World Geography and Science Quiz

Proton transfer reactions

• Hydrogen atoms consist of one proton and one electron.
• Ionization involves losing an electron, leaving only the proton (H⁺).
• Proton transfer reactions occur when a proton is transferred from one reactant to another.
• Reactants that lose H⁺ are proton donors, and are Brønsted-Lowry acids.
• Reactants that accept H⁺ are proton acceptors, and are Brønsted-Lowry bases.
• Species differing by only a single proton are conjugate acid-base pairs.

Brønsted-Lowry acids and bases

• A Brønsted-Lowry acid is a proton donor.
• A Brønsted-Lowry base is a proton acceptor.

Acid-Base Theories Historically

• Lavoisier initially defined acids as compounds containing oxygen. This was later proven incorrect.
• Arrhenius defined acids as substances that release H+ ions in water, and bases as substances that release OH- ions in water. This definition is limited to aqueous solutions only.
• Brønsted and Lowry's theory generalizes the definition to include reactions in any solvent

Evolution of theories

• Definitions of acids and bases have evolved over time as more knowledge has been acquired and new phenomena encountered
• Early definitions were limited to specific cases. Modern definitions apply more generally.

Proton transfer reactions in terms of Brønsted-Lowry theory

• An acid can only behave as a proton donor if it is also a base present to accept the proton
• Acid reacts to form base and vice versa
• A conjugate acid-base pair differ by one proton

Reacting species in Brønsted-Lowry acid-base reactions form conjugate pairs

• Conjugate acid-base pairs differ by one proton
• The act of donating cannot happen in isolation
• An acid can only behave as a proton donor if there is also a base present to accept the proton

Deducing the formulas of conjugate acids and bases

• In a conjugate pair the acid always has one proton more than its conjugate base.
• Make it easy to predict the formula of the corresponding conjugate for any given acid or base

The pH scale

• pH is a logarithmic scale used to express the concentration of hydrogen ions in a solution.
• pH = -log₁₀[H⁺]
• A lower pH indicates a higher concentration of hydrogen ions (more acidic).
• A higher pH indicates a lower concentration of hydrogen ions (more basic).

The ion product constant of water

• Water undergoes self-ionization to a small extent: H₂O(l) ⇌ H⁺(aq) + OH⁻(aq).
• The equilibrium constant for this reaction is called the ion product constant Kw.
• The product of [H⁺] and [OH⁻] is constant at a given temperature.

Strong and weak acids and bases

• Strong acids and strong bases dissociate almost completely in solution, generating a high concentration of ions.
• Weak acids and weak bases dissociate only partially in solution, generating a low concentration of ions.
• Strong acids and bases have large equilibrium constants (Ka or Kb).
• Weak acids and bases have small equilibrium expressions.

Neutralization reactions

• Neutralization reactions involve the transfer of protons between an acid and a base, producing water and a salt.
• Acidic and basic solutions have similar behaviors, with notable differences in the overall composition of the solution and their behaviors when reacting with other chemicals.
• The reaction between H⁺ ions and OH⁻ ions to form water is a neutralization reaction.
• This reaction is exothermic.

pH curves

• pH curves of neutralization reactions exhibit characteristic features.
• They can be used to determine the equivalence point, buffer range, and pKa.

Acid-base Indicators

• Indicators are weak acids or bases where the undissociated and dissociated forms have different colors.
• By changing the pH a noticeable change in color occurs.
• The end point of an indicator is the pH range.
• The equivalence point is the pH point where reaction is stoichiometrically balanced.

Buffer solutions

• Buffer solutions resist changes in pH upon adding small amounts of acid or base.
• Buffer solutions contain a weak acid and its conjugate base, or a weak base and its conjugate acid.
• The pH of the buffer solution is determined by the pKa/pKb and the ratio of the concentrations of the conjugate pair.
• Buffers are essential for many biological systems as well as for many industrial processes for maintaining a stable environment.

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