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Strong and weak acids and bases

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Intros
Lessons
  1. What is acid and base strength?
  2. Difference between concentration and strength.
  3. The degree of dissociation.
  4. How dissociation affects acid and base strength.
  5. Structure and properties that affect strength.
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Examples
Lessons
  1. Understand the difference between acid strength and concentration.
    1. Explain the difference between a concentrated weak acid and a dilute strong acid.
    2. Give an example of a strong acid and the chemical process that occurs when it is dissolved in water.
Topic Notes
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Introduction to Strong and Weak Acids and Bases

Welcome to our exploration of strong and weak acids and bases! This fundamental chemistry concept is crucial for understanding many chemical reactions. Strong bases and acids dissociate completely in water, releasing all their hydrogen or hydroxide ions. Examples include hydrochloric acid (HCl) and sodium hydroxide (NaOH). In contrast, weak acids and bases only partially dissociate, maintaining an equilibrium between the dissociated and undissociated forms. Acetic acid (CH3COOH) and ammonia (NH3) are common examples. The introduction video we'll watch shortly provides an excellent visual representation of these concepts, making it easier to grasp the differences in dissociation. Understanding strong and weak acids and bases is essential for predicting chemical behaviors, balancing equations, and solving pH problems. As we delve deeper into this topic, you'll discover how these principles apply to various real-world scenarios, from household cleaners to biological systems. Let's begin this exciting journey into the world of acids and bases!

Understanding Dissociation in Strong and Weak Acids

What is Acid Dissociation?

Acid dissociation is a fundamental concept in chemistry that describes how acids break apart in water. When an acid dissolves in water, it releases hydrogen ions (H+), which are responsible for the acidic properties. The extent to which an acid dissociates is crucial in determining its strength and behavior in solution.

Strong Acids vs. Weak Acids

Acids are classified as either strong or weak based on their ability to dissociate in water. Strong acids dissociation completely or almost completely in aqueous solutions, while weak acids only partially dissociate. This difference in dissociation behavior leads to significant variations in their properties and reactivity.

Degree of Dissociation

The degree of dissociation refers to the extent to which an acid breaks apart into ions when dissolved in water. Strong acids dissociation have a high degree of dissociation, typically close to 100%, meaning almost all of their molecules separate into ions. Weak acids, on the other hand, have a much lower degree of dissociation, often less than 5%.

Examples of Strong Acids

Two common examples of strong acids are hydrochloric acid (HCl) and nitric acid (HNO3). These acids are known for their complete dissociation in water. When HCl dissolves in water, it forms H+ and Cl- ions, while HNO3 dissociates into H+ and NO3- ions. Other examples of strong acids include sulfuric acid (H2SO4) and hydrobromic acid (HBr).

Why Strong Acids Dissociate Completely

Strong acids dissociate completely in water due to their molecular structure and the nature of their chemical bonds. The bond between the hydrogen atom and the rest of the molecule in strong acids is relatively weak compared to the strong attraction between the ions and water molecules. This allows water to easily break the bond, separating the acid into its constituent ions.

Factors Affecting Dissociation

Several factors influence the dissociation of acids in water. These include the strength of the acid-base pair, the polarity of the solvent, temperature, and concentration. Strong acids have a greater tendency to dissociate because their conjugate bases are very weak. The polar nature of water also plays a crucial role in facilitating the dissociation process.

Implications of Dissociation

The degree of dissociation has significant implications for an acid's properties and behavior in solution. Strong acids dissociation, due to their complete dissociation, produce a high concentration of hydrogen ions, resulting in very low pH values. They react vigorously with bases and metals, and their solutions are highly conductive. Weak acids, with their partial dissociation, have less extreme properties and are often used in buffer solutions.

Measuring Acid Strength

The strength of an acid is typically measured by its acid dissociation constant (Ka) or its negative logarithm (pKa). Strong acids have very large Ka values (or very small pKa values), indicating a high degree of dissociation. Weak acids have much smaller Ka values, reflecting their partial dissociation in solution.

Applications in Chemistry and Beyond

Understanding acid dissociation is crucial in various fields, including chemistry, biology, and environmental science. It plays a vital role in pH regulation in living organisms, water treatment processes, and industrial applications. The concept of dissociation is also important in developing pharmaceuticals, as it affects drug absorption and efficacy. Applications of acid dissociation are vast and varied, impacting many aspects of science and technology.

Conclusion

The concept of dissociation in strong and weak acids is fundamental to understanding acid-base chemistry. Strong acids like HCl and HNO3 dissociate completely in water, while weak acids only partially dissociate. This difference in behavior leads to distinct properties and applications for each type of acid. By comprehending the factors that influence dissociation and its implications, we can better predict and control chemical reactions in various applications of acid dissociation.

Characteristics of Strong and Weak Bases

Properties of Strong Bases

Strong bases are ionic compounds that completely dissociate in water, releasing hydroxide ions (OH-). They have a high pH value, typically above 12, and are highly reactive. Strong bases can conduct electricity efficiently in aqueous solutions due to the presence of free ions. They have a bitter taste and feel slippery to the touch. Common examples of strong bases include potassium hydroxide (KOH) and calcium hydroxide (Ca(OH)2).

Properties of Weak Bases

Weak bases, on the other hand, only partially dissociate in water. They have a lower pH value compared to strong bases, usually between 8 and 11. Weak bases are less reactive and conduct electricity less efficiently in solution. They often have a slightly bitter taste and may feel slightly slippery. Examples of weak bases include ammonia (NH3) and methylamine (CH3NH2).

How to Determine if a Base is Strong or Weak

Several methods can be used to determine whether a base is strong or weak:

  • pH measurement: Strong bases have a pH closer to 14, while weak bases have a pH between 8 and 11.
  • Conductivity test: Strong bases conduct electricity more efficiently than weak bases in aqueous solutions.
  • Reaction rate: Strong bases react more vigorously with acids compared to weak bases.
  • Dissociation constant (Kb): Strong bases have a very large Kb value, while weak bases have a smaller Kb value.

Dissociation of Strong Bases in Water

When a strong base is dissolved in water, it undergoes complete dissociation. This means that all of the base molecules separate into their constituent ions. For example, when potassium hydroxide (KOH) dissolves in water, it completely dissociates into potassium ions (K+) and hydroxide ions (OH-):

KOH (s) K+ (aq) + OH- (aq)

This complete dissociation results in a high concentration of hydroxide ions in the solution, leading to a high pH and strong basic properties.

Examples of Strong Bases

Potassium Hydroxide (KOH)

KOH is a common strong base used in various industrial applications. When dissolved in water, it completely dissociates, releasing potassium ions (K+) and hydroxide ions (OH-). KOH has a high pH and is highly reactive, making it useful in soap production, battery manufacturing, and as a cleaning agent.

Calcium Hydroxide (Ca(OH)2)

Also known as slaked lime, Ca(OH)2 is another strong base that dissociates completely in water. It releases calcium ions (Ca2+) and hydroxide ions (OH-) when dissolved. Ca(OH)2 is used in water treatment, construction (as lime mortar), and in the production of paper.

Importance of Understanding Base Strength

Distinguishing between strong and weak bases is crucial in various fields, including chemistry, biology, and environmental science. The strength of a base affects its reactivity, pH, and overall behavior in chemical reactions. This knowledge is essential for:

  • Predicting reaction outcomes in chemical processes
  • Selecting appropriate bases for specific applications
  • Understanding and controlling pH levels in various systems
  • Ensuring safety when handling basic substances
  • Optimizing industrial processes that involve bases

In conclusion, understanding the properties of strong and weak bases, their dissociation behavior, and how to identify them is fundamental in chemistry and related fields. Strong bases like KOH and Ca(OH)2 play significant roles in various applications due to their complete dissociation and high reactivity.

Factors Influencing Acid and Base Strength

Understanding the factors that determine whether an acid or base is strong or weak is crucial in chemistry. The strength of acids and bases plays a significant role in various chemical reactions and processes. In this section, we'll explore the key factors that influence acid and base strength, focusing on electronegativity, resonance, and inductive effects. We'll also use examples like HNO3 (nitric acid) and H3PO4 (phosphoric acid) to illustrate these concepts.

Electronegativity

Electronegativity is a fundamental factor in determining acid strength. It refers to an atom's ability to attract electrons in a chemical bond. In acids, the more electronegative the atom bonded to hydrogen, the stronger the acid tends to be. This is because a highly electronegative atom pulls electrons away from the hydrogen, making it easier for the hydrogen to be released as a proton (H+).

For example, HNO3 is a strong acid because the nitrogen atom, bonded to three oxygen atoms, becomes highly electronegative. This increased electronegativity makes it easier for HNO3 to donate a proton, resulting in its classification as a strong acid.

Resonance

Resonance is another critical factor affecting acid and base strength. It refers to the delocalization of electrons across multiple atoms in a molecule. When an acid or base can distribute its charge over a larger area through resonance, it becomes more stable. For acids, this stability means they are less likely to hold onto their protons, making them stronger acids.

Consider the example of H3PO4 (phosphoric acid). While it's not as strong as HNO3, it's still a relatively strong acid. The phosphate ion (PO4^3-) formed after losing protons has resonance structures that distribute the negative charge, contributing to its acid strength.

Inductive Effect

The inductive effect refers to the transmission of charge through a chain of atoms in a molecule. This effect can either withdraw or donate electrons, influencing the molecule's ability to release or accept protons. Electron-withdrawing groups increase acid strength by making it easier for the molecule to lose protons, while electron-donating groups have the opposite effect.

In the case of HNO3, the oxygen atoms attached to nitrogen create a strong electron-withdrawing effect. This inductive effect enhances the acid's ability to donate protons, contributing to its classification as a strong acid.

Size and Polarizability

The size and polarizability of atoms also play a role in acid and base strength. Generally, as we move down a group in the periodic table, acid strength increases for binary acids (HX, where X is a halogen). This is because larger atoms are more polarizable, meaning their electron clouds can be more easily distorted. This increased polarizability makes it easier for the acid to release protons.

Bond Strength

The strength of the bond between hydrogen and the rest of the molecule is another crucial factor. Weaker bonds make it easier for the hydrogen to be released as a proton, resulting in stronger acids. Conversely, stronger bonds lead to weaker acids. This factor is closely related to bond strength in acids but also considers the overall molecular structure.

Solvent Effects

While not an intrinsic property of the acid or base itself, the solvent in which a substance is dissolved can significantly affect its apparent strength. For example, in water, HNO3 is a strong acid, but in a less polar solvent, its behavior might change. The solvent's ability to stabilize the resulting ions plays a crucial role in determining the observed acid or base strength.

Conclusion

In conclusion, the strength of acids and bases is determined by a complex interplay of factors, including electronegativity, resonance, inductive effects, size and polarizability, bond strength in acids, and solvent effects. Understanding these factors helps chemists predict and explain the behavior of acids and bases in various reactions and environments. Examples like HNO3 and H3PO4 illustrate these concepts effectively.

Polyprotic Acids and Their Behavior

Polyprotic acids are a fascinating class of compounds in chemistry that can donate more than one proton per molecule during acid-base reactions. These acids undergo a process called stepwise dissociation, where they release protons in distinct stages. To understand this concept better, let's explore the behavior of polyprotic acids, using sulfuric acid (H2SO4) as our primary example.

Polyprotic acids are characterized by their ability to donate multiple protons in aqueous solutions. This property sets them apart from monoprotic acids, which can only donate a single proton. The stepwise dissociation of polyprotic acids occurs because each successive proton becomes increasingly difficult to remove due to the growing negative charge on the conjugate base.

Let's examine H2SO4, a diprotic acid, to illustrate the concept of stepwise dissociation. The first dissociation of H2SO4 can be represented as:

H2SO4 + H2O H3O+ + HSO4-

This first dissociation is very strong, with a large acid dissociation constant (Ka1) of about 1 × 10^3. This means that H2SO4 readily donates its first proton in aqueous solutions, making it a strong acid for this step.

The second dissociation, however, tells a different story:

HSO4- + H2O H3O+ + SO4^2-

This step has a much smaller acid dissociation constant (Ka2) of about 1.2 × 10^-2, indicating that the second proton is much less readily donated. This significant difference in dissociation strength between the first and second steps is a hallmark of polyprotic acids.

The stark contrast between the first and second dissociations of H2SO4 has important implications for acid strength classification. While H2SO4 is considered a strong acid due to its first dissociation, the HSO4- ion produced in the first step behaves as a weak acid in the second dissociation. This dual nature highlights the complexity of polyprotic acids and their behavior in solution.

The stepwise dissociation of polyprotic acids affects various aspects of their chemistry, including their titration curves of polyprotic acids, buffer capacity, and interactions with other substances. For instance, when titrating H2SO4 with a strong base, two distinct equivalence points can be observed, corresponding to the neutralization of each proton.

Understanding the behavior of polyprotic acids is crucial in many areas of chemistry and biology. For example, phosphoric acid (H3PO4), another polyprotic acid, plays a vital role in biological systems, including DNA structure and energy metabolism. Its three-step dissociation process influences its function in these complex biochemical environments.

In conclusion, the concept of polyprotic acids and their stepwise dissociation, as exemplified by H2SO4, reveals the intricate nature of acid-base chemistry. The significant difference between the first and second dissociations of H2SO4 demonstrates how a single compound can exhibit both strong and weak acid characteristics. This understanding is essential for accurately predicting and interpreting the behavior of polyprotic acids in various chemical and biological processes, making it a fundamental concept in advanced chemistry studies and research.

The Leveling Effect in Strong Acids and Bases

The leveling effect is a fundamental concept in acid-base chemistry that plays a crucial role in understanding the behavior of strong acids and bases in aqueous solutions. This phenomenon explains why all strong acids of the same concentration exhibit identical strength when dissolved in water, despite their inherent differences in molecular structure and composition.

To comprehend the leveling effect, it's essential to first understand the nature of strong acids and bases. Strong acids are substances that completely dissociate in water, releasing all of their protons (H+ ions) into the solution. Similarly, strong bases fully dissociate, producing hydroxide ions (OH-) in aqueous environments. Common examples of strong acids include hydrochloric acid (HCl), sulfuric acid (H2SO4), and nitric acid (HNO3), while strong bases include sodium hydroxide (NaOH) and potassium hydroxide (KOH).

The leveling effect occurs because water acts as a limiting factor in the dissociation process. When a strong acid is dissolved in water, it transfers its protons to water molecules, forming hydronium ions (H3O+). However, water can only accept a certain number of protons, regardless of the acid's initial strength. This limitation results in all strong acids reaching the same maximum acidity in aqueous solutions, which is determined by the concentration of hydronium ions.

For example, when hydrochloric acid (HCl) and perchloric acid (HClO4) are dissolved in water at the same concentration, they both dissociate completely. Despite perchloric acid being inherently stronger than hydrochloric acid in their pure forms, they produce identical concentrations of hydronium ions in water. This is because water acts as a proton acceptor, leveling the acidity of both acids to the same strength.

The same principle applies to strong bases, where the hydroxide ion concentration becomes the limiting factor. All strong bases of equal concentration will produce the same concentration of hydroxide ions in aqueous solutions, regardless of their intrinsic strength in their pure forms.

The leveling effect has significant implications for comparing strong acids and bases. Since all strong acids (or bases) of the same concentration exhibit identical strength in water, it becomes impossible to differentiate between them based on their acidity or basicity alone. This means that traditional methods of comparing acid strength, such as acid dissociation constants (Ka values), become irrelevant for strong acids in aqueous solutions.

Instead, when comparing strong acids or bases, other factors must be considered, such as their reactivity with specific substances, their ability to act as oxidizing or reducing agents, or their impact on reaction kinetics. For instance, while hydrochloric acid and sulfuric acid may have the same strength in water due to the leveling effect, sulfuric acid is a stronger oxidizing agent and can participate in additional reactions that hydrochloric acid cannot.

Understanding the leveling effect is crucial for chemists and researchers working with acids and bases in various applications. It explains why the pH scale has practical limits (typically 0 to 14) in aqueous solutions and why extremely strong acids or bases cannot be accurately measured using conventional pH meters. This knowledge is particularly important in fields such as analytical chemistry, environmental science, and chemical engineering, where precise control and measurement of acidity and basicity are essential.

In conclusion, the leveling effect demonstrates that water plays a significant role in determining the strength of acids and bases in aqueous solutions. It explains why all strong acids of the same concentration have identical strength in water and highlights the limitations of comparing strong acids and bases based solely on their acidity or basicity. By understanding this phenomenon, scientists can better interpret experimental results, design more effective chemical processes, and develop more accurate analytical techniques for working with acids and bases in various applications.

Conclusion: Key Takeaways on Strong and Weak Acids and Bases

In this article, we've explored the fundamental concepts of strong acids, weak acids, strong bases, and weak bases in acid-base chemistry. Understanding the differences between these substances is crucial for grasping their behavior in chemical reactions. We've learned that strong acids and bases dissociate completely in solution, while weak acids and bases only partially dissociate. This distinction affects their pH levels, reactivity, and overall chemical properties. The article emphasized the importance of dissociation constants and their role in determining acid or base strength. We also discussed common examples of each type and their applications in various fields. The introduction video provided a visual foundation for these concepts, making them more accessible. As you continue your journey in chemistry, remember that mastering these fundamental principles is essential for tackling more advanced topics. We encourage you to further explore acid-base chemistry and apply these concepts in practical scenarios to deepen your understanding.

Example:

What is acid and base strength? Difference between concentration and strength.

Step 1: Understanding Acid and Base Strength

Acid and base strength refers to the degree of ionization or dissociation of an acid or base in water. Strong acids and bases completely dissociate in water, meaning they release all of their hydrogen ions (H+) or hydroxide ions (OH-) respectively. Examples of strong acids include hydrochloric acid (HCl) and sulfuric acid (H2SO4), while examples of strong bases include sodium hydroxide (NaOH) and potassium hydroxide (KOH). On the other hand, weak acids and bases only partially dissociate in water. This means that only a fraction of their molecules release H+ or OH- ions. Examples of weak acids include acetic acid (CH3COOH) and citric acid, while examples of weak bases include ammonia (NH3) and methylamine (CH3NH2).

Step 2: Difference Between Concentration and Strength

It is crucial to distinguish between the concentration of an acid or base and its strength. Concentration refers to the amount of acid or base dissolved in a given volume of solution, typically expressed in moles per liter (M). For instance, a 1 M solution of hydrochloric acid contains 1 mole of HCl dissolved in 1 liter of water. Strength, as mentioned earlier, refers to the degree of ionization or dissociation of the acid or base in water. A strong acid or base will completely dissociate regardless of its concentration, while a weak acid or base will only partially dissociate even if it is highly concentrated. Therefore, you can have a dilute solution of a strong acid or a concentrated solution of a weak acid, and their behaviors in terms of ionization will still follow their respective definitions of strength.

Step 3: Practical Implications

Understanding the difference between concentration and strength is essential in various practical applications. For example, in titration experiments, knowing whether an acid or base is strong or weak helps in selecting appropriate indicators and calculating the pH of the solution. In industrial processes, the choice between using a strong or weak acid or base can affect the efficiency and safety of chemical reactions. Additionally, in biological systems, the strength of acids and bases can influence enzyme activity and metabolic pathways, as many biochemical reactions are pH-dependent.

Step 4: Examples and Comparisons

To further illustrate the difference, consider a 0.1 M solution of hydrochloric acid (a strong acid) and a 0.1 M solution of acetic acid (a weak acid). The hydrochloric acid will completely dissociate into H+ and Cl- ions, resulting in a high concentration of hydrogen ions and a low pH. In contrast, the acetic acid will only partially dissociate, resulting in a lower concentration of hydrogen ions and a higher pH compared to the hydrochloric acid solution of the same concentration. This example highlights how strength and concentration independently affect the properties of acidic and basic solutions.

Step 5: Summary

In summary, acid and base strength refers to the degree of ionization or dissociation in water, with strong acids and bases completely dissociating and weak acids and bases only partially dissociating. Concentration, on the other hand, refers to the amount of acid or base dissolved in a solution. Understanding the distinction between these two concepts is crucial for various scientific and practical applications, as it influences the behavior and properties of acidic and basic solutions.

FAQs

Here are some frequently asked questions about strong and weak acids and bases:

1. What are the 7 strong acids?

The 7 strong acids are: Hydrochloric acid (HCl), Hydrobromic acid (HBr), Hydroiodic acid (HI), Nitric acid (HNO3), Perchloric acid (HClO4), Sulfuric acid (H2SO4), and Chloric acid (HClO3). These acids completely dissociate in aqueous solutions.

2. How can you tell if an acid is strong or weak?

You can determine if an acid is strong or weak by:

  • Checking if it's one of the 7 strong acids
  • Looking at its acid dissociation constant (Ka) - strong acids have very large Ka values
  • Measuring its conductivity in solution - strong acids conduct electricity better
  • Observing its reaction rate with metals - strong acids react more vigorously

3. What is the difference between a strong acid and a weak acid dissociation?

Strong acids completely dissociate in water, releasing all their hydrogen ions. For example, HCl H+ + Cl-. Weak acids only partially dissociate, maintaining an equilibrium between the acid and its ions. For instance, CH3COOH H+ + CH3COO-. This partial dissociation is why weak acids have lower concentrations of hydrogen ions in solution.

4. Do strong bases dissociate in water?

Yes, strong bases completely dissociate in water. For example, sodium hydroxide (NaOH) dissociates into Na+ and OH- ions. This complete dissociation results in a high concentration of hydroxide ions, giving strong bases their characteristic properties.

5. What is the leveling effect in acids and bases?

The leveling effect occurs when a solvent (usually water) limits the strength of an acid or base. For strong acids in water, they all appear to have the same strength because water acts as a base, accepting protons to form H3O+. Similarly, strong bases are leveled to the strength of OH- in water. This effect makes it impossible to distinguish between the strengths of different strong acids or bases in aqueous solutions.

Prerequisite Topics for Understanding Strong and Weak Acids and Bases

To fully grasp the concept of strong and weak acids and bases, it's crucial to have a solid foundation in several key areas of chemistry. One of the most important prerequisites is understanding the acid dissociation constant. This fundamental concept helps explain why some acids and bases are stronger than others, and how they behave in solution.

The acid dissociation constant is closely related to the relative strength of acids and bases. By exploring this topic, students can gain insight into how molecular structure and bond strength influence the acidity or basicity of a compound. This knowledge is essential for predicting and explaining the behavior of various acids and bases in different chemical reactions.

Another critical area of study is titration curves. These graphical representations provide valuable information about the strength of acids and bases, as well as their behavior during neutralization reactions. Understanding titration curves is particularly important when dealing with polyprotic acids, which have multiple dissociable protons.

While it might not seem immediately obvious, an introduction to kinetics is also highly relevant to the study of strong and weak acids and bases. Reaction kinetics play a crucial role in acid-base reactions, influencing the rate at which protons are transferred and equilibrium is established. This knowledge helps explain why some acids and bases react more quickly or completely than others.

By mastering these prerequisite topics, students will be better equipped to tackle the complexities of strong and weak acids and bases. The acid dissociation constant provides a quantitative measure of acid strength, while the study of relative strengths offers insights into structural factors affecting acidity and basicity. Titration curves demonstrate the practical applications of these concepts in laboratory settings, and kinetics helps explain the dynamic nature of acid-base reactions.

Moreover, these foundational concepts are interconnected. For instance, the acid dissociation constant is directly related to the shape of titration curves, and kinetics influences how quickly equilibrium is reached in acid-base reactions. Understanding these relationships allows for a more comprehensive grasp of acid-base chemistry as a whole.

In conclusion, a thorough understanding of these prerequisite topics is essential for anyone looking to master the concept of strong and weak acids and bases. By building a strong foundation in these areas, students will be better prepared to tackle more advanced topics in acid-base chemistry and apply their knowledge to real-world problems in fields such as environmental science, biochemistry, and materials engineering.

In this lesson, we will learn:

  • To understand the difference between acid and base strength and concentration.
  • To identify strong and weak acids/bases by their degree of dissociation.
  • To understand how degree of dissociation leads to varying acid and base strength.
  • To understand the chemical structures and properties that influence dissociation and acid/base strength.

Notes:

  • In acid-base chemistry, there’s an important difference between strength and concentration. It is possible to have a highly concentrated ‘weak acid’ and a very dilute, quite harmless ‘strong acid’. Recall that:
    Concentration measures the amount of substance in a sample, compared to the amount of solvent (e.g. water) it is dissolved in. This can be changed easily: add more solvent to decrease concentration or add more solute to increase concentration.
    Acid and base strength (acidity or basicity) measures dissociation, where the degree of dissociation is how much a chemical compound splits from the complete compound (e.g. HCl) into its ions or components (e.g. H+ and Cl-). This can’t be changed because degree of dissociation is a core chemical property of something:
    • A strong acid or base experiences 100% dissociation into its ions when put in water. This means, in theory, every single molecule of the substance becomes aqueous ions: every HX molecule becomes H+ and X- (the conjugate base) while the base B becomes the conjugate acid HB+which accepted H+ (aq) in solution. Many mineral (inorganic) acids are strong acids, including:
      • Hydrochloric acid, HCl.
      • Sulfuric acid, H2SO4.
      • Nitric acid, HNO3.
    • Some strong bases include:
      • Potassium hydroxide, KOH.
      • Sodium hydroxide, NaOH.
      • Calcium hydroxide, Ca(OH)2.
    • A weak acid or base experiences partial dissociation into aqueous ions. This means that less than 100% of the molecules of acid or base dissociates into aqueous ions in water.
      • Carboxylic and other organic acids are weak acids.
      • Ammonia is an example of a weak base as it does not completely dissolve forming ammonium and hydroxide ions.

  • Recall that in our lesson on conjugate acids and bases, we learned that in a conjugate pair, the stronger the acid, the weaker the base. You will notice the effect of this in any chemistry information tables showing acid or base strength:
    • For the strongest acids, the dissociation into H+ and the conjugate base is shown by a single headed reaction arrow → instead of the equilibrium arrows. This is because the conjugate base of a strong acid is extremely poor at accepting a proton, so it will not go back to being the conjugate acid. This process won’t reach equilibrium like weak acids and bases will.
    • For the strongest bases, the addition of H+ to form the conjugate acid is also depicted by a single headed reaction arrow → because the reverse process of the poor conjugate acid returning to the original strong base simply will not occur.
    • This is true of CONJUGATE PAIRS, not of individual molecules; for example if compound A is a poor base, it doesn’t mean it is automatically a strong acid.

  • Be careful with polyprotic acids – those that have more than one proton to donate (e.g. H2SO4). Protons dissociate one at a time, and the second proton has a much lower degree of dissociation than the first proton. As far as sulfuric acid, a strong acid, is concerned, it has 100% dissociation of the first proton only:

    H2SO4 → H+ + HSO4-

    The effects of the second proton dissociating will be dealt with when looking at the acid dissociation constant, Ka.

  • As stated above, the degree of dissociation is how acidity and basicity is measured. With this however, once you have two strong acids/bases which both experience 100% dissociation, they are considered to have identical strength. This is because they are equivalent to solutions of H3O+ (aq) (or OH- for bases). All strong acids/bases have identical strength as far as the degree of dissociation goes because they all completely dissociate to H3O+ or OH- solutions.

  • One of the other differences between strong and weak acids and bases is in measurements like the enthalpy of neutralization. Remember that neutralization is the reaction:
    H+ (aq) + OH- (aq) → H2O (l)

    Enthalpy change of neutralization is the enthalpy change when an acid and base react in a neutralization to produce 1 mole of water. Strong acids have 100% dissociation into H+ (aq) and X- (aq), and strong bases will completely react to form OH- (aq) and B-H+ (aq). This means that all strong acids and bases have the same reaction to the same degree which is why their enthalpy of neutralization is a very similar exothermic value.
    In a weak acid however, not all the substance ionizes in solution (usually less than 1% does). Most of the acid isn’t ionized and there may be other enthalpy changes occurring rather than just H+ reacting with OH-. This leads to the enthalpy of neutralization being less exothermic in weak acids and bases than in strong acids and bases.

  • (AP) Analysing the structure of an acid or base helps to explain its relative strength. There are several structural factors to consider.
    Before looking at them, remember the definition of a strong acid/base is 100% dissociation. A strong acid will have a weak conjugate base that, due to these factors, tends to dissociate and not re-form the strong acid once it does dissociate.
    • Electronegative atoms: an electronegative atom can help to stabilize a conjugate base and make H atoms more likely to be lost as H+. Electronegativity will reduce the ability of lone pairs to attack and accept protons, so the conjugate base is weakened and in comparison, the acid is stronger.
      • An example is HCl. The H-Cl bond is quite polar, with Cl commanding most of the electron density and hydrogen being quite δ\delta+ or ‘acidic. Once dissociated, Cl- is a stable ion due to the electronegativity of Cl and it will not accept protons to re-form HCl. This contributes to HCl’s strength as an acid.

    • The inductive effect: this is where electron density is ‘pushed’ or drawn towards an atom in a chemical bond. If dissociation occurs, this can cause atoms to be more stable as ions (such as carbocations) than they would be without this effect, again stabilising them as conjugate bases.

    • Resonance: resonance is when a system delocalizes its electron density. Instead of different atoms of higher and lower charge density (think of mountains and valleys), it is more evenly spread out over the whole system (now think of these mountains/valleys flattened out).
      Resonance stabilizes molecules because electron density is less pronounced and available at any one atom. Less availability of electrons means less ability of a lone pair to attack a proton, so the conjugate base is less strong, more stable, and unlikely to change.
      • An example is the strong acid HNO3. Once it dissociates into H+ and NO3-, the three resonance forms of NO3- stabilize the structure and it does not re-form HNO3. Because it doesn’t “re-associate”, it stays fully dissociated and qualifies as a strong acid. This is also true of H2SO4 or sulfuric acid.

      Most species have a combination of these three effects that lead to their overall degree of dissociation and acid/base strength.
      • Phosphoric acid (H3PO4) is an example. It has resonance forms that stabilize its conjugate base (H2PO4-) just like the strong acids nitric acid (HNO3) and sulfuric acid (H2SO4). Unlike nitric and sulfuric acid though, it has only two resonance forms instead of three. In this way, we say it has less resonance stabilization.
        In addition to this, P is less electronegative than N and S. It does not pull electrons toward it with as much force as N or S, so the H atoms are less ‘acidic’, with a lower tendency to be lost as H+.
        In summary, compared to nitric and sulfuric acid, phosphoric acid has a less stable conjugate base and less tendency to dissociate in the first place. This is why phosphoric acid is a weak acid unlike its nitric and sulfuric analogues.