14.1 Brønsted-Lowry Acids and Bases
3 min read•june 25, 2024
Acids and bases are key players in chemistry, going beyond just sour and bitter tastes. The theory defines them by their ability to donate or accept protons, forming conjugate pairs in reactions. This framework helps us understand how substances interact and change in solution.
Water's unique properties make it central to acid-base chemistry. Its allows us to calculate ion concentrations and measure acidity using the . Some substances, like water itself, can act as both acids and bases, adapting to their chemical surroundings.
Brønsted-Lowry Acids and Bases
Brønsted-Lowry acids and bases
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- donates a () to another substance
- accepts a (H⁺) from another substance
- Examples of Brønsted-Lowry acids: (hydrochloric acid), (sulfuric acid), (acetic acid)
- Examples of Brønsted-Lowry bases: (ammonia), (), (carbonate ion)
- is determined by the extent to which an acid donates protons in solution
- is determined by the extent to which a base accepts protons in solution
Acid-base ionization equations
- differ by a single proton (H⁺) and are formed during acid-base reactions
- Conjugate acid is formed when a base gains a proton (H⁺), while conjugate base is formed when an acid loses a proton (H⁺)
- Example: HCl + H₂O ⇌ + (HCl and Cl⁻ are a conjugate acid-base pair, H₂O and H₃O⁺ are another conjugate acid-base pair)
- Acid ionization equation: + H₂O ⇌ H₃O⁺ + (HA is the acid, H₂O is the base, H₃O⁺ is the conjugate acid of water, A⁻ is the conjugate base of the acid)
- equation: + H₂O ⇌ + OH⁻ (B is the base, H₂O is the acid, BH⁺ is the conjugate acid of the base, OH⁻ is the conjugate base of water)
- occurs when an acid and base react to form water and a salt
Ion concentrations from water's ion-product
- Ion-product constant of water () equals the product of concentration [H₃O⁺] and hydroxide ion concentration [OH⁻] at a given temperature
- At 25°C, Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴
- In pure water at 25°C, [H₃O⁺] = [OH⁻] = 1.0 × 10⁻⁷ M (mol/L)
- Acidic solutions have [H₃O⁺] > 1.0 × 10⁻⁷ M and [OH⁻] < 1.0 × 10⁻⁷ M, while basic solutions have [H₃O⁺] < 1.0 × 10⁻⁷ M and [OH⁻] > 1.0 × 10⁻⁷ M
- Calculate unknown [H₃O⁺] or [OH⁻] using Kw and the known concentration of the other ion
- The pH scale is used to measure the acidity or basicity of a solution, with values ranging from 0 to 14
Amphiprotic substances in reactions
- substances can act as both an acid and a base in chemical reactions depending on the relative strengths of the other reactants
- Examples of substances: H₂O (water), (bicarbonate ion), (dihydrogen phosphate ion)
- Water (H₂O) acts as a base when reacting with an acid: HA + H₂O ⇌ H₃O⁺ + A⁻
- Water (H₂O) acts as an acid when reacting with a base: B + H₂O ⇌ BH⁺ + OH⁻
- Other amphiprotic substances behave similarly, with their role determined by the relative strengths of the acids and bases in the reaction
Lewis Acids and Bases
- Lewis acids are electron pair acceptors
- Lewis bases are electron pair donors
- This concept expands the definition of acids and bases beyond proton transfer
Key Terms to Review (40)
A⁻: A⁻ is the conjugate base of an acid, A, in the context of Brønsted-Lowry acid-base theory. It is the species formed when an acid loses a proton (H+) in a chemical reaction.
Acid strength: Acid strength refers to the ability of an acid to donate protons (H+) in a chemical reaction. It is a measure of how readily an acid can dissociate in water to release these protons, which is crucial for understanding the behavior of acids in various chemical processes. Strong acids completely ionize in solution, while weak acids only partially dissociate, highlighting the importance of this concept in predicting the reactivity and properties of different acids.
Alpha (α) decay: Alpha (α) decay is a type of radioactive decay where an unstable nucleus emits an alpha particle, consisting of 2 protons and 2 neutrons. This process reduces the atomic number by 2 and the mass number by 4.
Amphiprotic: An amphiprotic substance can act as both a proton donor (acid) and a proton acceptor (base) depending on the chemical environment. Water is a common example of an amphiprotic substance.
Amphiprotic: Amphiprotic, also known as amphoteric, refers to a substance that can act as both an acid and a base in chemical reactions. These substances are able to donate and accept protons (H+ ions) depending on the pH of the surrounding environment.
Amphoteric: An amphoteric substance can act as both an acid and a base. This dual behavior depends on the chemical environment it encounters.
Autoionization: Autoionization is the process by which molecules of a pure substance react with each other to form ions. A common example is the autoionization of water, where two water molecules produce hydronium and hydroxide ions.
B: B is a fundamental concept that spans multiple topics in chemistry, including atomic structure, acid-base chemistry, and the relative strengths of acids and bases. It is a versatile and essential term that underpins our understanding of various chemical phenomena.
Base ionization: Base ionization is the process in which a base accepts protons from water molecules, producing hydroxide ions ($OH^-$) and the conjugate acid of the base. This reaction occurs in aqueous solutions.
Base Strength: Base strength refers to the ability of a Brønsted-Lowry base to accept protons (H+ ions) and form a conjugate acid. The strength of a base is determined by its propensity to attract and stabilize protons, which is directly related to its ability to donate electrons and form covalent bonds.
BH⁺: BH⁺ is the conjugate acid of a Brønsted-Lowry base, B. It is formed when a Brønsted-Lowry base accepts a proton (H⁺) from an acid, resulting in the creation of the conjugate acid, BH⁺. This process is a fundamental concept in the Brønsted-Lowry theory of acids and bases.
Boyle: Boyle's Law states that the pressure of a gas is inversely proportional to its volume when temperature and the number of moles are held constant. It can be mathematically expressed as $PV = k$, where $P$ is pressure, $V$ is volume, and $k$ is a constant.
Brønsted-Lowry: Brønsted-Lowry is a theory of acid-base behavior that defines acids as proton (H+) donors and bases as proton acceptors. This theory provides a more comprehensive understanding of acid-base reactions compared to the earlier Arrhenius definition, which was limited to aqueous solutions.
Brønsted-Lowry acid: A Brønsted-Lowry acid is a substance that can donate a proton ($H^+$) to another substance. It is one-half of the Brønsted-Lowry acid-base theory, which focuses on proton transfer reactions.
Brønsted-Lowry base: A Brønsted-Lowry base is a substance that can accept a proton (hydrogen ion, $H^+$) from another substance. It plays a crucial role in acid-base reactions by undergoing protonation.
CH₃COOH: CH₃COOH, also known as acetic acid, is a weak acid that plays a crucial role in the context of Brønsted-Lowry acid-base theory. It is a key compound in various chemical reactions and has widespread applications in both scientific and everyday settings.
Cl⁻: Cl⁻ is the chloride ion, a negatively charged species formed when the element chlorine (Cl) gains an electron. It is an important ion in various chemical and biological processes, particularly in the context of Brønsted-Lowry acid-base theory.
CO₃²⁻: CO₃²⁻ is the carbonate ion, a polyatomic ion with a charge of -2. It is an important species in the context of Brønsted-Lowry acid-base theory, as it can act as both a base and a conjugate base.
Conjugate Acid-Base Pairs: Conjugate acid-base pairs are related chemical species that differ by the presence or absence of a single proton (H+). When an acid donates a proton, it becomes a conjugate base, and when a base accepts a proton, it becomes a conjugate acid. These pairs are fundamental to understanding the Brønsted-Lowry theory of acids and bases, as well as the concepts of pH, relative acid-base strengths, hydrolysis, polyprotic acids, and acid-base titrations.
Davy: Sir Humphry Davy was an English chemist and inventor known for his contributions to electrochemistry and the discovery of several alkali and alkaline earth metals. His work laid foundational principles in the study of acids, bases, and their reactions.
H⁺: H⁺ is the symbol used to represent a hydrogen ion, which is a proton (a positively charged subatomic particle) that has lost its electron. It is a crucial concept in the Brønsted-Lowry theory of acids and bases, as the transfer of H⁺ ions is central to acid-base reactions.
H₂PO₄⁻: H₂PO₄⁻ is the dihydrogen phosphate ion, a Brønsted-Lowry conjugate base that forms when a hydrogen ion (H⁺) is removed from phosphoric acid (H₃PO₄). It is an important species in acid-base equilibria and plays a key role in biological buffer systems.
H₂SO₄: H₂SO₄, also known as sulfuric acid, is a strong, corrosive, and highly reactive chemical compound that plays a crucial role in the context of Brønsted-Lowry acid-base theory. It is a diprotic acid, meaning it can donate two protons (H⁺) in aqueous solutions.
H₃O⁺: H₃O⁺, also known as the hydronium ion, is a polyatomic ion with the chemical formula H₃O⁺. It is the conjugate acid of water and is a key concept in the understanding of Brønsted-Lowry acid-base theory.
HA: HA, or Brønsted-Lowry acid, is a chemical species that can donate a proton (H+) to another substance, thereby acting as an acid in a chemical reaction. This term is particularly relevant in the context of understanding Brønsted-Lowry acid-base theory and the relative strengths of acids and bases.
HCl: HCl, or hydrogen chloride, is a chemical compound consisting of one hydrogen atom and one chlorine atom. It is a colorless, corrosive gas that has a wide range of applications and plays a crucial role in various chemical processes and reactions.
HCO₃⁻: HCO₃⁻ is the bicarbonate ion, a polyatomic ion with the chemical formula HCO₃⁻. It is an important species in acid-base chemistry and plays a crucial role in the regulation of pH in the body.
Hydronium Ion: The hydronium ion, represented as H3O+, is a positively charged ion formed when a proton (H+) combines with a water molecule (H2O). It is a key concept in understanding Brønsted-Lowry acid-base theory and the behavior of buffers in aqueous solutions.
Hydroxide Ion: The hydroxide ion, represented by the chemical formula OH-, is a negatively charged ion consisting of one oxygen atom and one hydrogen atom. It is a key player in the concept of Brønsted-Lowry acids and bases, where it acts as a base by accepting a proton (H+) to form water.
Ion-Product Constant: The ion-product constant, also known as the equilibrium constant for the autoionization of water, is a fundamental concept in the study of Brønsted-Lowry acids and bases. It represents the equilibrium constant for the self-ionization of water, which is a key process in determining the pH of aqueous solutions.
Ion-product constant for water, Kw: The ion-product constant for water, $K_w$, is the equilibrium constant for the self-ionization of water. It is defined as the product of the molar concentrations of hydrogen ions and hydroxide ions in water at a given temperature.
Kw: Kw, or the equilibrium constant for water, is a fundamental concept in chemistry that describes the self-ionization of water and its relationship to the acidity or basicity of a solution. This term is crucial in understanding Brønsted-Lowry acid-base theory, pH and pOH calculations, as well as the relative strengths of acids and bases.
Lewis Acids and Bases: Lewis acids are electron-pair acceptors, while Lewis bases are electron-pair donors. They are defined based on the ability to accept or donate electron pairs, rather than the ability to donate or accept protons as in the Brønsted-Lowry definition of acids and bases.
Neutralization: Neutralization is a chemical process in which an acid and a base react to form a salt and water. This process is characterized by the mutual cancellation or counteraction of the acidic and basic properties, resulting in a neutral solution.
Neutralization reaction: A neutralization reaction is a chemical reaction in which an acid and a base react to form water and a salt. This type of reaction typically involves the combination of hydrogen ions ($H^+$) from the acid and hydroxide ions ($OH^-$) from the base.
NH₃: NH₃, also known as ammonia, is a chemical compound composed of one nitrogen atom and three hydrogen atoms. It is a colorless gas with a pungent odor and is an important compound in the context of Brønsted-Lowry acid-base theory.
OH⁻: OH⁻ is the hydroxide ion, a negatively charged species consisting of one oxygen atom and one hydrogen atom. It is an important concept in the context of Brønsted-Lowry acid-base theory, as it acts as a base by accepting protons to form water.
PH Scale: The pH scale is a measure of the acidity or basicity of a solution, ranging from 0 to 14. It is a logarithmic scale that quantifies the concentration of hydrogen ions (H+) in a solution, with lower values indicating higher acidity and higher values indicating higher basicity or alkalinity. The pH scale is a fundamental concept in understanding Brønsted-Lowry acids and bases, as well as the relationship between pH and pOH, and the behavior of polyprotic acids.
Proton: A proton is a subatomic particle found in the nucleus of an atom, carrying a positive electric charge. Protons contribute to the atomic number and define the element.
Proton: A proton is a subatomic particle that carries a positive electric charge and is found in the nucleus of an atom. Protons are fundamental to the structure and behavior of atoms, and they play crucial roles in various areas of chemistry, including the evolution of atomic theory, acid-base chemistry, the properties of hydrogen, and nuclear physics.