Review of Chemical Reactions and Stoichiometry

Precipitation Reactions and Ionic Equations

Chemical reactions in aqueous solutions often involve the formation of a precipitate, which is an insoluble solid formed from the reaction of two soluble salts. Writing balanced equations for these reactions is essential for understanding stoichiometry and the conservation of mass.

• Complete Ionic Equation: Shows all soluble ionic substances dissociated into ions.

• Net Ionic Equation: Includes only the ions and compounds directly involved in the chemical change.

• Example: When magnesium chloride reacts with potassium carbonate, magnesium carbonate precipitates:

Net Ionic Equation:

• Redox Check: This is not a redox reaction; the oxidation numbers of all elements remain unchanged.

• Stoichiometry Application: The amount of precipitate can be used to determine the concentration of ions in solution.

• Calculation Example: If 21 g of MgCO3 (molar mass 84.3 g/mol) forms from 0.5 L solution, then:

Number of moles:

Molarity:

Oxidation Numbers

Assigning oxidation numbers helps identify redox reactions and track electron transfer. Key rules include:

• Elements in their standard state have oxidation number 0.

• Group 1 metals: +1; Group 2 metals: +2; Oxygen: usually -2; Hydrogen: +1 (with nonmetals), -1 (with metals).

• The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion's charge.

Examples:

• Br in BrO3-: +5

• C in K2CO3: +4

• N in NH3: -3

• O in O2 and O3: 0

• Mn in MnO4-: +7

• P in P2O5: +5

• N in NO: +2

• Ti in CaTiO3: +4

• H in CaH2: -1

Thermochemistry: Energy, Heat, and Work

Nature of Energy

Thermochemistry studies the relationships between chemical reactions and energy changes, especially heat. Energy is the capacity to do work or transfer heat.

• Work (w): Energy transfer not as heat; force acting over a distance.

• Heat (q): Energy transfer due to temperature difference.

Forms of Energy

• Kinetic Energy (Ek): Energy of motion.

• Potential Energy: Energy due to position or composition. In chemistry, electrostatic potential energy is most important:

• Thermal Energy: Energy associated with temperature, arising from molecular motion (a form of kinetic energy).

Units of Energy

• Joule (J): SI unit of energy.

• Calorie (cal): (food Calorie = 1 kcal = 1000 cal)

System and Surroundings

The system is the part of the universe under study (e.g., reactants and products), while the surroundings are everything else (e.g., container, air).

Work and Heat

• Work: (force times distance)

• Heat: Energy transferred due to temperature difference; always flows from hot to cold objects.

Conversion of Energy

Energy can be converted between forms (e.g., potential to kinetic, kinetic to thermal). Friction is a common way kinetic energy is converted to heat.

First Law of Thermodynamics and Internal Energy

First Law of Thermodynamics

The First Law of Thermodynamics states that energy is neither created nor destroyed; the total energy of the universe is constant. Energy lost by the system is gained by the surroundings, and vice versa.

Internal Energy (E)

Internal energy is the sum of all kinetic and potential energies of all particles in a system. It is a state function, meaning it depends only on the current state, not the path taken to reach it.

• Change in Internal Energy:

• Energy Exchange:

Endergonic and Exergonic Processes

• Endergonic: (system absorbs energy)

• Exergonic: (system releases energy)

Heat (q) and Work (w) Are Not State Functions

Although internal energy is a state function, heat and work are not; their values depend on the process path.

Sign Conventions for q, w, and ΔE

Enthalpy and Heat Flow at Constant Pressure

Definition of Enthalpy (H)

Enthalpy is a thermodynamic quantity defined as the internal energy plus the product of pressure and volume:

• Most chemical reactions occur at constant pressure, so enthalpy is useful for tracking heat flow.

Change in Enthalpy (ΔH)

The change in enthalpy at constant pressure equals the heat gained or lost by the system:

Since and , then at constant pressure.

Endothermic and Exothermic Processes

• Endothermic: (system absorbs heat from surroundings)

• Exothermic: (system releases heat to surroundings)

Summary Table: Key Thermodynamic Quantities