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Introduction to acid-base theory

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Intros
Lessons
  1. What are acids and bases?
  2. Arrhenius theory.
  3. Lewis and Brønsted-Lowry acids and bases.
  4. What happens to H+ ions?
  5. What is a basic solution?
  6. Measuring pH using H3O+ ions.
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Examples
Lessons
  1. Recall how pH is measured in solutions.
    1. A solution has an H3O+ concentration of 0.07 mol dm-3. What is its pH?
    2. What is the pH of the solution if the concentration is doubled?
    3. How many times more acidic is a solution at pH 2 than a solution at pH 5?
    1. Understand water's properties to act as an acid and a base.
      Water (H2O) is an amphiprotic molecule.
      1. Explain what this means in terms of acids and bases.
      2. Write two chemical equations to show this property.
      Topic Notes
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      Introduction to Acid-Base Theory

      Acid-base theory is a fundamental concept in chemistry that explains the behavior of acids and bases in various chemical reactions. The introduction video provides an essential foundation for understanding this crucial topic. In this article, we'll explore the key concepts of acid-base theory, including the definitions of acids and bases, their properties, and how they interact with each other. We'll delve into the different models of acid-base theory, such as the Arrhenius, Brønsted-Lowry, and Lewis theories, each offering unique perspectives on acid-base interactions. You'll learn about pH scales, neutralization reactions, and the importance of acids and bases in everyday life and industrial applications of acids and bases. By the end of this article, you'll have a comprehensive understanding of acid-base theory and its significance in chemistry and beyond.

      Historical Development of Acid-Base Theories

      The evolution of acid-base theories marks a significant progression in our understanding of chemical reactions. This journey begins with the groundbreaking Arrhenius theory and advances to the more comprehensive Bronsted-Lowry theory, each offering unique insights into the nature of acids and bases.

      Svante Arrhenius, a Swedish chemist, proposed the Arrhenius theory in 1884. This theory defined acids as substances that produce hydrogen ions (H+) when dissolved in water, while bases were substances that produce hydroxide ions (OH-) in aqueous solutions. Arrhenius' concept was revolutionary for its time, providing a clear explanation for many acid-base reactions observed in aqueous environments.

      The Arrhenius theory successfully explained numerous phenomena, such as the neutralization reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH) to form water and salt. It also clarified why acids conduct electricity in water due to the presence of free hydrogen ions. However, the theory had limitations. It was confined to aqueous solutions and couldn't explain acidic or basic behavior in non-aqueous environments. Additionally, it failed to account for substances like ammonia (NH3), which exhibit basic properties without producing hydroxide ions.

      Recognizing these limitations, Johannes Nicolaus Brønsted and Thomas Martin Lowry independently proposed a more expansive theory in 1923, known as the Brønsted-Lowry theory. This theory redefined acids and bases in terms of proton transfer. According to Brønsted-Lowry, an acid is a proton donor, while a base is a proton acceptor. This definition broadened the scope of acid-base chemistry beyond aqueous solutions and provided a more general framework for understanding these reactions.

      The Brønsted-Lowry theory introduced several key improvements over the Arrhenius theory. Firstly, it could explain acid-base reactions in non-aqueous solvents, addressing a major limitation of the Arrhenius theory. Secondly, it accounted for the behavior of amphoteric substances compounds that can act as both acids and bases depending on the reaction context. Water, for instance, can act as either a proton donor (acid) or a proton acceptor (base), a concept not fully captured by the Arrhenius theory.

      Another significant advancement of the Brønsted-Lowry theory was the introduction of conjugate acid-base pairs. When an acid donates a proton, it becomes its conjugate base, and when a base accepts a proton, it becomes its conjugate acid. This concept provided a deeper understanding of the reversibility of acid-base reactions and the equilibrium that exists in these systems.

      The Brønsted-Lowry theory also elegantly explained the behavior of Lewis bases like ammonia. While ammonia doesn't produce hydroxide ions in water (as required by the Arrhenius definition of a base), it clearly acts as a base by accepting protons from water molecules, forming ammonium ions (NH4+) and hydroxide ions. This observation aligns perfectly with the Brønsted-Lowry concept of bases as proton acceptors.

      Despite its improvements, the Brønsted-Lowry theory still had limitations. It couldn't explain all types of acid-base reactions, particularly those not involving proton transfer. This led to further developments in acid-base theory, such as the Lewis acid-base theory, which defines acids and bases in terms of electron pair acceptance and donation.

      In conclusion, the progression from Arrhenius theory to Brønsted-Lowry theory represents a significant advancement in our understanding of acid-base chemistry. While Arrhenius laid the foundation by explaining acid-base behavior in aqueous solutions, Brønsted and Lowry expanded this concept to a broader range of chemical environments and reactions. This evolution in acid-base theories has been crucial in developing our modern understanding of chemical reactions and continues to be fundamental in fields ranging from biochemistry to materials science.

      The Role of Water in Acid-Base Reactions

      Water plays a crucial role in acid-base reactions, serving as both a solvent and an active participant in the chemical processes. Its importance in these reactions stems from its unique molecular structure and properties, particularly its ability to form hydronium ions (H3O+) and hydroxide ions (OH-). Understanding the behavior of water in acid-base reactions is fundamental to grasping the principles of aqueous chemistry and the concept of pH.

      At the heart of water's significance in acid-base reactions is its ability to undergo autoionization. In pure water, a small fraction of water molecules spontaneously dissociate into hydronium ions and hydroxide ions. This process is represented by the equation: 2H2O H3O+ + OH-. The hydronium ion, formed when a water molecule accepts a proton (H+), is the primary species responsible for acidic properties in aqueous solutions. Conversely, the hydroxide ion, formed when a water molecule loses a proton, is the key player in basic or alkaline solutions.

      The formation of hydronium ions is particularly important in acid-base reactions. When an acid is dissolved in water, it releases protons (H+) which immediately combine with water molecules to form hydronium ions. For example, when hydrochloric acid (HCl) is added to water, the reaction can be written as: HCl + H2O H3O+ + Cl-. The hydronium ion is the actual acid species in aqueous solutions, responsible for the characteristic properties of acids such as sour taste and reactivity with metals.

      Similarly, the hydroxide ion plays a crucial role in basic solutions. When a base is dissolved in water, it either directly provides hydroxide ions or accepts protons from water molecules, thereby increasing the concentration of OH- ions. For instance, when sodium hydroxide (NaOH) is dissolved in water, it dissociates completely: NaOH Na+ + OH-. The presence of these hydroxide ions gives basic solutions their characteristic properties, such as a slippery feel and the ability to neutralize acids.

      One of the most fascinating aspects of water in acid-base chemistry is its amphoteric nature. An amphoteric substance is one that can act as both an acid and a base, depending on the context of the reaction. Water exemplifies this property perfectly. In the presence of an acid, water acts as a base by accepting protons to form hydronium ions. Conversely, in the presence of a base, water acts as an acid by donating protons, leading to an increase in hydroxide ion concentration.

      This amphoteric behavior of water is crucial in many biological and chemical processes. It allows water to participate in a wide range of reactions, maintaining pH balance in living organisms and enabling the formation of buffer solutions. Buffer solutions, which resist changes in pH when small amounts of acid or base are added, rely on the amphoteric nature of water to function effectively.

      The concept of pH, which measures the acidity or basicity of a solution, is directly related to the concentrations of hydronium and hydroxide ions in water. In pure water at 25°C, the concentrations of H3O+ and OH- are equal, each at 1 × 10^-7 mol/L, resulting in a neutral pH of 7. Any increase in hydronium ion concentration lowers the pH below 7, indicating an acidic solution, while an increase in hydroxide ion concentration raises the pH above 7, indicating a basic solution.

      Understanding the role of water in acid-base reactions is essential for various fields, including chemistry, biology, environmental science, and medicine. It explains phenomena such as the natural pH of rainwater, the functioning of antacids in the stomach, and the importance of maintaining proper pH levels in aquatic ecosystems. Moreover, this knowledge is crucial in industrial processes, water treatment, and the development of pharmaceuticals.

      In conclusion, water's ability to form hydronium and hydroxide ions, coupled with its amphoteric nature, makes it an indispensable component in acid-base reactions. Its unique properties enable a wide range of chemical processes that are fundamental to life and numerous technological applications. As we continue to explore and understand the intricacies of acid-base chemistry, the central role of water remains a cornerstone of our scientific knowledge.

      Understanding pH and pOH

      pH and pOH are fundamental concepts in chemistry that help us quantify the acidity and basicity of solutions. These measurements are crucial in various fields, from environmental science to biology and industrial processes. Let's delve into the definitions, mathematical expressions, and significance of pH and pOH.

      pH is defined as the negative logarithm (base 10) of the hydrogen ion concentration [H+] in a solution. Mathematically, it is expressed as:

      pH = -log[H+]

      This logarithmic scale ranges from 0 to 14, with 7 being neutral. Values below 7 indicate acidity, while values above 7 indicate basicity. The logarithmic scale nature of the pH scale means that each unit change represents a tenfold difference in hydrogen ion concentration. For example, a solution with a pH of 4 is ten times more acidic than a solution with a pH of 5, and 100 times more acidic than a solution with a pH of 6.

      To illustrate this concept, let's consider some common substances:

      • Lemon juice: pH 2 (highly acidic)
      • Coffee: pH 5 (mildly acidic)
      • Pure water: pH = 7 (neutral)
      • Baking soda solution: pH 9 (mildly basic)
      • Household ammonia: pH 11 (strongly basic)

      pOH, on the other hand, is the negative logarithm of the hydroxide ion concentration [OH-] in a solution. It is mathematically expressed as:

      pOH = -log[OH-]

      pOH is complementary to pH, and their sum always equals 14 in aqueous solutions at 25°C:

      pH + pOH = 14

      This relationship allows us to calculate pOH from pH and vice versa. For instance, if a solution has a pH of 4, its pOH would be 10 (14 - 4 = 10).

      The significance of pH and pOH lies in their ability to provide a standardized measure of acidity and basicity. This standardization is crucial for:

      1. Comparing different solutions: The logarithmic scale allows for easy comparison of vastly different concentrations.
      2. Biological processes: Many biochemical reactions are pH-dependent, and organisms require specific pH ranges to function properly.
      3. Environmental monitoring: pH is a key indicator of water quality in natural and industrial settings.
      4. Industrial applications: Many manufacturing processes, such as food production and wastewater treatment, rely on precise pH control.

      The logarithmic scale nature of the pH scale is particularly useful when dealing with very small concentrations of hydrogen or hydroxide ion concentration. For example, pure water at 25°C has a hydrogen ion concentration of 1 × 10^-7 mol/L. Instead of working with this small number, we can express it as a pH of 7, which is much more manageable.

      Understanding pH and pOH is essential for anyone working with chemical solutions or studying chemical reactions. These concepts provide a quantitative way to describe the acidity or basicity of a solution, which is crucial in many scientific and practical applications. By mastering these fundamental principles, one can better understand and control chemical processes in various fields, from environmental science to medicine and beyond.

      Acid-Base Properties and Their Causes

      Acids and bases are fundamental chemical substances that play crucial roles in various aspects of our lives, from the food we eat to the products we use daily. Understanding their properties and the underlying causes of their behavior is essential in chemistry and many practical applications. This article will explore the properties of acids and bases, focusing on how hydronium ions contribute to acidic properties and hydroxide ions to basic properties.

      Acids are substances that can donate protons (H+ ions) in aqueous solutions. When an acid dissolves in water, it releases these protons, which combine with water molecules to form hydronium ions (H3O+). The presence of these hydronium ions is responsible for the characteristic acidic properties. Some common examples of acids include hydrochloric acid (HCl), sulfuric acid (H2SO4), and acetic acid (CH3COOH).

      The acidic properties caused by hydronium ions include:

      • Sour taste: Acids have a distinctive sour flavor, which is why we perceive foods like lemons and vinegar as tart.
      • Reactivity with metals: Acids can react with certain metals to produce hydrogen gas.
      • pH values below 7: The concentration of hydronium ions determines the pH of a solution, with lower values indicating higher acidity.
      • Ability to conduct electricity: Aqueous solutions of acids can conduct electricity due to the presence of mobile ions.

      On the other hand, bases are substances that can accept protons or donate hydroxide ions (OH-) in aqueous solutions. When a base dissolves in water, it either accepts protons from water molecules, leaving behind hydroxide ions, or directly releases hydroxide ions. The presence of these hydroxide ions is responsible for the characteristic basic properties. Common examples of bases include sodium hydroxide (NaOH), potassium hydroxide (KOH), and ammonia (NH3).

      The basic properties caused by hydroxide ions include:

      • Bitter taste: Bases have a characteristic bitter taste, though it's important to note that tasting chemicals is not recommended for safety reasons.
      • Slippery feel: Many bases feel slippery to the touch, which is why soap (a mild base) has this property.
      • pH values above 7: The concentration of hydroxide ions determines the pH of a solution, with higher values indicating higher basicity.
      • Ability to conduct electricity: Like acids, aqueous solutions of bases can conduct electricity due to the presence of mobile ions.

      The interaction between acids and bases is a fundamental concept in chemistry. When an acid and a base are combined, they undergo a neutralization reaction, producing water and a salt. This process involves the hydronium ions from the acid combining with the hydroxide ions from the base to form water molecules.

      In everyday life, we encounter acids and bases in various forms. Citric acid in fruits, ascorbic acid (vitamin C) in supplements, and carbonic acid in carbonated beverages are examples of common acids. Bases are found in many cleaning products, such as ammonia-based cleaners and baking soda (sodium bicarbonate). Understanding the properties of these substances helps us use them safely and effectively in different applications.

      The strength of acids and bases is determined by their ability to produce hydronium or hydroxide ions in solution. Strong acids and bases ionize completely in water, while weak acids and bases only partially ionize. This concept is crucial in understanding chemical reactions and predicting the behavior of different substances in various environments.

      In conclusion, the properties of acids and bases are fundamentally linked to the presence of hydronium and hydroxide ions, respectively. These ions give rise to characteristic behaviors that we observe in everyday life and utilize in numerous scientific and industrial applications. By understanding the role of these ions, we can better appreciate the chemistry behind many natural phenomena and human-made processes.

      Neutralization Reactions

      Neutralization reactions are fundamental chemical processes that occur between acids and bases, resulting in the formation of water and salts. These reactions play a crucial role in various applications, from everyday life to industrial processes. Understanding neutralization is essential for anyone interested in chemistry or its practical applications.

      At its core, a neutralization reaction involves the combination of an acid and a base, which results in the production of water and a salt. The term "neutralization" refers to the process of an acidic solution becoming less acidic (or a basic solution becoming less basic) as the reaction progresses. In a complete neutralization, the resulting solution has a pH of 7, which is considered neutral.

      The general equation for a neutralization reaction can be written as:

      Acid + Base Salt + Water

      For example, when hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH), the result is sodium chloride (NaCl) and water (H2O):

      HCl + NaOH NaCl + H2O

      This reaction demonstrates how the hydrogen ions (H+) from the acid combine with the hydroxide ions (OH-) from the base to form water, while the remaining ions form a salt.

      Neutralization reactions are significant in various applications. In the medical field, antacids work by neutralizing excess stomach acid to relieve indigestion and heartburn. In agriculture, farmers use lime (calcium hydroxide) to neutralize acidic soils, improving crop growth. Environmental scientists employ neutralization techniques to treat acid rain and manage industrial waste, ensuring the pH of water bodies remains balanced for aquatic life.

      In the laboratory, titration is a common analytical technique that utilizes neutralization reactions to determine the concentration of unknown acid or base solutions. This method is crucial in quality control processes across various industries, including pharmaceuticals and food production.

      The importance of neutralization extends to industrial processes as well. In water treatment plants, acids and bases are used to adjust the pH of drinking water to safe levels. The paper industry relies on neutralization reactions to control the acidity of paper products, ensuring their quality and longevity.

      Understanding acid-base neutralization is also vital in the development of buffer solutions, which resist changes in pH when small amounts of acid or base are added. These buffers are essential in biological systems, maintaining the proper pH for enzymatic reactions and cellular functions.

      In conclusion, neutralization reactions between acids and bases are fundamental processes in chemistry with wide-ranging applications. From maintaining our health to preserving the environment and supporting various industries, these reactions play a crucial role in our daily lives and scientific advancements. The formation of water and salts through these reactions underscores the importance of understanding and harnessing the power of acid-base chemistry in solving real-world problems and advancing technological innovations.

      Conclusion

      In conclusion, this article has provided a comprehensive overview of acid-base theory, a fundamental concept in chemistry. We've explored the key definitions, including acids and bases, and pH, and discussed the three major models: Arrhenius, Brønsted-Lowry, and Lewis. Understanding these theories is crucial for grasping chemical reactions and their applications in various fields, from biology to environmental science. We've also touched on the importance of pH in everyday life and industrial processes. To reinforce your understanding and get a visual representation of these concepts, we strongly encourage you to watch the introduction video accompanying this article. It offers a concise yet thorough explanation of acid-base theory, making complex ideas more accessible. By mastering these principles, you'll be better equipped to tackle more advanced chemistry topics and appreciate the role of acids and bases in our world.

      Understanding the role of chemical reactions in various processes is essential for students and professionals alike. This knowledge not only helps in academic pursuits but also in practical applications in industries such as pharmaceuticals, agriculture, and environmental management.

      Introduction to Acid-Base Theory

      What are acids and bases? Arrhenius theory.

      Step 1: Understanding Bonding and Intermolecular Forces

      To understand acids and bases, we first need to grasp the concept of bonding and intermolecular forces. Bonding refers to the forces of attraction between atoms and molecules, which are attractive in nature. These bonds can be covalent, where atoms share electrons, or ionic, where atoms transfer electrons to achieve stability. Intermolecular forces, on the other hand, are the forces of attraction between molecules. These include van der Waals forces (such as London dispersion forces and dipole-dipole interactions) and hydrogen bonding. These forces play a crucial role in determining the properties of substances, including their behavior as acids or bases.

      Step 2: Differentiating Between Intermolecular and Intramolecular Forces

      It's important to distinguish between intermolecular and intramolecular forces. Intermolecular forces act between molecules, while intramolecular forces act within a molecule. For example, in a carbon dioxide (CO2) molecule, the covalent bonds between carbon and oxygen atoms are intramolecular forces. In contrast, the forces of attraction between different CO2 molecules are intermolecular forces. Understanding this distinction helps us comprehend how substances interact at a molecular level, which is essential for understanding acid-base behavior.

      Step 3: Exploring the Arrhenius Theory of Acids and Bases

      The Arrhenius theory is one of the earliest and simplest theories to define acids and bases. According to this theory, an acid is a substance that increases the concentration of hydrogen ions (H+) in an aqueous solution, while a base is a substance that increases the concentration of hydroxide ions (OH-) in an aqueous solution. This theory helps explain the behavior of acids and bases in water and provides a foundation for understanding their chemical properties.

      Step 4: Applying the Particle Model to Acids and Bases

      Using the particle model, we can visualize how acids and bases behave at a molecular level. In an aqueous solution, an acid dissociates to release H+ ions, which can interact with water molecules to form hydronium ions (H3O+). For example, hydrochloric acid (HCl) dissociates in water to form H+ and Cl- ions. Similarly, a base dissociates to release OH- ions, which can interact with H+ ions to form water. For instance, sodium hydroxide (NaOH) dissociates in water to form Na+ and OH- ions.

      Step 5: Understanding the Relationship Between Bonding, Structure, and Properties

      The bonding structure of a substance significantly influences its properties, including its behavior as an acid or base. For example, the strength of an acid or base depends on how easily it can dissociate in water to release H+ or OH- ions. Strong acids, like sulfuric acid (H2SO4), dissociate completely in water, releasing a high concentration of H+ ions. In contrast, weak acids, like acetic acid (CH3COOH), only partially dissociate, releasing fewer H+ ions. Similarly, strong bases, like potassium hydroxide (KOH), dissociate completely to release OH- ions, while weak bases, like ammonia (NH3), only partially dissociate.

      Step 6: Summarizing the Key Concepts

      In summary, understanding acids and bases through the Arrhenius theory involves recognizing how substances dissociate in water to release H+ or OH- ions. This behavior is influenced by the bonding structure and intermolecular forces within the substance. By applying the particle model, we can visualize these interactions at a molecular level, helping us comprehend the properties and behavior of acids and bases in various contexts.

      FAQs

      Here are some frequently asked questions about acid-base theory:

      1. What is the concept of acids and bases?

        Acids and bases are fundamental chemical substances with distinct properties. Acids are proton donors that release hydrogen ions (H+) in solution, while bases are proton acceptors that can neutralize acids. Their interaction forms the basis of acid-base chemistry, which is crucial in many chemical processes.

      2. What is an acid and a base for beginners?

        For beginners, acids can be thought of as sour-tasting substances that can react with metals and turn litmus paper red. Bases, on the other hand, taste bitter, feel slippery, and turn litmus paper blue. Common examples of acids include lemon juice and vinegar, while examples of bases include soap and baking soda.

      3. What is the theory of acid and base?

        There are three main theories of acids and bases: Arrhenius, Brønsted-Lowry, and Lewis. The Arrhenius theory defines acids as substances that produce hydrogen ions in water and bases as substances that produce hydroxide ions. The Brønsted-Lowry theory expands this, defining acids as proton donors and bases as proton acceptors. The Lewis theory further broadens the concept, defining acids as electron pair acceptors and bases as electron pair donors.

      4. What does the Arrhenius theory explain?

        The Arrhenius theory explains acid-base behavior in aqueous solutions. It states that acids are substances that dissociate in water to produce hydrogen ions (H+), while bases are substances that dissociate to produce hydroxide ions (OH-). This theory successfully explains many acid-base reactions but is limited to aqueous solutions.

      5. How do you introduce a lesson on acids and bases?

        To introduce a lesson on acids and bases, start with everyday examples like lemon juice (acid) and soap (base). Demonstrate their properties using indicators like litmus paper. Explain the basic concepts of pH and neutralization reactions. Then, progress to more complex theories and applications in chemistry and daily life. Use visual aids, experiments, and real-world examples to make the concepts more engaging and understandable.

      Prerequisite Topics

      Understanding the fundamental concepts that lay the groundwork for more advanced topics is crucial in chemistry. When delving into the "Introduction to acid-base theory," it's essential to have a solid grasp of several prerequisite topics. These foundational concepts not only enhance your understanding but also provide the necessary context for comprehending the intricacies of acid-base interactions.

      One of the key prerequisites is solubility and ion concentration. This topic is vital because it helps you understand how substances dissolve in solutions and how the concentration of ions, particularly hydroxide ions, affects the acidity or basicity of a solution. Knowing how to calculate and interpret ion concentrations is fundamental when studying acid-base reactions and equilibria.

      Another crucial prerequisite is the concept of the logarithmic scale. In acid-base theory, we often deal with very small concentrations of hydrogen ions, which are more conveniently expressed using the pH scale a logarithmic measure of hydrogen ion concentration. Understanding logarithms is essential for interpreting pH values and performing calculations related to acid-base equilibria.

      Lastly, familiarity with conjugate acids and bases is indispensable when studying acid-base theory. This topic introduces you to the properties of acids and bases and how they interact in solution. It helps you understand the concept of proton transfer, which is at the heart of acid-base reactions, and prepares you for more complex discussions on acid-base equilibria and buffer solutions.

      By mastering these prerequisite topics, you'll be well-equipped to tackle the intricacies of acid-base theory. The knowledge of ion concentrations will help you understand how acids and bases dissociate in solution and affect pH. Your grasp of logarithmic scales will enable you to interpret and calculate pH values with ease. And your understanding of conjugate acids and bases will provide insight into the behavior of acids and bases in various reactions and equilibria.

      As you progress in your study of acid-base theory, you'll find these prerequisite concepts continually resurfacing, reinforcing their importance. They form the building blocks upon which more advanced concepts are constructed. For instance, when you encounter buffer solutions or titration curves, your understanding of ion concentrations and conjugate pairs will prove invaluable.

      In conclusion, taking the time to thoroughly understand these prerequisite topics will significantly enhance your learning experience as you delve into acid-base theory. It will allow you to approach the subject with confidence, making connections between new information and your existing knowledge base, ultimately leading to a deeper and more comprehensive understanding of this fundamental area of chemistry.

      In this lesson, we will learn:

      • To understand how acids, bases and salts were originally and currently defined in chemistry.
      • To understand how acid-base theory is related to pH and the amphoteric nature of water.
      • To understand what H+ ions do in solution and how acid-base reactions depend on them.

      Notes:

      • Even though acids, bases and salts are common and well-studied in chemistry, there are different sets of definitions that depend on what you focus on in an acid-base reaction. The chemicals that we call ‘acids’ (e.g. hydrochloric acid) and ‘bases’ (such as sodium hydroxide) all have something in common which leads to the acidic/basic properties that we know.

      • The original definitions of acid, base and salt come from Arrhenius theory (named after Svante Arrhenius), which contains a few basic ideas about how acids and bases are related:
        • An acid is any substance which, in water, produces hydrogen ions (H+).
        • A base is any substance which, in water, produces hydroxide ions (OH-).
        • A salt is any substance which is the product of an acid-base reaction.
          Practically, this means any ionic compound that isn't an acid or base is a salt.

      • These definitions were the first of their kind in describing the properties of acids and bases (Arrhenius did a lot for chemistry!) and are the reason why many chemistry teachers give the following (not perfect) simple rule to 'spot' acids and bases:
        • Any ionic compound beginning with H is an acid.
        • Any ionic compound with an OH group is a base.

      • Today there are two theories or 'points of view’ to acid-base reactions. Brønsted-Lowry theory which is closely related to the original Arrhenius theory is much more common:
        • Brønsted-Lowry acid-base theory is about protons.
          • A Brønsted-Lowry acid is a proton (H+) donor such as HCl.
          • A Brønsted-Lowry base is a proton acceptor.

        • There is also Lewis theory is about electrons.
          • A Lewis acid is an electron pair acceptor such as BH3.
          • A Lewis base is an electron pair donor.

      • Looking at the Arrhenius and Brønsted-Lowry definitions of acids and bases, you will see they are about the hydrogen ion (H+). This is actually not what causes the properties acids are known for. The properties that ‘acids’ have are from the hydronium ion, H3O+. This gets produced rapidly when protons mix with water.
        • Hydrogen atoms have only one proton and one electron. Therefore when 'hydrogen ions (H+) are produced' by a substance (like hydrochloric acid, HCl) dissolving in water, what is actually released is just a single proton without the electron it used to have – quite literally just one tiny proton with nothing surrounding it!
        • The charge density of a lone proton with nothing surrounding it is INCREDIBLY high – an individual proton of +1 charge is extremely small compared to a hydrogen atom with the same positive nucleus with the electron 'cloud' orbiting it. This makes H+ extremely reactive and it will interact with anything remotely negative nearby. Remember, this is all happening dissolved in water.
        • This tiny ultra-concentrated proton immediately interacts with a negative lone pair on the oxygen atom of a nearby water molecule. In doing this, a hydronium ion, H3O+, is formed. This process can be described with the equation:
          • H+(aq) + H2O (l) → H3O+
        • The process can be explained in two ways:
          • Acids release H+ ions in solution which 'protonate' water molecules.
          • Acids release H+ in solution which are ‘hydrated’ (added to by water).

      • The reverse process occurs when bases are added to solution. When bases (that produce OH-) are dissolved, their strong negative charge means they react to neutralize H+ and H3O+ species, which decreases their concentration in the solution. This can also deprotonate neutral water molecules. See the equation below:
        H3O+ (aq) + OH- (aq) → 2H2O (l)
        • This process of decreasing hydronium ion concentration is what makes the pH rise. Eventually a lack of hydronium ions to neutralize the hydroxide ions means there will be free hydroxide (OH-) ions in solution, which causes the basic properties chemists observe.

      • We have seen above that water can act as a base in acidic conditions; it accepts H+ ions to form the hydronium ion. However, water also can act as an acid in basic conditions; it donates H+ to bases in solution, forming a hydroxide ion as a result.
        • Water’s ability to act as both acid and base makes it an amphoteric molecule.

      • Acid-base theory is important for understanding how pH is measured. Remember, the definition of acidic is pH < 7, while basic is pH > 7.
        • The equation to find pH of a solution is:
          pH = -log[H3O+]

        • This is an inverse logarithmic expression which means the following:
          • Inverse: as the concentration of hydronium ions increases, pH decreases.
          • Logarithmic: To change the pH by 1, you have to change [H3O+] by a factor or ten. For example, a solution of pH 4 is ten times more concentrated with H3O+ than a solution of pH 5.
        • This equation explains why acidic solutions have a low pH value, while basic solutions (where [H3O+] is low) have a high pH value.