCH, 2001. ISBN: 9783527304134. See also the study guide and other publisher resources .

Brandis, Kerry. Acid-Base Physiology . See Chapter 1 of this online tutorial/textbook, which applies acid-base chemistry to physiology.

Antoine Lavoisier

Svante Arrhenius — 1903 Nobel Prize in Chemistry

Johannes Nicolaus Brønsted

Martin Lowry

Gilbert N. Lewis

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How to teach acids, bases and salts

By Naomi Hennah 2018-10-10T13:25:00+01:00

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Naomi Hennah suggests ideas, resources and activities for your classroom

Sodium chloride is the source of sodium in our diets, essential for the transmission of nerve impulses and the maintenance of a proper fluid balance in the body. Throughout history, humans have been using this salt to preserve meat, clean wounds and make soap.

Sodium chloride is one example of a salt. In chemistry, the term salt refers to a group of ionic compounds formed from the neutralisation reaction between an acid and a base.

The concepts of acids, bases and salts are introduced early in secondary school science, are developed and refined as students progress, and underpin many future topics. Here are some ideas to engage students, avoid misconceptions, and connect practical work to underlying concepts.

What students need to know

  • Acids are hydrogen-containing substances with a sour taste that form solutions with pH values less than 7. Common examples include hydrochloric acid, sulfuric acid, citric acid and ethanoic acid (vinegar/acetic acid).
  • Bases are a group of substances that neutralise acids.
  • Soluble bases are called alkalis. They have a slippery, soapy feel and form solutions with pH values greater than 7. Common examples include sodium hydroxide, magnesium hydroxide, sodium hydrogen carbonate (sodium bicarbonate), sodium hypochlorite and ammonia.
  • Neutralisation is a reaction between an acid and an alkali that forms a salt and water.
  • Salts are odourless and have a salty taste, and many are soluble in water. Common examples include sodium chloride, potassium iodide, calcium carbonate and copper sulfate.
  • The pH scale is used to measure acidity and alkalinity.
  • Indicators are substances that change colour with a change in acidity/alkalinity. Litmus is a common indicator; alkaline solutions turn red litmus blue and acid solutions turn blue litmus red.
  • Acids can react with some metals to form a salt and hydrogen gas.

Ideas for the classroom

Students have a wealth of experience of acids, bases and salts accumulated from both school and everyday life. It is worth starting the topic by eliciting their existing ideas by constructing a mind-map together. Be alert to miscomprehensions and aim to address these before moving on.

Using real-life examples and anecdotes in your classroom can help make ideas stick. Download a selection of acid, base and salt anecdotes ( MS Word or pdf ).

Using real-life examples and anecdotes in your classroom can help make ideas stick. Download a selection of acid, base and salt anecdotes from the  Education in Chemistry  website: rsc.li/2Oj0lQk.

One particular problem that can arise when introducing laboratory acids and alkalis is that they both look like water. Students can struggle to use chemical properties to characterise these solutions. The ‘Sage and scribe’ task can be used to demonstrate the limitations of visual description, and reinforce the need to look for the presence or absence of defined characteristics or properties. Download the student information ( MS PowerPoint or pdf ) and teacher notes ( MS Word or pdf ) for this activity.

One particular problem that can arise when introducing laboratory acids and alkalis is that they both look like water. Students can struggle to use chemical properties to characterise these solutions. The ‘Sage and scribe’ task can be used to demonstrate the limitations of visual description, and reinforce the need to look for the presence or absence of defined characteristics or properties. Download the student information and teacher notes for this activity: rsc.li/2Oj0lQk.

Universal indicator paper

Source: Bjoern Wylezich / Shutterstock.com

In this activity, the sage has to describe a simple image to two scribes. One scribe records the sage’s description, while the other tries to make a drawing from it. Can the scribes identify the image? Descriptions are subjective and can be misinterpreted, whereas identification requires an objective approach.

Link this idea to how we can positively distinguish between two solutions that look like water, specifically with the use of indicators with acids and alkalis.

Aim to introduce acids and alkalis equally rather than focus on acids alone. For homework, ask students to identify household substances that are acids and alkalis. These will be commonly found in the kitchen and bathroom. For example, vinegar and lemon juice are acids whereas baking powder and toothpaste are alkalis. Student can bring in samples and test them using the Acid or alkali? Acidic or alkaline?  activity. The Colour creactions  activity can then be used to introduce universal indicator and the pH scale before moving on to neutralisation.

Aim to introduce acids and alkalis equally rather than focus on acids alone. For homework, ask students to identify household substances that are acids and alkalis. These will be commonly found in the kitchen and bathroom. For example, vinegar and lemon juice are acids whereas baking powder and toothpaste are alkalis. Student can bring in samples and test them using the ‘Acid or alkali? Acidic or alkaline?’ activity (rsc.li/2PueBT7). The ‘Colour reactions’ activity can then be used to introduce universal indicator and the pH scale before moving on to neutralisation (rsc.li/2CaCriy).

It is important to consider what you want students to learn from these activities and how this can be achieved. Time needs to be split equally between hands-on activities (including watching demonstrations and videos) and the subsequent minds-on activity of discussing the underlying concepts and ideas. Students can work collaboratively to build their understanding using structured talk activities that scaffold the making of links between practical work and the underlying concepts.

Linking practical tasks to underpinning concepts

Remember that time allocated to practical work should be split equally between hands-on activities and making links between phenomena and the underlying concepts.

Encourage students to both observe and interpret by asking questions. Try repeating their answers using the correct language. For example, replace ‘made a hole’ or ‘burnt through’ with ‘corroded’.

With practice, students can manage their own practical talk within their group. Structured talk activities should have a strict time limit. Scaffolding, such as sentence stems, encourage both listening and responding to each other.

  • What did you observe?
  • What can you infer/work out from this?
  • What do you think of …?
  • What are the implications for …?
  • How might …?

Sentence stems

  • I think this shows …
  • This suggests that …
  • I think this tells us …
  • But what about … ?
  • I disagree with you because …
  • I don’t understand …

Common misconceptions

Students often believe that only acids are corrosive and can be identified because they eat away/burn materials. However, alkalis can also be highly corrosive, as demonstrated in the Coke cans in acid and base  video.

Students often believe that only acids are corrosive and can be identified because they eat away/burn materials. However, alkalis can also be highly corrosive, as demonstrated in the ‘Coke cans in acid and base’ video (bit.ly/2C5YyXv).

Flasks with coloured liquid

Source: Rattiya Thongdumhu / Shutterstock.com

Use the video as a prompt to draw out ideas of corrosion. Explain that corrosion can be defined as ‘the degradation or breakdown of a material due to a reaction with its environment’. Degradation is a change in the bulk properties of a material, which will look different, become weaker or even break apart because of chemical changes. This description encourages students to see materials at both the macroscopic (bulk) level and the submicroscopic (particle) level . For example, the sodium hydroxide solution is added to the aluminium can, which breaks apart with the release of gas (macroscopic), because the sodium hydroxide particles are reacting with the aluminium particles to form an aluminium salt and hydrogen gas (sub-microscopic).

Use the video as a prompt to draw out ideas of corrosion. Explain that corrosion can be defined as ‘the degradation or breakdown of a material due to a reaction with its environment’. Degradation is a change in the bulk properties of a material, which will look different, become weaker or even break apart because of chemical changes. This description encourages students to see materials at both the macroscopic (bulk) level and the submicroscopic (particle) level (find out more about getting your students to think about how they learn: rsc.li/2C596X0). For example, the sodium hydroxide solution is added to the aluminium can, which breaks apart with the release of gas (macroscopic), because the sodium hydroxide particles are reacting with the aluminium particles to form an aluminium salt and hydrogen gas (sub-microscopic).

Try to avoid the use of anthropomorphic descriptions such as ‘attack’ or ‘eat’ away. These terms tend to lead to ideas about chemicals ‘wanting’ or ‘needing’ to react. This will make it harder for students to develop a solid understanding of how and why chemical reaction proceed.

Formative assessment

Concept mapping is a useful tool for reinforcing how this topic links across the curriculum. The Revising acids  activity has been adapted as an Assessment for Learning activity .

Concept mapping is a useful tool for reinforcing how this topic links across the curriculum. The ‘Revising acids’ activity (rsc.li/2OTINtu) has been adapted as an Assessment for Learning activity (rsc.li/2A365VQ).

Provide students with opportunities to practice both longer answer (four and six mark) and multiple-choice questions (one mark) in preparation for exams. Display the question and give the students two minutes to write their answer on mini-whiteboards. Importantly, ask them to write why they chose their answer. This provides the opportunity to identify miscomprehension and provide immediate feedback.

Progression to 14–16

In 14–16 teaching, a more sophisticated model of acidity is used, based on hydrogen and hydroxide ions. Acids release hydrogen ions (H + ) in solution, and alkalis hydroxide ions (OH - ).

pH is formally defined as a logarithmic measure of hydrogen ion concentration. Neutralisation is defined as the reaction of hydrogen ions and hydroxide ions to produce water. The distinction between dilute/concentrated (amount of substance) and weak/strong (degree of ionisation) is also made. Finally, there are specific apparatus and techniques that must be used and understood including rates of reaction and titration (see our guides to practical work in GCSE specifications ).

pH is formally defined as a logarithmic measure of hydrogen ion concentration. Neutralisation is defined as the reaction of hydrogen ions and hydroxide ions to produce water. The distinction between dilute/concentrated (amount of substance) and weak/strong (degree of ionisation) is also made. Finally, there are specific apparatus and techniques that must be used and understood including rates of reaction and titration (see our guides to practical work in GCSE specifications: rsc.li/2pIPwbD).

  • Acidic and alkaline solutions can be identified by their chemical properties.
  • Enforce the particulate nature of acids, bases and salts so students move away from just referring to bulk properties and anthropomorphic descriptions such as ‘ate holes in’.
  • Make explicit the link between practical work and underpinning concepts using the macroscopic and sub-microscopic representations.
  • This topic underpins many future topics including synthesis and analysis of chemical substances.

Acids, bases and salts: anecdotes

Sage and scribe: student activity, sage and scribe: teacher notes.

  • Acids and bases

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  • Acids, Bases and Salts

Acids, Bases, and Salts are the main chemical compounds that exist in our surroundings. Acids, Bases, and Salts are compounds that occur naturally and can also be created artificially. They are found in various substances including our food. Vinegar or acetic acid is used as a food preservative. Citrus fruits have citric acid and etc.

Other than food they also have a wide variety of uses such as in various industries, manufacturing plants, processing plants, laboratories, and others. In this article, we will learn about Acids, Bases, and Salts, their properties, types, uses, and others in detail. The image given below shows acid and base which when reacting form salt.

Table of Content

What are Acids?

Types of acids, uses of acids, what are bases, types of bases, uses of bases, what are salts, types of salts.

Acid, Base and Salts

Acids are a type of chemical substance that is characterized by their ability to donate hydrogen ions (H+) when dissolved in water. Acids can be found naturally in many foods and beverages, including citrus fruits, vinegar, and fermented products, and they are also used in various industrial processes. Here are some key points about acids:

  • Acids can be classified into two categories: organic acids, which are derived from living organisms, and inorganic acids, which are derived from non-living sources.
  • Acids have a sour taste and can cause a tingling or burning sensation on the tongue or skin.
  • The strength of an acid is determined by its pH value, which ranges from 0 to 14. Acids with a pH less than 7 are considered to be acidic, while those with a pH greater than 7 are basic or alkaline.
  • Acids can react with bases to form salts and water in a process called neutralization.
  • Acids are used in a variety of industrial applications, such as the production of fertilizers, dyes, and pharmaceuticals.
  • Some acids can be harmful or corrosive to living tissue, such as hydrochloric acid, which is found in the stomach and is necessary for digestion, but can cause severe burns if it comes into contact with the skin.
  • Acids are also used in various chemical reactions, such as in the production of polymers and plastics, as well as in the cleaning and sterilization of medical equipment.

Acid

Physical Properties of Acids

Acids are chemical substances that have unique physical and chemical properties. Here are some of the physical properties of acids:

Acids have a sour taste: Most acids have a distinctly sour taste, such as lemon juice and vinegar.

Acids change the color of indicators: Acids change the color of certain indicators, such as litmus paper, which turns red in the presence of an acid.

Acids are corrosive: Many acids are corrosive and can cause damage to living tissue or corrode metal and other materials.

Acids have a low pH: The pH scale measures the acidity or alkalinity of a solution, with a pH of 7 being neutral, and lower pH values indicating greater acidity. Acids typically have a pH of less than 7.

Acids react with bases to form salts and water: When an acid reacts with a base, the two substances neutralize each other, forming a salt and water.

Acids have a higher boiling point than water: Acids generally have a higher boiling point than water, which means that they require more energy to boil than water.

Acids are good conductors of electricity: In a solution, acids can conduct electricity because they contain charged particles known as ions.

Chemical Properties of Acid

Acid has various chemical properties few of the following chemical properties of acids include, 

Reaction of acids with metal: When an acid reacts with a metal, it produces hydrogen gas and the corresponding salt. 

Metal + Acid → Salt + Hydrogen

Example: When hydrochloride acid combines with zinc metal, it produces hydrogen gas and zinc chloride.

Zn + 2HCl  → ZnCl 2 + H 2  

Reaction of acids with metal carbonate: When acids react with metal carbonates, they produce carbon dioxide gas and salts as well as water.

Metal carbonate + Acid → Salt + Carbon dioxide + Water

Example: When hydrochloric acid combines with sodium carbonate, it produces carbon dioxide gas, sodium chloride, and water.

 Na 2 CO 3 + 2HCl → 2NaCl + H 2 O + CO 2

Reaction of acid with hydrogen carbonates (bicarbonates): When acids react with metal hydrogen carbonates, they produce carbon dioxide gas, salt, and water.

Acid + Metal hydrogen carbonate → Salt + Carbon dioxide + Water

Example : Sulfuric acid gives sodium sulfate, Carbon dioxide gas and water when it reacts with sodium bicarbonate.

2NaHCO 3 + H 2 SO 4 → NaCl + CO 2 + H 2 O

Acids are classified into different categories, the classification of acids is discussed below in the article.

On the basis of their Occurrence  

On the basis of their Occurrence  acid are subdivided into two categories

Natural Acid

Mineral Acids

Natural acids, often known as organic acids, are acids derived from natural sources. For example Methanoic acid (HCOOH), Acetic acid (CH 3 COOH), Oxalic acid (C 2 H 2 O 4 ), etc.

Mineral acids are acids that are created from minerals. Inorganic acids, man-made acids, and synthetic acids are all examples of Mineral Acids. For example Hydrochloric acid (HCl), Sulphuric acid (H 2 SO 4 ), Nitric acid (HNO 3 ), Carbonic acid (H 2 CO 3 ), Phosphoric acid (H 3 PO 4 ), etc.

On the basis of Concentration

On the basis of Concentration, acids are categorized into two categories

  • Strong Acid

Strong Acids

Strong Acid is an acid that is totally ionized in water and produces (H + ). For example Hydrochloric acid (HCl), Sulphuric acid (H 2 SO 4 ), Nitric acid (HNO 3 ) etc.

A weak acid is partially ionized in water, creating a tiny amount of hydrogen ions (H + ). For example Acetic acid (CH 3 COOH), Carbonic acid (H 2 CO 3 ) etc. 

Acids have various uses some of the important uses of acid are,

  • Vinegar is a diluted solution of acetic acid that has a variety of uses in the home. It’s mostly utilized in the food industry as a preservative.
  • Orange and lemon juice contain a significant amount of citric acid. It can also be used for food preservation.
  • In batteries, sulfuric acid is commonly utilized. This acid is typically found in the batteries used to start vehicle motors.
  • Sulfuric and nitric acid is used in the industrial production of dyes, explosives, paints, and fertilizers.
  • Many soft drinks contain phosphoric acid as the main ingredient.

Bases are chemical compounds that react chemically with acids, they produce salts and hydroxide ions (OH – ) in water. For example Potassium hydroxide (caustic potash or KOH), Calcium hydroxide (Ca(OH) 2 ), Sodium hydroxide (caustic soda or NaOH) etc. The image given below shows a base in its aqueous medium.

Base

Physical Properties of Base

Bases have specific physical properties and they can easily be distinguished by their physical properties. Some physical properties of bases are:

  • Base has a bitter taste
  • Bases are soapy to touch
  • Base change red litmus to blue
  • Aquous solution of base conducts electricity

Chemical Properties of Base

Bases have various chemical properties few of the following chemical properties of bases are,

Reaction of Base with Metals: When alkali (base) reacts with metal, salt and hydrogen gas is produced.

Alkali + Metal → Salt + Hydrogen

Example: When sodium hydroxide interacts with aluminium metal, sodium aluminate and hydrogen gas are generated.

2NaOH + 2Al + 2H 2 O → 2NaAlO 2 + 2H 2

Reaction of Non-Metallic Oxides with Base : Salt and water are formed when non-metallic oxides react with a base.

Non-metallic oxide + Base → Salt + Water

Example: When calcium hydroxide reacts with carbon dioxide calcium carbonate is formed along with water.

Ca(OH) 2 + CO 2 → CaCO 3 + H 2 O

Action of Alkalis/Base with Ammonium Salts: Ammonia is produced when alkalis react with ammonium salts.

Alkali + Ammonium salt   →   Salt   +  Water  +  Ammonia

Example: When calcium hydroxide reacts with ammonium chloride, calcium chloride, water, and ammonia are produced.

Ca(OH) 2 + NH 4 Cl  →  CaCl 2 + H 2 O + NH 3

Acidity, concentration, and degree of ionization are three variables that can be used to classify bases.

Types of Bases Based on Acidity

The number of hydroxyl ions presents determines acidity in bases. Based on acidity, bases are classified into three categories:

Mono-acidic bases are those that contain only one hydroxyl ion and interact with only one hydrogen ion. Mono-acidic bases include NaOH, KOH, and others.

Diacidic base is a base with two hydroxyl ions that interact with two hydrogen ions. Ca(OH) 2 , Mg(OH) 2 , and other di-acidic bases are examples.

Triacidic base is a type of base that comprises three hydroxyl ions and three hydrogen ions. Triacidic bases include Al(OH) 3 , Fe(OH) 2 , and others.

Types of Bases Based on their Concentration

Based on their concentration in an aqueous solution, bases are divided into two categories:

  • Concentrated

Concentrated: The concentration of base in these types of bases is higher in the solution. Concentrated NaOH solution, for example.

Diluted: These types of bases have a lower concentration of base in their aqueous solution. For instance, dilute NaOH, dilute KOH, and so on.

Types of Bases Based on their Degree of Ionization

The degree of ionization of bases in solution can be used to classify them. It’s also known as foundation strength. When dissolved in water, it produces a certain quantity of hydroxyl ions. The degree of ionization distinguishes two types of bases.

  • Strong Base

Strong Base: A strong base is one that dissociates entirely or to a large extent in water. For example, NaOH, KOH, and strong bases.

Weak Base: A weak base is one that does not dissolve entirely or only dissociates to a very little level. For example, NH 4 OH, and others are weak bases.

Base has various uses some of the important uses of the base are,

  • Sodium hydroxide is used in the making of paper and soap. Sodium hydroxide (NaOH) is also utilized in the production of rayon.
  • Bleaching powder is made from Ca(OH) 2 , commonly known as calcium hydroxide or slaked lime.
  • Calcium hydroxide is used to create dry mixtures for painting and decorating.
  • Magnesium hydroxide, popularly known as milk of magnesia, is a laxative that is extensively used. It is also used as an antacid since it decreases excess acidity in the human stomach.
  • In laboratories, ammonium hydroxide is a critical reagent.
  • Slaked lime can be used to neutralize any excess acidity in soils.

Bases that are easily dissolved in water are called Alkali, in other words, water-soluble bases are called Alkali. For example, NaOH is an alkali as it dissolves in water forming Na + and OH – ions.

Difference between Alkali and Base

The difference between Alkali and Base can easily be understood with the help of the table given below,

Learn, Acids and Bases

Arrhenius’s Theory of Acid and Base

What is acid, what is base, and what is the difference between acid and base? these questions are nightmares to chemists in the early 15 and 16 centuries. To solve these questions a chemist name Arrhenius came up with a theory called  Arrhenius theory. According to Arrhenius’s theory, a substance that gives H + ion in its aqueous solution is called acid whereas the substance that ionizes OH – ion in the aqueous solution is called a Base.

HCl (aquaous solution) ⇌ H + + Cl – NaOH (aquaous solution) ⇌ Na + + OH –

Bronsted-Lowry Theory of Acid and Base

Bronsted-Lowry Theory also provides an explanation of acid and base, according to this theory, acid is an H + ion or a proton donor and it forms its conjugate base whereas the base is a substance that accepts an H + ion or a proton to form conjugate acid.

Bronsted-Lowry Acid

According to Bronsted-Lowry acids are substances that donate a proton or H + ion to the other compound.

Acid ⇌ Proton + Conjugate Base

Example: H 2 SO 4 ⇌  H + + HSO 4 –

Bronsted-Lowry Base

According to Bronsted-Lowry bases are substances that accept a proton or H + ion from other compounds.

Base + Proton ⇌ Conjugate Acid

Example: OH – + H + ⇌  H 2 O 

Strength of Acids and Bases

The strength of an acid or a base is measured by the amount of H + ions or OH – ions present in their aqueous solution.

Strong acids have a higher concentration of H + ions per unit volume in their aqueous solution whereas weaker acids have a lower concentration of H + ions per unit volume in their aqueous solution. An example of a strong acid is H 2 SO 4 and an example of a weak acid is CH 3 COOH. 

Similarly, strong bases have a higher concentration of OH – ions per unit volume in their aqueous solution whereas weaker bases have a lower concentration of OH – ions per unit volume in their aqueous solution. An example of a strong base is KOH and an example of a weak acid is CaO.

The strength of Acids and Bases can easily be measured using a pH scale.

pH scale is used to measure the basicity and acidity of a solution. It gives the strength of any solution. pH is determined by the amount of hydrogen ion concentration in the solution.

It is calculated using the formula,

pH = -log[H + ]

For an acid, pH ranges from 0 to 7 whereas for a base it ranges between 7 and 14. The lower the pH higher is the strength of the acid and the higher the pH higher the strength of the base.

Note: pH ranges of acids and bases.

0 < Acid < 7 7 < Base< 14

For more detail on pH read, Importance of pH in Everyday Life .

Indicators are chemical compounds which help to indicate the presence of acid or base in a chemical reaction. They possess different colours in acidic solutions and different colours in basic solutions. Indicators are made naturally by plants and animals or artificially by humans. The image shows a litmus test of acids and bases.

Litmus test of acid and base

An indicator indicating the pH

  • The range of 0 to 7 indicates an acidic solution. 
  • The range of 7 to 14 indicates the basic solution.
  • 7 is a neutral solution.

Types of Indicators

There are various types of indicators used for various purposes some of which are,

  • Natural Indicators: Indicators derived from plants, animals or any living organism are natural indicators. Examples, Red Cabbage, Litmus paper and others.
  • Synthetic Indicators:  Indicators made artificially in laboratories and factories are synthetic indicators. Examples, are Phenopthelien, Methyl orange, and others.
  • Olfactory Indicators: Substances that have different smells in acidic or basic mediums are called Olfactory Indicators. Example onions, olives and others.

When an acid and a base react to neutralize one another, they generate sales, which are ionic substances. Salts do not have an electrical charge. Salts come in a variety of forms, the most common of which is sodium chloride. Table salt and common salt are both terms for sodium chloride. Sodium chloride is used to make dishes taste better. The image given below shows a salt and its cation and anion.

Salts

Physical Properties of Salt

Salts have various physical properties and some of following physical properties of salts are,

  • In nature, the bulk of the salts is crystalline.
  • Salts that are transparent or opaque are available.
  • The bulk of salts is soluble in water.
  • Salt solutions, in their molten state, also transmit electricity.
  • The flavour of salt can be salty, sour, sweet, bitter, or umami (savoury).
  • There is no odour to neutral salts.
  • Salts that are colourless or coloured are available.
  • Because it contains ions, salt water is an excellent conductor of electricity.
  • Electrostatic attraction holds the ions together, and a chemical bond is established between them.

Check, Neutralization Reaction

Salts are categorised into various categories some of the important categories are given below

Acidic Salt

Basic or alkali salt, neutral salts.

A partial neutralisation of a diprotic or polyprotic acid produces an acidic salt. These salts contain H + cations or strong cations in their aqueous solution. The ionizable H + makes up the majority of the ions. Some examples of acidic salts are NaHSO 4 ­, KH 2 PO 4 etc. These salts are formed by the neutralization of strong acids and weak bases.

Ammonium Chloride

Ammonium chloride is formed when hydrochloric acid (a strong acid) interacts with ammonium hydroxide (a weak base).

NH 4 OH + HCl → NH 4 Cl + H 2 O

Ammonium Sulphate

Ammonium sulphate is formed when ammonium hydroxide (a weak base) reacts with sulphuric acid (a strong acid).

2NH 4 OH + H 2 SO 4 → (NH 4 ) 2 SO 4 + 2H 2 O

A basic salt is formed when a strong base reacts with a weak acid to partially neutralise it. When they are hydrolyzed, they decompose into a basic solution. This is because when a basic salt is hydrolyzed, it produces the conjugate base of a weak acid in the solution. e.g. Sodium Carbonate (Na 2 CO 3 ), Sodium Acetate (CH 3 COONa)

Sodium Carbonate

Sodium carbonate is formed when sodium hydroxide (a strong base) reacts with carbonic acid (a weak acid)

H 2 CO 3 + 2NaOH → Na 2 CO 3 + H 2 O

Sodium Acetate

Sodium acetate is formed when a strongly basic, sodium hydroxide (a strong base), reacts with acetic acid (a weak acid)

CH 3 COOH + NaOH → CH 3 COONa + H 2 O

Salts generated by the reaction of a strong acid with a strong base are neutral in nature. The pH of these salts is 7, which is considered neutral. Potassium Chloride, Sodium Chloride, and others are examples of neutral salts.

Sodium Chloride (NaCl)

Sodium Chloride is formed when hydrochloric acid (a strong acid) mixes with sodium hydroxide (a strong base).

NaOH + HCl → NaCl + H 2 O

Sodium Sulphate (Na 2 SO 4 )

It’s made when sulphuric acid combines with sodium hydroxide (a strong basic) ( a strong acid).

2NaOH + H 2 SO 4 → Na 2 SO 4 + 2H 2 O

Salts can also be categorised into other categories which include,

Mixed Salts

Double salt.

Salts with more than one cation or anion are known as double salts. They’re created by mixing two different salts that crystallised in the same ionic lattice. e.g. Potassium Sodium Tartrate (KNaC 4 H 4 O 6 .4H 2 O) also known as Rochelle salt.

Salts which are produced by mixing two salts, which generally share a common cation or anion, are called mixed salt. e..g. Bleaching Powder CaOCl 2 .

What Causes the Formation of Acidic, Basic, and Neutral Salts?

The causes of the formation of Acidic, Basic, and Neutral Salts are discussed below,

When a strong acid reacts with a weak base, the base is unable to completely neutralise the acid. As a result, a salt that is acidic forms. When a strong base is combined with a weak acid, the acid is unable to completely neutralise it. As a result, you get a simple salt. When an equal-strength acid and base react, they totally neutralise each other. A neutral salt is formed as a result of this process.

Check, What is meant by Family of Salts?

Some Common Salts

Salts are chemical compounds that are formed as a result of a neutralization reaction between acids and bases. When we hear salt we only think about common salt which is Sodium chloride that we eat in our daily life but there are several other salts also which are widely useful. Here in this article, we will learn about some common salt that is widely used.

  • Baking Soda or Sodium Bicarbonate
  • Washing Soda or Sodium Carbonate
  • Bleaching Powder or Calcium Hypochlorite

Baking Soda

Baking soda also called Sodium Hydrogen Carbonate, is a chemical compound whose chemical formula is NaHCO 3 . Baking soda has a sodium cation (Na + ) and a bicarbonate anion (HCO 3 – ). Sodium bicarbonate is a white, crystalline powder and as the name suggests is used for baking.

Chemical Name: Sodium hydrogen carbonate

Chemical Formula: NaHCO 3

Preparation:

Baking soda can be prepared with the help of the reaction given below.

NaCl(aq) + NH 3 (g) + CO 2 (g) + H 2 O(l) → NaHCO 3 (aq) + NH 4 Cl(aq)

A few of the uses of Baking Soda are,

  • It is used as an antacid in case of acidity.
  • It is used for baking purposes.
  • It is used as a water softener.

Washing Soda

Washing soda also called Sodium Carbonate, is a chemical compound whose chemical formula is Na 2 CO 3 . Washing soda has two sodium cations (Na + ) and a carbonate anion (CO 3 2- ). Sodium carbonate is a white, crystalline powder and as the name suggests is used for washing purposes.

Chemical Name: Sodium Carbonate

Chemical Formula: Na 2 CO 3

A few of the uses of Washing Soda are,

  • It is used in the glass, soap and paper industries.
  • It is used as washing powder.

Bleaching Powder

Bleaching Powder also called Calcium Hypochlorite, is a chemical compound whose chemical formula is CaOCl 2 . Bleaching Powder is used for bleaching purposes. In its aqueous solution bleaching powder releases chlorine which is responsible for the bleaching action.

Chemical Name: Calcium Hypochlorite

Chemical Formula: CaOCl 2

Bleaching Powder can be prepared with the help of the reaction given below.

Ca(OH) 2 (aq) + Cl 2 (g) → CaOCl 2 (aq) + H 2 O(l)

A few of the uses of Bleaching Powder are,

  • It is used for bleaching the laundry.
  • It is used as an oxidizer in many industries.
  • It is used as a disinfectant to clean water

Crystals of Salts

Some salts combining with water form crystals and these water molecules which are required to form crystals are called water of crystallisation. Some examples of crystal salts are Table salt (sodium chloride crystals), Sugar (sucrose crystals).

Plaster of Paris

Plaster of Paris is a widely used chemical compound is used for various purposes such as sculpting materials, gauze bandages, building and furnishing houses and others. Plaster of Paris is hydrated calcium sulphate obtained by calcining gypsum. It is a white powdery chemical compound.

Chemical Name: Calcium Sulphate Hemi Hydrate.

Chemical Formula: CaSO4. ½ H 2 O

Preparation of Plaster of Paris

Plaster of Paris can easily be prepared with the help of the equation given below,

CaSO 4 .2H 2 O (s) (heating at 100°C ) —> CaSO 4 . ½ H 2 O + 3/2 H 2 O

Also, Check

Ionization of Acids and Bases Acids and bases

FAQs on Acids, Bases, and Salts

What are salts in acids, bases, and, salts.

The neutralization reaction of acids and bases results in a substance called salt. Salts are made of cations and anions. Some examples of salt are NaCl, Na 2 SO 4

What are the two types of acids?

Acids can easily be categorised as, Inorganic Acids: Examples, Sulphuric Acid (H 2 SO 4 ), Nitric acid (HNO 3 ), and others. Organic Acids: Examples, Acetic acid Citric acid, and others.

What is the difference between an acid and a base?

Acids and bases are two types of corrosive chemicals.  Acids are ionic chemicals that break down in water to create the hydrogen ion (H + ) and they have a pH value between 0 and 7, while base are ionic chemicals that break down in water to create the hydronium ion (OH – ) and they have a pH value between 7 and 14.

What are the physical properties of bases?

The physical properties of bases are, They have a bitter taste to them. Their aqueous solutions have a soapy quality to them. They change the colour of litmus paper from red to blue. Their aqueous solutions are electrically conductive. In an aqueous solution, they release OH – ions.

What are the physical properties of acids?

The physical properties of the acids are Acids have a sour flavour to them. Blue litmus turns red. Electricity can be conducted through an acidic solution. In an aqueous solution, they release H + ions.

What happens when hydrochloric acid reacts with sodium carbonate?

When hydrochloric acid combines with sodium carbonate, it produces carbon dioxide gas, sodium chloride, and water.                                               Na 2 CO 3 + 2HCl → 2NaCl + H 2 O + CO 2

Are salt basic or acidic?

A salt can either be basic or be acidic depending on the types of acid and base that react to form a salt.

Is NH 4 Cl a basic salt?

No, Ammonium chloride (NH 4 Cl) is an acidic salt because it is a salt of a strong acid (hydrochloric acid) and a weak base (ammonium hydroxide).

What happens when metal reacts with HCl?

Metal reacting with acid produces salt and hydrogen. Acid + Metal → Salt + Hydrogen

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3.E: Acid–Base (more practice questions with answers)

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16.1: Acids and Bases: A Brief Review

16.2: brønsted–lowry acids and bases, conceptual problems.

  • \(HSO^−_{4}\,(aq)+H_2O\,(l) \rightleftharpoons SO^{2−}_{4}\,(aq)+H_3O^{+}\,(aq)\)
  • \(C_{3}H_{7}NO_{2}\,(aq)+H_{3}O^{+}\,(aq) \rightleftharpoons C_{3}H_{8}NO^{+}_{2}\,(aq)+H_{2}O\,(l)\)
  • \(CH_{3}O_{2}H\,(aq)+NH_{3}\,(aq) \rightleftharpoons CH_{3}CO^{−}_{2}\,(aq)+NH^{+}_{4}\,(aq)\)
  • \(SbF_{5}\,(aq)+2\,HF\,(aq) \rightleftharpoons H_{2}F^{+}\,(aq)+SbF^{−}_{6}\,(aq)\)
  • \(HF\,(aq)+H_{2}O\,(l) \rightleftharpoons H_3O^{+}\,(aq)+F^{−}\,(aq)\)
  • \(CH_3CH_2NH_{2}\,(aq)+H_{2}O\,(l) \rightleftharpoons CH_3CH_2NH^{+}_{3}\,(aq)+OH^{−}\,(aq)\)
  • \(C_3H_7NO_{2}\,(aq)+OH^{−}\,(aq) \rightleftharpoons C_3H_6NO^{−}_{2}\,(aq)+H_{2}O\,(l)\)
  • \(CH_3CO_2H\,(aq)+2\,HF\,(aq) \rightleftharpoons CH_3C(OH)_{2}^{+}\,(aq)+HF^{−}_{2}\,(aq)\)
  • Salts such as NaH contain the hydride ion (\(H^−\)). When sodium hydride is added to water, it produces hydrogen gas in a highly vigorous reaction. Write a balanced chemical equation for this reaction and identify the conjugate acid–base pairs.
  • \(HCO^−_{3}\,(aq)+H_2O\,(l) \rightleftharpoons CO^{2−}_{3}\,(aq)+H_3O^{+}\,(aq)\)
  • \(formic\;acid\,(aq)+H_2O\,(l) \rightleftharpoons formate\,(aq)+H_3O^+\,(aq)\)
  • \(H_3PO_{4}\,(aq)+H_2O\,(l) \rightleftharpoons H_2PO^−_{4}\,(aq)+H_3O^+\,(aq)\)
  • \(OCH^{−}_{3}\,(aq)+H_2O\,(l) \rightleftharpoons HOCH_{3}\,(aq)+OH^{-}\,(aq)\)
  • \(NH^−_{2}\,(aq)+H_2O\,(l) \rightleftharpoons NH_{3}\,(aq)+OH^{−}\,(aq)\)
  • \(S^{2−}\,(aq)+H_2O\,(l) \rightleftharpoons HS^−\,(aq)+OH^−\,(aq)\)
  • \(HBr\,(aq)+H_2O\,(l) \rightleftharpoons H_3O^+\,(aq)+Br^−\,(aq)\)
  • \(NaH\,(s)+NH_{3}\,(aq) \rightleftharpoons H_{2}\,(g)+NaNH_{2}\,(s)\)
  • \(OCH^{−}_{3}\,(aq)+NH_{3}\,(aq) \rightleftharpoons CH_{3}OH\,(aq)+NH^−_{2}\,(aq)\)
  • \(NH_{3}\,(aq)+HCl\,(aq) \rightleftharpoons NH^{+}_{4}\,(aq)+Cl^−\,(aq)\)
  • Species that are strong bases in water, such as \(CH_3^−\), \(NH_2^−\), and \(S^{2−}\), are leveled to the strength of \(OH^−\), the conjugate base of \(H_2O\). Because their relative base strengths are indistinguishable in water, suggest a method for identifying which is the strongest base. How would you distinguish between the strength of the acids \(HIO_3\), \(H_2SO_4\), and \(HClO_4\)?
  • Is it accurate to say that a 2.0 M solution of \(H_2SO_4\), which contains two acidic protons per molecule, is 4.0 M in \(H^+\)? Explain your answer.
  • The alkalinity of soil is defined by the following equation: alkalinity = \([HCO_3^−] + 2[CO_3^{2−}] + [OH^−] − [H^+]\). The source of both \(HCO_3^−\) and \(CO_3^{2−}\) is \(H_2CO_3\). Explain why the basicity of soil is defined in this way.
  • Why are aqueous solutions of salts such as \(CaCl_2\) neutral? Why is an aqueous solution of \(NaNH_2\) basic?
  • \(C_2H_5NH_3Cl\)
  • \((CH_3)_2NH_2^+Br^−\)
  • Which complex ion would you expect to be more acidic: \(Pb(H_2O)_4^{2+}\) or \(Sn(H_2O)_4^{2+}\)? Why?
  • Would you expect \(Sn(H_2O)_4^{2+}\) or \(Sn(H_2O)_6^{4+}\) to be more acidic in aqueous solutions? Why?
  • Is it possible to arrange the hydrides \(LiH\), \(RbH\), \(KH\), \(CsH\), and \(NaH\) in order of increasing base strength in aqueous solution? Why or why not?

Conceptual Answer

a. \(\underset{\text{acid}}{HSO^−_{4}\,(aq)} + \underset{\text{base}}{H_2O\,(l)} \rightleftharpoons \underset{\text{conjugate base}}{SO^{2−}_{4}\,(aq)} + \underset{\text{conjugate acid}}{H_3O^+\,(aq)}\)

b. \(\underset{\text{base}}{C_3H_7NO_{2}\,(aq)} + \underset{\text{acid}}{H_3O^+\,(aq)} \rightleftharpoons \underset{\text{conjugate acid}}{C_3H_8NO^+_{2}\,(aq)} + \underset{\text{conjugate base}}{H_2O\,(l)}\)

c. \(\underset{\text{acid}}{HOAc\,(aq)} + \underset{\text{base}}{NH_{3}\,(aq)} \rightleftharpoons \underset{\text{conjugate base}}{CH_3CO^−_{2}\,(aq)} + \underset{\text{conjugate acid}}{NH^+_{4}\,(aq)}\)

d. \(\underset{\text{acid}}{SbF_{5}\,(aq)} + \underset{\text{base}}{2\,HF\,(aq)} \rightleftharpoons \underset{\text{conjugate acid}}{H_2F^+\,(aq)} + \underset{\text{conjugate base}}{SbF_6^−(aq)}\)

a. \(\underset{\text{acid}}{HF\,(aq)} + \underset{\text{base}}{H_2O\,(l)} \rightleftharpoons \underset{\text{conjugate acid}}{H_{3}O^{+}\,(aq)} + \underset{\text{conjugate base}}{F^{-}\,(aq)}\)

b. \(\underset{\text{base}}{CH_{3}CH_{2}NH_{2}\,(aq)} + \underset{\text{acid}}{H_{2}O\,(l)} \rightleftharpoons \underset{\text{conjugate acid}}{CH_{3}CH_{2}NH_{3}^{+}\,(aq)} + \underset{\text{conjugate base}}{OH^{-}\,(aq)}\)

c. \(\underset{\text{acid}}{C_{3}H_{7}NO_{2}\,(aq)} + \underset{\text{base}}{OH^{-}\,(aq)} \rightleftharpoons \underset{\text{conjugate base}}{C_{3}H_{6}NO_{2}^{-}\,(aq)} + \underset{\text{conjugate acid}}{H_{2}O\,(l)}\)

d. \(\underset{\text{base}}{CH_{3}CO_{2}H\,(aq)} + \underset{\text{acid}}{2\,HF\,(aq)} \rightleftharpoons \underset{\text{conjugate acid}}{CH_{3}C(OH)_{2}^{+}\,(aq)} + \underset{\text{conjugate base}}{HF_{2}^{-}(aq)}\)

3. \(\underset{\text{base}}{NaH\,(s)} + \underset{\text{acid}}{H_{2}O\,(l)} \rightleftharpoons \underset{\text{conjugate acid}}{H_{2}\,(g)} + \underset{\text{conjugate base}}{NaOH\,(aq)}\)

a. \(K_a=\frac{[CO_{3}^{2-}][H_{3}O^{+}]}{[HCO_{3}^{-}]}\)

b. \(K_a=\frac{[formate][H_{3}O^{+}]}{[formic\,acid]}\)

c. \(K_a=\frac{[H_{2}PO_{4}^{-}][H_{3}O^{+}]}{[H_{3}PO_{4}]}\)

a. \(K_b=\frac{[CO_{3}^{2-}][H_{3}O^{+}]}{[HCO_{3}^{-}]}\)

b. \(K_b=\frac{[NH_{3}][OH^{-}]}{[NH_{2}^{-}]}\)

c. \(K_b=\frac{[HS^{-}][OH^{-}]}{[S^{2-}]}\)

6. Strong acids have the smaller \(pK_a\).

a. Equilibrium lies primarily to the right because \(HBr\) (\(pK_a=-8.7\)) is a stronger acid than \(H_{3}O^{+}\) (\(pK_a=-1.7\)) and \(H_{2}O\) (\(pK_a=14\)) is a stronger base than \(Br^-\) ( \(pK_a=-8.7\)).

b. Equilibrium lies primarily to the left because \(H_{2}\) (\(pK_a=36\)) is a stronger acid than \(NH_{3}\) (\(pK_a=38\)) and (\(NaNH_2\)) (\(pK_a=38\)) is a stronger base than \(NaH\) (\(pK_a=35\)).

c. Equilibrium lies primarily to the left because \(CH_{3}OH\) (\(pK_a=17\)) is a stronger acid than \(NH_{3}\) (\(pK_a=38\)) and \(NH_{2}^{-}\) (\(pK_a=38\)) is a stronger base than \(OCH_{3}^{-}\) (\(pK_a=25\)).

d. Equilibrium lies to the right because \(HCl\) (\(pK_a=-7\)) is a stronger acid than \(NH_{4}^{+}\) (\(pK_a=9.3\)) and \(NH_{3}\) is a stronger base than \(Cl^{-}\) (\(pK_a=-7\)).

7. To identify the strongest base we can determine their weakest conjugate acid. The conjugate acids of \(CH_{3}^{-}\), \(NH_{2}^{-}\), and \(S_{2}^{-}\) are \(CH_{4}\), \(NH_{3}\), and \(HS^{-}\), respectively. Next, we consider that acidity increases with positive charge on the molecule, thus ruling out that \(S_{2}^{-}\) is the weakest base. Finally, we consider that acidity increases with electronegativity, therefore \(NH_{3}\) is the second most basic and \(CH_{4}\) is the most basic. To distinguish between the strength of the acids \(HIO_3\), \(H_{2}SO_{4}\), and \(HClO_4\) we can consider that the higher electronegativity and oxidation state of the central nonmetal is the more acidic, therefore the order of acidity is: \(HIO_3\)<\(H_{2}SO_{4}\)<\(HClO_4\) because electronegativity and oxidation state increases as follows: \(I(+5)<S(+6)<Cl(+7)\).

8. It is not accurate to say that a 2.0 M solution of \(H_2SO_4\), which contains two acidic protons per molecule, is 4.0 M in \(H^+\) because a 2.0 M solution of \(H_2SO_4\) is equivalent to 4.0 N in \(H^+\).

\(\frac{2.0\,mol\,H_{2}SO_{4}}{1\,L} \times \frac{2\,eq\,H^{+}}{1\,mol\,H_{2}SO_{4}}=\frac{4\,eq\,H^{+}}{L}=4\,N\,H^{+}\)

9. Alkalinity is a measure of acid neutralizing capability. The basicity of the soil is defined this way because bases such as \(HCO_{3}^{-}\) and \(CO_{3}^{2-}\) can neutralize acids in soil. Because most soil has a pH between 6 and 8, alkalinity can be estimated by its carbonate species alone. At a near neutral pH, most carbonate species are bicarbonate.

10. Aqueous solutions of salts such as \(CaCl_{2}\) are neutral because it is created from hydrochloric acid (a strong acid) and calcium hydroxide (a strong base). An aqueous solution of \(NaNH_2\) is basic because it can deprotonate alkynes, alcohols, and a host of other functional groups with acidic protons such as esters and ketones.

a. \(Li_3N\) is a base because the lone pair on the nitrogen can accept a proton.

b. \(NaH\) is a base because the hydrogen has a negative charge.

c. \(KBr\) is neutral because it is formed from \(HBr\) (a strong acid) and \(KOH\) (a strong base).

d. \(C_2H_5NH_3Cl\) is acidic because it can donate a proton.

a. The pH is expected to increase. \(\underset{\text{acid}}{LiCH_{3}\,(aq)} + \underset{\text{base}}{H_2O\,(l)} \rightleftharpoons \underset{\text{conjugate base}}{LiOH\,(aq)} + \underset{\text{conjugate acid}}{CH_{4}\,(aq)}\)

b. The pH is expected to increase. \(\underset{\text{acid}}{MgCl_{2}\,(aq)} + \underset{\text{base}}{H_2O\,(l)} \rightleftharpoons \underset{\text{conjugate acid}}{2\,HCl\,(aq)} + \underset{\text{conjugate base}}{MgO\,(aq)}\)

c. The pH is expected to remain the same. \(K_{2}O\,(aq)+H_2O\,(l) \rightleftharpoons 2\,KOH\,(aq)\)

d. The pH is expected to increase. \(\underset{\text{acid}}{(CH_3)_2NH_2^+Br^−\,(aq)} + \underset{\text{base}}{H_2O\,(l)} \rightleftharpoons \underset{\text{conjugate acid}}{H_3O^{+}\,(aq)} + \underset{\text{conjugate base}}{(CH_{3})_{2}NH\,(aq)}\)

13. \(Sn(H_2O)_4^{2+}\) is expected to be more acidic than \(Pb(H_2O)_4^{2+}\) because \(Sn\) is more electronegative than \(Pb\).

14. \(Sn(H_2O)_6^{4+}\) is expected to be more acidic than \(Sn(H_2O)_4^{2+}\) because the charge on \(Sn\) is greater (\(4^+>2^+\)).

15. Yes, it is possible the order of increasing base strength is: \(LiH<NaH<RbH<CsH\) because increasing base strength is dependent on decreasing electronegativity.

Numerical Problems

  • acid A: \(pK_a = 1.52\)
  • acid B: \(pK_a = 6.93\)
  • acid C: \(pK_a = 3.86\)

Given solutions with the same initial concentration of each acid, which would have the highest percent ionization?

  • base A: \(pK_b = 13.10\)
  • base B: \(pK_b = 8.74\)
  • base C: \(pK_b = 11.87\)

Given solutions with the same initial concentration of each base, which would have the highest percent ionization?

  • −2.50
  • Benzoic acid is a food preservative with a \(pK_a\) of 4.20. Determine the \(K_b\) and the \(pK_b\) for the benzoate ion.
  • Determine \(K_a\) and \(pK_a\) of boric acid \([B(OH)_3]\), solutions of which are occasionally used as an eyewash; the \(pK_b\) of its conjugate base is 4.80.

Numerical Answers

1. Acids in order of increasing strength: \(acid\,B<acid\,C<acid\,A\). Given the same initial concentration of each acid, the highest percent of ionization is acid A because it is the strongest acid.

2. Bases in order of increasing strength: \(base\,A<base\,C<base\,B\). Given the solutions with the same initial concentration of each base, the higher percent of ionization is base A because it is the weakest base.

\(pK_a+pK_b=14 \rightarrow pK_a=14-pK_b=14-3.80=10.2\)

\(K_a=10^{-pK_a}=10^{-10.2}=6.31 \times 10^{-11}\)

\(pK_a+pK_b=14 \rightarrow pK_a=14-pK_b=14-7.90=6.10\)

\(K_a=10^{-pK_a}=10^{-6.10}=7.94 \times 10^{-7}\)

\(pK_a+pK_b=14 \rightarrow pK_a=14-pK_b=14-7.90=3.000 \times 10^{-1}\)

\(K_a=10^{-pK_a}=10^{-3.000 \times 10^{-1}}=-5.012 \times 10^{-1}\)

\(pK_a+pK_b=14 \rightarrow pK_a=14-pK_b=14-1.40=12.6\)

\(K_a=10^{-pK_a}=10^{-12.6}=2.51 \times 10^{-13}\)

e. \(pK_a+pK_b=14 \rightarrow pK_a=14-pK_b=14-7.90=16.5\)

\(K_a=10^{-pK_a}=10^{-16.5}=3.16 \times 10^{-17}\)

\(pK_a+pK_b=14 \rightarrow pK_b=14-pK_a=14-4.20=9.80\)

\(K_b=10^{-pK_b}=10^{-9.80}=1.58 \times 10^{-10}\)

\(pK_a+pK_b=14 \rightarrow pK_a=14-pK_b=14-4.80=9.20\)

\(K_a=10^{-pK_a}=10^{-9.20}=6.31 \times 10^{-10}\)

16.3: The Autoionization of Water

  • What is the relationship between the value of the equilibrium constant for the autoionization of liquid water and the tabulated value of the ion-product constant of liquid water (\(K_w\))?
  • The density of liquid water decreases as the temperature increases from 25°C to 50°C. Will this effect cause \(K_w\) to increase or decrease? Why?
  • Show that water is amphiprotic by writing balanced chemical equations for the reactions of water with \(HNO_3\) and \(NH_3\). In which reaction does water act as the acid? In which does it act as the base?
  • Nitric acid is added to water.
  • Potassium hydroxide is added to water.
  • Calcium hydroxide is added to water.
  • Sulfuric acid is added to water.
  • Show that \(K\) for the sum of the following reactions is equal to \(K_w\).

\[HMnO_{4}\,(aq) \rightleftharpoons H^+\,(aq) + MnO^−_{4}\,(aq)\]

\[MnO^−_{4}\,(aq)+H_2O\,(l) \rightarrow HMnO_{4}\,(aq) + OH^−\,(aq)\]

Conceptual Answers

\[K_{auto} = \dfrac{[H_3O^+][OH^−]}{[H_2O]^2}\]

\[K_w = [H_3O^+][OH^−] = K_{auto}[H_2O]^2\]

2. This will affect \(K_w\) as it is dependent on temperature. As the temperature increases, an endothermic process occurs (energy must be absorbed to break the bonds). Consequently, according to Le Chatelier, an increase in temperature favors the forward reaction thus the position of equilibrium shifts toward the right-hand side and \(K_w\) becomes larger.

a. \(HNO_3\,(aq)+H_2O\,(l) \rightleftharpoons H_3O^{+}\,(aq)+ HNO_{2}^{-}\,(aq)\)

b. \(KOH\,(s)+H_2O\,(l) \rightleftharpoons K^{-}\,(aq)+OH^{-}\,(aq)\)

c. \(Ca(OH)_{2}\,(s)+H_2O\,(l) \rightleftharpoons Ca^{2+}\,(aq)+2\,OH^{-}\,(aq)\)

d. \(H_2SO_4\, (aq)+H_2O\,(l) \rightleftharpoons HSO_4^{-}\,(aq)+H^{+}\,(aq)\)

\(H_{2}O\,(l) \rightleftharpoons H^{+}\,(aq)+OH^{-}\,(aq)\)

\(K_w=[H^{+}][OH^{-}]\)

  • The autoionization of sulfuric acid can be described by the following chemical equation: \[H_2SO_{4}\,(l)+H_2SO_{4}\,(aq) \rightleftharpoons H_3SO^+_{4}\,(aq)+HSO_{4}^{-}\,(aq)\] At 25°C, \(K = 3 \times 10^{−4}\). Write an equilibrium constant expression for \(K_{H_2SO_4}\) that is analogous to \(K_w\). The density of \(H_2SO_4\) is \(1.8\frac{g}{cm^{3}}\) at 25°C. What is the concentration of \(H_3SO_{4}^{+}\) ? What fraction of \(H_2SO_4\) is ionized?
  • An aqueous solution of a substance is found to have \([H_3O]^+ = 2.48 \times 10^{−8}\; M\). Is the solution acidic, neutral, or basic?
  • The pH of a solution is 5.63. What is its pOH? What is the \([OH^{−}]\)? Is the solution acidic or basic?
  • \([H_3O^+] = 8.6 \times 10^{−3}\; M\)
  • \([H_3O^+] = 3.7 \times 10^{−9}\; M\)
  • \([H_3O^+] = 2.1 \times 10^{−7}\; M\)
  • \([H_3O^+] = 1.4 \times 10^{−6}\; M\)
  • 0.15 \(M\,HBr\)
  • 0.03 \(M\,KOH\)
  • \(2.3 \times 10^{−3}\; M\; HNO_3\)
  • \(9.78 \times 10^{−2} \;M\; NaOH\)
  • 0.00017 \(M\,HCl\)
  • 5.78 \(M\,HI\)
  • 25.0 mL of \(2.3 \times 10^{−2}\;M\;HCl\), diluted to 100 mL
  • 5.0 mL of \(1.87\,M\,NaOH\), diluted to 125 mL
  • 5.0 mL of \(5.98\,M\,HCl\) added to 100 mL of water
  • 25.0 mL of \(3.7\,M\,HNO_3\) added to 250 mL of water
  • 35.0 mL of \(0.046\,M\,HI\) added to 500 mL of water
  • 15.0 mL of \(0.0087\,M\,KOH\) added to 250 mL of water.
  • The pH of stomach acid is approximately 1.5. What is the \([H^+]\)?
  • household bleach (11.4)
  • orange juice (3.5)
  • seawater (8.5)
  • tomato juice (4.2)
  • A reaction requires the addition of 250.0 mL of a solution with a pH of 3.50. What mass of HCl (in milligrams) must be dissolved in 250 mL of water to produce a solution with this pH?
  • If you require 333 mL of a pH 12.50 solution, how would you prepare it using a 0.500 M sodium hydroxide stock solution?

\[K_{H_2SO_4}=[H_3SO_4^+][HSO_4^−]=K[H_2SO_4]_2\]

\[[H_3SO_4^+] = 0.3\,M\]

So the fraction ionized is 0.02.

2. The solution is basic because the \(pH=-log([H_3O^{+}])=-log(2.48 \times 10^{−8})=7.61>7\).

\(pH+pOH=14 \rightarrow pOH=14-pH=14-5.63=8.37\)

\([OH^{-}]=10^{-pOH}=-4.27 \times 10^{-9}\)

The \(pH=5.63<7\), therefore the solution is acidic.

a. The solution is acidic. \(pH=-log([H_3O^{+}])=-log(8.6 \times 10^{−3})=2.1<7\)

b. The solution is basic. \(pH=-log([H_3O^{+}])=-log(3.7 \times 10^{−9})=8.4>7\)

c. The solution is acidic. \(pH=-log([H_3O^{+}])=-log(2.1 \times 10^{−7})=6.7<7\)

d. The solution is acidic. \(pH=-log([H_3O^{+}])=-log(1.4 \times 10^{−6})=5.9<7\)

\(pH=-log([H_3O^{+}])=-log(0.15)=0.82\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-0.82=13\)

\(pOH=-log([OH^{-}])=-log(0.03)=2\)

\(pH+pOH=14 \rightarrow pH=14-pOH=14-2=10\)

\(pH=-log([H_3O^{+}])=-log(2.3 \times 10^{−3})=2.6\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-2.6=11\)

\(pOH=-log([OH^{-}])=-log(9.78 \times 10^{−2})=1.01\)

\(pH+pOH=14 \rightarrow pH=14-pOH=14-1.01=13.0\)

\(pH=-log([H_3O^{+}])=-log(0.00017)=3.8\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-3.8=10\)

\(pH=-log([H_3O^{+}])=-log(5.78)=-0.762\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-(-0.762)=14.8\)

a. \(25.0\,mL \times \frac{1\,L}{1,000\,mL} \times \frac{2.3 \times 10^{-2}\,mol}{1\,L} \times \frac{1}{100\,mL \times \frac{1\,L}{1,000\,mL}}=0.060\,M\,HCl\)

\(pH=-log([H_3O^{+}])=-log(0.060)=1.22\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-1.22=12.78\)

b. \(5.0\,mL \times \frac{1\,L}{1,000\,mL} \times \frac{1.87\,mol}{1\,L} \times \frac{1}{125\,mL \times \frac{1\,L}{1,000\,mL}}=7.5 \times 10^{-2}\,M\,NaOH\)

\(pOH=-log([OH^{-}])=-log(7.5 \times 10^{-2})=1.1\)

\(pH+pOH=14 \rightarrow pH=14-pOH=14-1.1=12.9\)

c. \(5.0\,mL \times \frac{1\,L}{1,000\,mL} \times \frac{5.98\,mol}{1\,L} \times \frac{1}{100\,mL \times \frac{1\,L}{1,000\,mL}}=0.20\,M\,HCl\)

\(pH=-log([H_3O^{+}])=-log(0.20)=0.70\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-0.70=13.3\)

d. \(25.0\,mL \times \frac{1\,L}{1,000\,mL} \times \frac{3.7\,mol}{1\,L} \times \frac{1}{250\,mL \times \frac{1\,L}{1,000\,mL}}=0.370\,M\,HNO_3\)

\(pH=-log([H_3O^{+}])=-log(0.370)=0.432\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-0.432=13.568\)

e. \(35.0\,mL \times \frac{1\,L}{1,000\,mL} \times \frac{0.046\,mol}{1\,L} \times \frac{1}{500\,mL \times \frac{1\,L}{1,000\,mL}}=3 \times 10^{-3}\,M\,HI\)

\(pH=-log([H_3O^{+}])=-log(3 \times 10^{-3})=2.52\)

\(pH+pOH=14 \rightarrow pOH=14-pH=14-2.52=11.48\)

f. \(15.0\,mL \times \frac{1\,L}{1,000\,mL} \times \frac{0.0087\,mol}{1\,L} \times \frac{1}{125\,mL \times \frac{1\,L}{1,000\,mL}}=5.20 \times 10^{-4}\,M\,KOH\)

\(pOH=-log([OH^{-}])=-log(5.20 \times 10^{-4})=3.28\)

\(pH+pOH=14 \rightarrow pH=14-pOH=14-3.28=10.72\)

7. \([H^+]=10^{-pH}=10^{-1.5}=3.2 \times 10^{-2}\,M\)

a. \([H^{+}]=10^{-11.4}=3.98 \times 10^{-12}\,M\)

b. \([H^{+}]=10^{-6.5}=3.2 \times 10^{-7}\,M\)

c. \([H^{+}]=10^{-3.5}=3.2 \times 10^{-4}\,M\)

d. \([H^{+}]=10^{-8.5}=3.2 \times 10^{-9}\,M\)

e. \([H^{+}]=10^{-4.2}=6.3 \times 10^{-5}\,M\)

9. 2.9 mg \(HCl\)

\([H^{+}]=10^{-pH}=10^{-3.50}=3.1622 \times 10^{-4}\,M\)

\(x\,mg\,HCl \times \frac{1\,g\,HCl}{1,000\,mg\,HCl} \times \frac{1\,mol\,HCl}{36.46\,g\,HCl} \times \frac{1}{250\,mL\,H_2O \times \frac{1\,L}{1,000\,mL\,H_2O}}=3.1622 \times 10^{-4}\,M\,HCl \rightarrow \frac{x\,mol\,HCl}{9115\,L\,HCl}=3.1622 \times 10^{-4}\,M\,HCl \rightarrow x\,mol\,HCl=3.1622\times 10^{-4}\,M \times 9115\,L\,HCl \rightarrow x\,mol\,HCl=2.9\,mol\,HCl \rightarrow x=2.9\)

10. To prepare the stock solution, \(2.11 \times 10^{-2}\,L\) of \(0.500\,M\,NaOH\) solution is required.

\(\frac{0.03162\,mol\,NaOH}{1\,L\,NaOH} \times \frac{1\,L\,NaOH}{0.5\,mol\,NaOH} \times 0.333\,L=2.11 \times 10^{-2}\,g\,NaOH\).

\([OH^{-}]=10^{-1.5}=0.03162\,M\)

\(pH+pOH=14 \rightarrow pOH=14-12.50=1.5\)

16.4: The pH Scale

16.5: strong acids and bases, 16.6: weak acids, 16.7: weak bases, 16.8: relationship between ka and kb, 16.9: acid-base properties of salt solutions, 16.10: acid-base behavior and chemical structure.

  • \(CH_3CCl_2CH_2CO_2H\) versus \(CH_3CH_2CH_2CO_2H\)
  • \(CH_3CO_2H\) versus \(CH_3CH_2OH\)
  • \(HClO\) versus \(HBrO\)
  • \(\ce{CH_3C(=O)NH_2}\) versus \(CH_3CH_2NH_2\)
  • \(H_3AsO_4\) versus \(H_3AsO_3\)
  • The stability of the conjugate base is an important factor in determining the strength of an acid. Which would you expect to be the stronger acid in aqueous solution—\(C_6H_5NH_3^+\) or \(NH_4^+\)? Justify your reasoning.
  • Explain why \(H_2Se\) is a weaker acid than \(HBr\).
  • Arrange the following in order of decreasing acid strength in aqueous solution: \(H_3PO_4\), \(CH_3PO_3H_2\), and \(HClO_3\).
  • Arrange the following in order of increasing base strength in aqueous solution: \(\ce{CH_3S−}\), \(OH^−\), and \(CF_3S^−\).
  • Arrange the following in order of increasing acid strength in aqueous solution: \(HClO_2\), \(HNO_2\), and \(HNO_3\).
  • Do you expect \(H_2SO_3\) or \(H_2SeO_3\) to be the stronger acid? Why?
  • Give a plausible explanation for why \(CF_3OH\) is a stronger acid than \(CH_3OH\) in aqueous solution. Do you expect \(CHCl_2CH_2OH\) to be a stronger or a weaker acid than \(CH_3OH\)? Why?
  • Do you expect \(Cl_2NH\) or \(NH_3\) to be the stronger base in aqueous solution? Why?

a. The most important factor in determining the stronger acid is considering the inductive effect. Chlorine is an electron-withdrawing group. It pulls electron density away from the compound by means of the inductive effect through the sigma bond. In considering the conjugate base of \(CH_3CCl_2CH_2CO_2H\), Chlorine absorbs some of the electron density or excess negative charge on the oxygen atom. This causes the C bonded to the attached Chlorine atoms to be partially positive. The conjugate base of \(CH_3CCl_2CH_2CO_2H\) is more stable, thus more acidic than the conjugate base of \(CH_3CH_2CH_2CO_2H\).

b. The most important factor in determining the stronger acid is knowing the \(pK_a\) values for functional groups. The \(pK_a\) of alcohol is about 16 while the \(pK_a\) of a carboxylic acid is about 5. Therefore, \(CH_3CO_2H\) is more acidic than \(CH_3CH_2OH\).

c. The most important factor in determining the stronger acid is electronegativity. The chlorine atom is more electronegative than the bromine atom, therefore \(HClO\) is more acidic than \(HBrO\).

d. The most important factor in determining the stronger acid is considering resonance. The \(\ce{CH_3C(=O)NH_2}\) has a resonance which increases the stability of the conjugate base (therefore increasing acidity) because the negative charge can be delocalized. Thus, \(\ce{CH_3C(=O)NH_2}\) is more acidic than \(CH_3CH_2NH_2\).

e. The most important factor in determining the stronger acid is considering oxidation states on the central nonmetal. \(H_3AsO_4\) has an oxidation state of +5 which is larger and thus more acidic than \(H_3AsO_3\) which has an oxidation state of +3.

\(CF_3S^− < CH_3S^− < OH^−\) (strongest base)

\(NH_3\); \(Cl\) atoms withdraw electron density from \(N\) in \(Cl_2NH\).

2. It is expected that the stronger acid is \(C_6H_5NH_3^+\) because in considering the conjugate base the lone pair of electrons on nitrogen is involved in resonance, hence the molecule is stable.

3. \(H_2Se\) is a weaker acid than HBr because \(Br\) is more electronegative than \(Se\) thus more stable.

4. \(HClO_3>CH_3PO_3H_2>H_3PO_4\)

This is because \(H_3PO_4\) is a polyprotic acid which contains more than one ionizable proton, and the protons are lost in a stepwise manner. The fully protonated species is always the strongest acid because it is easier to remove a proton from a neutral molecule than from a negatively charged ion. Thus, acid strength decreases with the loss of subsequent protons, and, correspondingly, the \(pK_a\) increases which indicate it is the most basic. The conjugate base of \(ClO^{3-}\) has a much smaller charge to volume ratio, thus most stable and acidic.

5. \(CF_{3}S^{-}<CH_{3}S^{-}<OH^{-}\)

\(CF_{3}S^{-}\) is the most acidic because of the three electronegative Fluorines. \(OH^{-}\) is the strongest base in water. Thus, \(CH_{3}S^{-}\) is in between these aqueous solutions.

6. \(HNO_2<HClO_2<HNO_3\)

\(HNO_3\) has resonance stabilization, therefore it is the most acidic. Between \(HClO_2\) and \(HNO_2\), \(Cl\) is the most electronegative therefore \(HClO_2\) is more acidic than \(HNO_2\).

7. I expect \(H_2SO_3\) to be the stronger acid because it is more electronegative or has greater attraction which means it’ll be less inclined to share their electrons with a proton.

8. \(CF_3OH\) is a stronger acid than \(CH_3OH\) in aqueous solution because \(F\) is more electronegative than \(H\).

9. It would be expected that \(NH_3\) be a stronger base than \(Cl_2NH\) because it is electronegative due to the two \(Cl\) atoms.

16.11: Lewis Acids and Bases

1. Identify the nature of each of the following as either a Lewis Acid or a Lewis Base:

  • \(Ni^{2+}\)
  • \(Pt^{4+}\)

2. Explain why \(SiF_4\) can act as a Lewis Acid.

3. Identify the nature of each of the following as either a Lewis Acid or a Lewis Base:

  • \([Fe(CN)_6]^{3-}\)
  • \([Ni(NH_3)_6]^{2+}\)
  • \(CdBr_4^{2-}\)

4. What is the product of the reaction of \(CO_2 + OH^- \rightarrow \) ?

5. In the reactions below, which is the Lewis Acid and/or which is the Lewis Base?

  • \(NH_3 + H^+ \rightarrow NH_4^+\)
  • \(H_2O + H^+ \rightarrow H_3O^+\)

6. In the complex ion, \([PtCl_6]^{2-}\) w hich is the Lewis Acid and which is the Lewis base?

7. The reaction of \(AgCl\) + \(NH_3\) pr oduces what complex ion?

a. \(NH_3\) is a lewis base because nitrogen has a lone pair of electrons to "donate."

b. \(Ag^{+}\) is a lewis acid because it has an unfilled octet and thus is able to accept a pair of electrons.

c. \(Ni^{+}\) is a lewis acid because it has an unfilled octet and thus is able to accept a pair of electrons.

d. \(Pt^{4+}\) is a lewis acid because it has an unfilled octet and thus is able to accept a pair of electrons.

e. \(H_2O\) is a lewis base because oxygen has two lone pairs of electrons to "donate."

f. \(SO_2\) is a lewis acid because sulfur has an unfilled octet and thus is able to accept a pair of electrons.

2. \( SiF _4\) has a central Silicon Atom which can expand its octet to 12 (compared to the typical 8) so that it forms \([ SiF _6]^{2-}\).

  • Lewis Acid: \(Fe^{3+}\), Lewis Base: \(CN^-\)
  • Lewis Acid: \(Ni^{2+}\), Lewis Base: \(NH_3\)
  • Lewis Acid: \(Cd^{2+}\), Lewis Base: \(Br^-\)

4. This reaction forms a bicarbonate ion. \(CO_2 + OH^- \rightarrow O--COH=O\).

a. Lewis Acid: \(H^+\), Lewis Base: \(NH_3\)

b. Lewis Acid: \(H^+\), Lewis Base: \(H_2O\)

6. \(Pt^{4+}\) is the Lewis acid and \(Cl^-\) is the Lewis base.

7. \[AgCl + 2\,NH_3 \rightarrow [Ag(NH_3)_2]^+ + Cl^-\]

Construct a table comparing how OH − , NH 3 , H 2 O, and BCl 3 are classified according to the Arrhenius, the Brønsted–Lowry, and the Lewis definitions of acids and bases

Describe how the proton \((H^{+})\) can simultaneously behave as an Arrhenius acid, a Brønsted–Lowry acid, and a Lewis acid.

Would you expect aluminum to form compounds with covalent bonds or coordinate covalent bonds? Explain your answer.

Classify each compound as a Lewis acid or a Lewis base and justify your choice.

a. \(AlCl_{3}\)

b. \(CH_{3}N\)

c. \(IO_{3}^{-}\)

Explain how a carboxylate ion \((RCO_{2}^{-})\) can act as both a Brønsted–Lowry base and a Lewis base.

An Arrhenius acid is a molecule that when dissolved in water it will donate an \(H^{+}\) in solution.

An Arrhenius base is a molecule that when dissolved in water it will donate an \(OH^{-}\) in solution.

A Brønsted–Lowry acid is a molecule that when dissolved in a solution it will donate an \(H^{+}\) in solution.

A Brønsted–Lowry base is a molecule that when dissolved in a solution it will donate an atom or ion capable of accepting or bonding to a free proton in solution.

A Lewis acid is an atom or molecule that accepts an electron pair.

A Lewis base is an atom or molecule that donates an electron pair.

The proton \((H^{+})\) can simultaneously behave as an Arrhenius acid because when it is dissolved in water it will donate itself. \(H^{+}+H_{2}O \rightleftharpoons H_{3}O^{+}\)

The proton \((H^{+})\) can simultaneously behave as a Brønsted–Lowry acid because when it is dissolved in solution it will donate itself.

\(H^{+}+B^{-} \rightleftharpoons HB\)

The proton \((H^{+})\) can simultaneously behave as a Lewis acid as it can accept an electron pair.

3. It is expected that Aluminum forms a coordinate covalent bond as it can participate in a Lewis acid and a Lewis base interaction. For example \(Al^{3+}+H_{2}O \rightleftharpoons [Al(OH_2)_{6}]^{3+}\)

a. \(AlCl_{3}\) is a Lewis acid as \(Al\) can accept an electron pair.

b. \(CH_{3}N\) is a Lewis base as \(N\) can donate an electron pair.

c. \(IO_{3}^{-}\) is a Lewis base as \(I\) can donate an electron pair.

5. The carboxylate ion \((RCO_{2}^{-})\) can act as Brønsted–Lowry base because when dissolved in a solution the electron rich \(O\) is capable of accepting a proton. The carboxylate ion \((RCO_{2}^{-})\) can act as a Lewis base because the electron rich \(O\) can donate an electron pair.

1. In each reaction, identify the Lewis acid and the Lewis base and complete the reaction by writing the products(s).

a. (CH 3 ) 2 O + AlCl 3

b. SnCl 4 + 2 Cl −

2. Use Lewis dot symbols to depict the reaction of BCl 3 with dimethyl ether [(CH 3 ) 2 O]. How is this reaction similar to that in which a proton is added to ammonia?

a. The Lewis acid is \(AlCl_{3}\) and the Lewis base is \((CH_{3})_{2}O\).

\((CH_{3})_{2}O+AlCl_{3} \rightleftharpoons AlCl_{3} \cdot O(CH_{3})_{2}\)

b. The Lewis acid is \(SnCl_{4}\) and the Lewis base is \(Cl^{-}\).

\(SnCl_{4}+2\,Cl^{-} \rightleftharpoons SnCl_{6}^{2-}\)

\(BCl_{3}+(CH_{3})_{2}O \rightleftharpoons BCl_{3}\cdot O(CH_{3})_{2}\)

This reaction is similar to that in which a proton is added to ammonia as it also involves a Lewis acid and a Lewis base interaction.

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Chemistry Worksheets Class 7 on Chapter 5 Acids, Bases and Salts with Answers - Set 1

Substances that we use daily can be classified as acids, bases, and neutral substances. For example, lemon juice is acidic, and baking soda solution is basic and common salt is neutral. Lemon juice is acidic since lemon contains citric acid. Similarly, baking soda contains sodium bicarbonate which acts as a base. Sodium chloride is a salt formed by the neutralisation reaction between an acid and a base.

Download Class 7 Chemistry Worksheet on Chapter 5 Acids, Bases and Salts Set 1 PDF

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Chemistry Worksheets Class 7 on Chapter 5 Acids, Bases and Salts with Answers - Set 1

CBSE Class 7 Chemistry Chapter 5 Acids, Bases and Salts Worksheet – Set 1

Q1. What is the common name of sodium bicarbonate?

(a) Antacid

(b) Baking soda

(c) Common salt

(d) Washing soda

Q2. Baking soda turns ______.

(a) Blue litmus red

(b) Red litmus blue

(c) Phenolphthalein colourless

(d) None of the above

Q3. A substance that turns red litmus blue is a/an ______.

(c) Neutral

Q4. A base ______.

(a) Has a bitter taste

(b) Turns blue litmus red

(c) Has a sour taste

(d) Does not affect turmeric

Q5. Salt is formed when ______.

(a) Acid reacts with water

(b) Base reacts with water

(c) Base reacts with acid

(d) All of the above

Q6. Sodium hydroxide is a ______.

Q7. _______ and _______ are used as indicators.

Q8. Grapes contain _______ acid, while apple has _______ acid.

Q9. State true or false.

Excessive use of fertilisers in the soil makes the soil basic.

Q10. State true or false.

Phenolphthalein and methyl orange are examples of indicators.

Q11. Why are antacids used to cure acidity?

Q12. Name any two antacids.

Q13. Why do we use calamine solution on ant bites?

Q14. Why is factory waste neutralised before disposing of?

Q15. What is the chemical name of milk of magnesia?

Q16. What is an indicator?

Q17. What is a universal indicator?

Q18. What is a pH scale?

Q19. Differentiate between acid and base.

Q20. What is a neutralisation reaction? Explain it with the help of an example.

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    A buffer resists sudden changes in pH. It has a weak acid or base and a salt of that weak acid or base. No. Combining a strong acid and a strong base will produce salt and water. Excess strong acid or strong base will not act as a buffer. not a buffer; buffer; not a buffer; buffer; 4. 1. not a buffer. 2. buffer. 3. not a buffer. 4. not buffer

  2. 17.2: Acids, Bases, and Salts

    This theory was developed by Svante Arrhenius in 1883. Later, two more sophisticated and general theories were proposed. These are the Brønsted-Lowry and the Lewis definitions of acids and bases. The Lewis theory is discussed elsewhere. Figure 17.2.1 17.2. 1 Acid base reaction theories as superset and subset models.

  3. Acids, Bases, and Salts

    Definitions. Acid:- An acid is defined as a substance whose water solution tastes sour, turns blue litmus red, and neutralizes bases. Base:- A substance is called base if its aqueous solution tastes bitter, turns red litmus blue, or neutralizes acids. Salt:- Salt is a neutral substance whose aqueous solution does not affect litmus.

  4. Acids, Bases and Salts Class 7 Science Notes

    Salts. Salt is the product formed from the neutralisation reaction of acids and bases. In the reaction between hydrochloric acid and sodium hydroxide, the salt formed is sodium chloride. HCl+N aOH→N aCl+H2O. Salt can be acidic, basic or neutral in nature. To know more about Salts, visit here.

  5. Acid-Base Chemistry

    The three main theories of acids and bases are the Arrhenius theory, Brønsted-Lowry theory, and Lewis theory. Each of these theories has its uses in chemistry. Arrhenius Theory. Acids: Substances that increase the concentration of H⁺ ions in water.; Bases: Substances that increase the concentration of OH⁻ ions in water.; Brønsted-Lowry Theory. Acids: Proton donors.

  6. Acids, Bases and Salts (practice)

    Salt + Water. Loading... Learn for free about math, art, computer programming, economics, physics, chemistry, biology, medicine, finance, history, and more. Khan Academy is a nonprofit with the mission of providing a free, world-class education for anyone, anywhere.

  7. 10.1: Introduction to Acids and Bases

    Arrhenius Definition. "an acidic substance is one whose molecular unit contains at least one hydrogen atom that can dissociate, or ionize, when dissolved in water, producing a hydrated hydrogen ion and an anion." hydrochloric acid. HCl → H +(aq) + Cl -(aq) sulfuric acid. H 2 SO 4 → H +(aq) + HSO 4-(aq)

  8. 26. Acids & Bases

    Resources. Lecture Slides (PDF - 2.2MB) Lecture Summary. This lecture introduces the chemical models and behaviors of acids and bases. Starting from the historical origins ("acid" derives from the Latin acidus, meaning "sour"), Prof. Sadoway discusses the evolving acid-base models of Lavoisier (1776), Arrhenius (1887), Brønsted and Lowry (1923), and Lewis (1923-1938).

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    What students need to know. Acids are hydrogen-containing substances with a sour taste that form solutions with pH values less than 7. Common examples include hydrochloric acid, sulfuric acid, citric acid and ethanoic acid (vinegar/acetic acid). Bases are a group of substances that neutralise acids. Soluble bases are called alkalis.

  10. Class 10 Chemistry Worksheet on Chapter 2 Acids, Bases and Salts

    Chemistry Worksheets Class 10 on Chapter 2 Acids, Bases and Salts with Answers - Set 1. Acids are substances that taste sour and are corrosive in nature. It turns blue litmus paper to red. These substances are chemically acidic in nature.E.g.:-orange juice, curd, vinegar, hydrochloric acid etc. Bases are substances that, in an aqueous solution ...

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    Acids, Bases, and Salts REACTIONS 19.1 Acid-Base Theories Essential Understanding Acids and bases can be classified in terms of hydrogen ions or ... salt basic They are extremely caustic. electronegative water c a b 0132525887_CHEM_WKBK_CH 19.indd 287 4/7/10 6:58:04 PM. Name Class Date

  12. Introduction to Acids and Bases (Worksheet ...

    The acid-base concept is a system of classifying chemical substances which permits both the organization as well as the prediction of a vast number of chemical reactions. A substance may be assigned to one our four conceivable categories. It may be an acid or a base, but in addition, it may be both an acid and a base or it may be neither an ...

  13. Acid Base and Salt Worksheet Chapter 4

    Question 13 You are provided with four test tubes containing sugar solution, baking soda solution, tamarind solution, salt solution. Write down an activity to find the nature (acidic/basic/neutral) of each solution. Also Read. Notes. Acid base And Salt Class 7 Notes; Assignments. Acid base And Salt NCERT Solutions Class 7; Acid base And Salt ...

  14. PDF Acids, Bases and Salts W

    The reaction between an acid and a base is known as neutralisation. Salt and water are produced in this process with the evolution of heat. Acid+Base Salt+Water (Heat is evolved) The following reaction is an example: Hydrochloric acid (HCl) + Sodium hydroxide (NaOH) Sodium chloride (NaCl) + Water (H 2 O) Boojho added dilute sulphuric acid to ...

  15. Properties of Acids and Bases Assignment and Bases

    Hydrogen gas is highly flammable. Identify two acids and two bases that you use or come into contact with in an average week. Identify uses for each substance. I use citric acid everyday when I eat foods with preservatives and citrus fruits. I use magnesium hydroxide everyday when I put on my deodorant.

  16. Acid-base properties of salts (video)

    The extend of salt hydrolysis depends on the strength of the conjugate base(A-) and conjugate acid (BH⁺) respectively. Kb (conj base) = Kw / Ka (acid) and Ka (conj acid) = Kw / Kb (base) If Kb (conj base) > Ka (conj acid), then the solution would be basic. One line answer, if Ka of the weak acid is more than the Kb of the weak base, the ...

  17. 7.8: Acid-Base Properties of Salts

    7.8: Acid-Base Properties of Salts. Page ID. Salts, when placed in water, will often react with the water to produce H 3 O + or OH -. This is known as a hydrolysis reaction. Based on how strong the ion acts as an acid or base, it will produce varying pH levels. When water and salts react, there are many possibilities due to the varying ...

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    Acids are substances that form hydrogen ions, when dissolved in water. Bases are oxides and hydroxides of metals that react and neutralize acids to form salts and water only. Bases can be strong or weak depending on the extent to which they dissociate and produce OH- ions in solution. Here also briefly describe on common strong and weak acids ...

  19. 7.1A: Acid-Base Theories and Concepts

    There are three primary theories of acid-base chemistry that are often taught together: Arrhenius theory, Brønsted-Lowry theory, and Lewis acid-base theory. Each theory is introduced below. Figure 7.1A. 1 7.1 A. 1: Hierarchal definitions of acids and bases via the three primary theories. These theories are designed to be "superset" of the ...

  20. Acids, Bases, and Salts

    Metal carbonate + Acid → Salt + Carbon dioxide + Water. Example: When hydrochloric acid combines with sodium carbonate, it produces carbon dioxide gas, sodium chloride, and water. Na 2 CO 3 + 2HCl → 2NaCl + H 2 O + CO 2. Reaction of acid with hydrogen carbonates (bicarbonates): When acids react with metal hydrogen carbonates, they produce carbon dioxide gas, salt, and water.

  21. 3.E: Acid-Base (more practice questions with answers)

    2. BCl3 + (CH3)2O ⇌ BCl3 ⋅ O(CH3)2. This reaction is similar to that in which a proton is added to ammonia as it also involves a Lewis acid and a Lewis base interaction. 3.E: Acid-Base (more practice questions with answers) is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

  22. Class 7 Chemistry Worksheet on Chapter 5 Acids, Bases and Salts

    Chemistry Worksheets Class 7 on Chapter 5 Acids, Bases and Salts with Answers - Set 1. Substances that we use daily can be classified as acids, bases, and neutral substances. For example, lemon juice is acidic, and baking soda solution is basic and common salt is neutral. Lemon juice is acidic since lemon contains citric acid.