Acids and Bases

Definitions of Acids and Bases

Acids and bases are fundamental chemical species with several definitions, each broadening the scope of what constitutes an acid or base.

• Arrhenius Definition: An acid produces H+ ions in aqueous solution; a base produces OH- ions.

• Brønsted-Lowry Definition: An acid is a proton (H+) donor; a base is a proton acceptor.

• Lewis Definition: An acid is an electron pair acceptor; a base is an electron pair donor.

Example: NH3 is a Brønsted-Lowry base (accepts H+) and a Lewis base (donates electron pair).

Conjugate Acids and Bases

Acids and bases form conjugate pairs during reactions.

• Conjugate Acid: Formed when a base gains a proton.

• Conjugate Base: Formed when an acid loses a proton.

• Identifying Conjugate Pairs: In a reaction, each acid/base pair differs by one H+.

Example: In the reaction: , NH3 and NH4+ are a conjugate pair.

Acid/Base Reactions and Amphoteric Substances

Acid/base reactions involve proton transfer. Amphoteric substances can act as both acids and bases.

• Amphoteric: Substances like H2O can donate or accept H+.

• Example: Water acts as an acid with NH3 and as a base with HCl.

Acid Strength and Molecular Structure

The strength of an acid depends on its molecular structure.

• Binary Acids (HX): Strength increases with bond polarity and decreases with bond strength.

• Oxyacids: For same central atom, more oxygens increase acid strength. For same number of oxygens, higher electronegativity of central atom increases strength.

Example: HClO4 > HClO3 > HClO2 > HClO

Acid/Base Reaction Favorability

Reactions favor the side with the weaker acid/base (lower tendency to donate/accept protons).

pH, pOH, and Ion Concentrations

pH and pOH quantify acidity and basicity.

• pH Equation:

• pOH Equation:

• Relationship: (at 25°C)

• Water Ion Product:

Example: If , then .

Strong vs. Weak Acids and Bases

Strong acids/bases dissociate completely; weak acids/bases only partially.

• Strong Acid: e.g., HCl, HNO3

• Weak Acid: e.g., CH3COOH

• Calculations: For strong acids/bases, or equals initial concentration.

• For weak acids/bases: Use or and ICE tables to solve for equilibrium concentrations.

Example: For 0.1 M HCl, .

Polyprotic Acids

Polyprotic acids can donate more than one proton.

• Stepwise Dissociation: Each step has its own value, with .

• Example: H2SO4 dissociates in two steps.

Stepwise Equations:

Ka and Kb Relationship

For conjugate acid-base pairs:

Hydrolysis of Ions and Salt Solutions

Some ions react with water (hydrolyze), affecting solution pH.

• Acidic, Basic, or Neutral Solutions: Determined by the nature of the ions from the salt.

• Example: NaCl yields neutral solution; NH4Cl yields acidic solution.

Buffers and the Common Ion Effect

Buffer Solutions

Buffers resist changes in pH when small amounts of acid or base are added.

• Made from: Weak acid and its conjugate base, or weak base and its conjugate acid.

• Example: CH3COOH/CH3COO-

• Buffer Capacity: Amount of acid/base a buffer can neutralize before pH changes significantly.

Buffer Equation:

Common Ion Effect

The addition of a common ion suppresses the ionization of a weak acid/base, affecting pH.

Buffer Calculations

• Use Henderson-Hasselbalch equation to calculate pH.

• Determine quantities needed for desired pH.

Titration of Acids and Bases

Titration is used to determine unknown concentrations or molar masses.

• Equivalence Point: Point at which stoichiometric amounts of acid and base have reacted.

• Half-Equivalence Point: pH = pKa for weak acid/strong base titration.

• Indicator Selection: Choose indicator whose color change matches equivalence point pH.

Solubility and Precipitation

Solubility and Ksp

Solubility product constant () quantifies the equilibrium between a solid and its ions in solution.

• Molar Solubility: Amount of solute that dissolves to form a saturated solution.

• Calculations: Use to find ion concentrations.

Predicting Precipitation

Mixing solutions may form a precipitate if the product of ion concentrations exceeds .

Effect of Additives on Solubility

• Adding common ions decreases solubility.

• Adding acids may increase solubility of salts containing basic anions.

Thermodynamics

Spontaneous and Non-Spontaneous Processes

Spontaneous processes occur without external input; non-spontaneous require energy.

• Example: Ice melting at room temperature is spontaneous.

Entropy and Disorder

Entropy () measures disorder/randomness. Greater disorder favors spontaneity.

• Phase Changes: Entropy increases from solid to liquid to gas.

• Chemical Reactions: Entropy increases if products are more disordered than reactants.

Second and Third Laws of Thermodynamics

• Second Law: Total entropy of the universe increases in spontaneous processes.

• Third Law: Entropy of a perfect crystal at absolute zero is zero.

Microstates

A microstate is a specific arrangement of particles; more microstates mean higher entropy.

Temperature Dependence of Entropy

• and depend on temperature; higher temperature means smaller $\mathrm{\Delta S_{surr}}$ for a given heat transfer.

Calculating Entropy and Free Energy

• Standard Entropy Change:

• Gibbs Free Energy:

• Spontaneity: : spontaneous; : non-spontaneous.

• Standard Free Energy Change:

• Nonstandard Conditions:

• Relationship with Equilibrium:

Electrochemistry

Redox Reactions

Redox reactions involve electron transfer.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation; is reduced.

• Reducing Agent: Causes reduction; is oxidized.

Example: Zn + Cu2+ → Zn2+ + Cu

Assigning Oxidation Numbers

• Rules for assigning oxidation numbers to atoms in compounds and ions.

Balancing Redox Reactions

• Balance in acidic or basic solution using half-reaction method.

Electrochemical Cells and Cell Potential

Electrochemical cells convert chemical energy to electrical energy.

• Cell Potential (Eocell): Calculated from standard reduction potentials.

• Hydrogen Electrode: Eo = 0.00 V by convention.

• Spontaneity: Eocell > 0: spontaneous.

Drawing and Labeling Electrochemical Cells

• Label anode, cathode, salt bridge, and direction of electron flow.

Predicting Redox Behavior

• Use reduction potential tables to predict which substances are oxidized/reduced.

• Predict metal reactivity with acids.

Relationships Between Eocell, Keq, and ΔG

Nernst Equation

Electrolysis

• Predict products for electrolysis of molten and aqueous salts.

Cathodic Protection and Sacrificial Anodes

• Use more easily oxidized metals as sacrificial anodes to protect structures from corrosion.

Organic Chemistry Fundamentals

Lewis Structures and Bond-Line Formulas

Organic compounds are represented by Lewis structures and bond-line (carbon skeleton) formulas.

• Lewis Structures: Show all atoms, bonds, and lone pairs.

• Bond-Line Formulas: Simplified representation showing carbon skeleton.

Hybridization in Organic Molecules

• Determine hybridization (sp, sp2, sp3) based on bonding and geometry.

Formal Charges of Reactive Intermediates

• Identify formal charges for carbocations (positive), carbanions (negative), oxonium, iminium, etc.

Summary Table: Acid/Base Definitions

Summary Table: Strong vs. Weak Acids and Bases

Summary Table: Thermodynamic Laws

Summary Table: Redox Terms

Additional info: Academic context and examples have been added to clarify and expand upon the original bullet points.