General Chemistry: Reactions, Stoichiometry, and Thermochemistry

Comprehensive practice flashcards featuring fundamental vocabulary and principles across chemical equations, stoichiometry, solution chemistry, acids, bases, and thermochemistry.

Chemical Equation

The shorthand notation of a chemical reaction using chemical formulas, symbols, and physical state indicators ($(s)$, $(l)$, $(g)$, $(aq)$) to represent the chemical changes occurring in a process.

Reactants

The initial substances consumed during a chemical reaction, positioned on the left side of a chemical equation.

Products

The new chemical substances formed as a result of a chemical reaction, positioned on the right side of a chemical equation.

Coefficients

The whole numbers placed directly in front of chemical formulas in an equation to denote the relative number of molecules, formula units, or moles of each substance involved.

Law of Conservation of Matter

A fundamental scientific law stating that matter is neither created nor destroyed during a chemical reaction; the total mass of the products must equal the total mass of the reactants, and total atom counts of each element remain unchanged.

Combustion Reaction

A rapid chemical reaction between a substance and oxygen gas ($\text{O}_2$) that typically releases heat and light; combustion of hydrocarbons produces carbon dioxide ($\text{CO}_2$) gas and water ($\text{H}_2\text{O}$) vapor.

Stoichiometry

The quantitative study and calculation of the numerical relationships between amounts of reactants and products involved in a balanced chemical reaction.

Mole Ratio

A stoichiometric conversion factor derived from the coefficients of a balanced chemical equation that establishes the proportional mole relationship between any two chemical species in the reaction.

Limiting Reactant

The reactant that is completely consumed first in a chemical reaction, thereby limiting and determining the maximum theoretical yield of product that can be generated.

Excess Reactant

A reactant present in a quantity greater than required to react with the limiting reactant, leaving an unreacted portion remaining after the reaction ceases.

Theoretical Yield

The maximum quantity of product that can be generated in a chemical reaction, calculated from the complete consumption of the limiting reactant based on the balanced equation.

Actual Yield

The amount of product isolated, collected, and measured experimentally in the laboratory from a given chemical reaction.

Percent Yield

The percentage efficiency of a chemical reaction, defined as: $\text{\% Yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$.

Solution

A homogeneous mixture of two or more substances that is uniform in composition and appearance throughout.

Solute

The substance dissolved in a solution, typically present in a smaller quantity than the solvent.

Solvent

The substance in which the solute dissolves, present in the greatest quantity; when water is the solvent, the resulting mixture is termed an aqueous solution ($aq$).

Like Dissolves Like

The empirical principle stating that substances with similar intermolecular forces dissolve in one another; polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents.

Ion-Dipole Forces

The electrostatic attractive forces that occur between dissolved ions and the partial electrical charges on the ends of polar solvent molecules (dipoles), facilitating the dissolution of ionic compounds in water.

Electrolyte

A substance that dissolves in water to form mobile ions, creating an aqueous solution capable of conducting electric current; strong electrolytes dissociate completely, whereas weak electrolytes dissociate only partially.

Colligative Properties

Physical properties of solutions (such as freezing point depression and boiling point elevation) that depend strictly on the total number of dissolved solute particles, regardless of their chemical identities.

Solubility

The maximum amount of solute that can dissolve in a specified quantity of solvent at a given temperature; substances forming solutions with concentrations greater than $0.1\,M$ are practically classified as soluble.

Precipitation Reaction

A reaction in aqueous solution wherein two soluble ionic compounds exchange ionic partners to form at least one insoluble solid product known as a precipitate.

Molecular Equation

A chemical equation written with all reactants and products represented by their complete, undissociated neutral chemical formulas.

Complete Ionic Equation

A chemical equation displaying all soluble ionic compounds and strong electrolytes dissociated into their constituent aqueous ions, leaving insoluble solids, liquids, and gases in their intact molecular/formula forms.

Spectator Ions

Ions in an aqueous reaction mixture that do not participate directly in the chemical change, appearing unchanged on both the reactant and product sides of a complete ionic equation.

Net Ionic Equation

A simplified chemical equation resulting from removing all spectator ions from a complete ionic equation, leaving only the chemical species that directly participate in the reaction.

Molarity ($M$)

A standard unit of solution concentration defined as the number of moles of solute per liter of total solution: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$.

Dilution

The process of preparing a solution of lower concentration from a more concentrated stock solution by adding solvent, described quantitatively by the relationship: $M_{\text{concentrated}} \times V_{\text{concentrated}} = M_{\text{dilute}} \times V_{\text{dilute}}$.

Arrhenius Acid

A compound that ionizes in aqueous solution to increase the concentration of hydrogen ions ($\text{H}^+$) or hydronium ions ($\text{H}_3\text{O}^+$).

Arrhenius Base

A compound that dissociates in aqueous solution to yield hydroxide ions ($\text{OH}^-$).

Brønsted-Lowry Acid

A proton donor; any substance that donates a hydrogen ion ($\text{H}^+$) in a chemical reaction.

Brønsted-Lowry Base

A proton acceptor; any substance that accepts a hydrogen ion ($\text{H}^+$) in a chemical reaction.

pH

A dimensionless scale measuring aqueous acidity, defined as the negative base-$10$ logarithm of the hydronium ion concentration: $\text{pH} = -\log[\text{H}_3\text{O}^+]$.

Neutralization Reaction

A chemical reaction between an acid and a base yielding water and an ionic salt, characterized by the net ionic equation: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$.

Thermodynamics

The scientific discipline focused on the study of energy, work, heat, and their mutual interconversions in physical and chemical systems.

Thermal Energy

A form of kinetic energy that originates from the constant, random movement and vibration of atoms and molecules within a substance, measured macroscopically as temperature.

Chemical Energy

A form of potential energy stored within the spatial arrangement, structural attachments, and chemical bonds holding atoms and molecules together.

First Law of Thermodynamics

The fundamental law stating that energy cannot be created or destroyed; it can only be converted between forms, maintaining a constant total energy across the universe ($\Delta E_{\text{universe}} = 0$).

Kinetic Energy ($KE$)

The energy an object possesses due to its motion, proportional to its mass and the square of its velocity: $E_k = \frac{1}{2} m v^2$.

Potential Energy ($PE$)

The stored energy an object possesses by virtue of its position, structural composition, or state relative to other interacting bodies.

Joule ($\text{J}$)

The SI unit of energy, defined as the work required to exert a force of $1\,\text{N}$ over a distance of $1\,\text{m}$ ($1\,\text{J} = 1\,\text{kg}\cdot\text{m}^2/\text{s}^2$).

Calorie ($\text{cal}$)

A non-SI energy unit defined as the heat required to raise the temperature of exactly $1\,\text{g}$ of pure liquid water by $1\,^\circ\text{C}$, precisely equivalent to $4.184\,\text{J}$.

System

The specific part of the universe singled out for observation and thermodynamic study (such as the chemical reactants and products of a reaction).

Surroundings

Everything outside the defined thermodynamic system with which the system can exchange energy or matter.

Open System

A thermodynamic system capable of freely exchanging both energy and matter across its boundary with the surroundings.

Closed System

A thermodynamic system that can exchange energy (as heat or work) with its surroundings, but cannot exchange matter.

Isolated System

A thermodynamic system that is completely insulated from its environment, exchanging neither energy nor matter with its surroundings.

Internal Energy ($E$)

The total sum of all microscopic kinetic and potential energies possessed by all the particles that constitute a thermodynamic system.

State Function

A mathematical property of a system whose value depends exclusively on the current state (initial and final conditions) of the system, completely independent of the path or mechanism taken to reach it.

Work ($w$)

An energy transfer mechanism involving a force acting across a distance ($w = F \times d$); for expanding gases at constant external pressure, pressure-volume work is given by: $w = -P\Delta V$.

Heat ($q$)

The thermal energy transferred between a system and its surroundings as a consequence of a temperature gradient, flowing spontaneously from higher temperature to lower temperature.

Thermal Equilibrium

The thermal state achieved when two bodies in contact reach identical temperatures, resulting in zero net transfer of heat energy between them.

Specific Heat Capacity ($C_s$)

An intrinsic physical constant representing the quantity of heat energy required to raise the temperature of $1\,\text{g}$ of a substance by $1\,\text{K}$ (or $1\,^\circ\text{C}$), expressed in units of $\text{J}/(\text{g}\cdot^\circ\text{C})$.

Enthalpy ($H$)

A thermodynamic state function defined as the sum of a system's internal energy plus the product of its pressure and volume: $H = E + PV$.

Change in Enthalpy ($\Delta H$)

The quantity of heat evolved or absorbed by a chemical process at constant pressure: $\Delta H = q_p$.

Exothermic Reaction

A reaction or process that releases thermal energy to its surroundings, characterized by a negative enthalpy change ($\Delta H < 0$) and products with lower chemical potential energy than reactants.

Endothermic Reaction

A reaction or process that absorbs thermal energy from its surroundings, characterized by a positive enthalpy change ($\Delta H > 0$) and products with higher chemical potential energy than reactants.

Hess's Law

A thermodynamic principle stating that if a chemical reaction can be broken down into a series of steps, the overall enthalpy change ($\Delta H_{\text{rxn}}$) equals the sum of the enthalpy changes of the individual steps.

Standard Enthalpy of Formation ($\Delta H_f^\circ$)

The enthalpy change occurring when exactly one mole of a compound is synthesized from its constituent elements in their standard states at $1.00\,\text{atm}$ and $298\,\text{K}$; elements in standard state have $\Delta H_f^\circ = 0$.

Bond Energy

The energy needed to break one mole of a particular chemical bond in a gaseous molecule; breaking bonds requires energy (endothermic), while making bonds releases energy (exothermic).