General Chemistry: Mass Relationships and Aqueous Reactions

Comprehensive vocabulary flashcards covering stoichiometry, chemical reactions, balancing equations, solutions, electrolyte classifications, solubility rules, precipitation, acid-base neutralization, redox reactions, oxidation numbers, molarity, dilution, gravimetric analysis, and titrations.

Atomic Mass

The mass of an atom expressed in atomic mass units ($\text{amu}$), also historically referred to as atomic weight.

Atomic Mass Unit ($\text{amu}$)

A mass unit defined as exactly equal to one-twelfth ($\frac{1}{12}$) the mass of one carbon-12 atom.

Carbon-12 Standard

The carbon isotope containing six protons and six neutrons, assigned an atomic mass of exactly $12.00\,\text{amu}$ to serve as the reference standard for all other elements.

Hydrogen Atomic Mass Basis

An experimentally determined value of $1.008\,\text{amu}$, calculated because a hydrogen atom has on average $8.400\%$ of the mass of a carbon-12 atom ($0.084 \times 12.00\,\text{amu}$).

Mole ($\text{mol}$)

The SI unit for the amount of substance containing as many elementary entities as there are atoms in exactly $12\,\text{g}$ (or $0.012\,\text{kg}$) of carbon-12.

Molar Mass ($\mathcal{M}$)

The mass (in grams or kilograms) of one mole of units (such as atoms or molecules) of a substance, expressed in $g\,mol^{-1}$.

Molar Mass of Carbon

$12.01\,g\,mol^{-1}$, representing the mass of one mole ($6.022 \times 10^{23}$ atoms) of naturally occurring carbon.

Molar Mass of Molecular Oxygen ($O_2$)

$32.00\,g\,mol^{-1}$, corresponding to the mass of one mole ($6.022 \times 10^{23}$ molecules) of diatomic oxygen gas.

Molar Mass of Water ($H_2O$)

$18.02\,g\,mol^{-1}$, representing the mass of one mole ($6.022 \times 10^{23}$ molecules) of water.

Atomic Mass vs. Molar Mass: Scale

Atomic mass is measured on a microscopic scale (single atom in $\text{amu}$), whereas molar mass is measured on a macroscopic scale (laboratory amounts in $g\,mol^{-1}$).

Average Atomic Mass

The weighted average of the atomic masses of all naturally occurring isotopes of an element, calculated using their fractional abundances: $\sum (\text{isotopic mass} \times \text{fractional abundance})$.

Fractional Abundance

The percent abundance of an isotope divided by $100\%$, used as a weighting factor in average atomic mass calculations.

Average Atomic Mass of Copper

$63.55\,\text{amu}$, calculated from its two stable isotopes: $^{63}_{29}\text{Cu}$ ($62.93\,\text{amu}$, $69.09\%$\text{ abundance}) and $^{65}_{29}\text{Cu}$ ($64.9278\,\text{amu}$, $30.91\%$\text{ abundance}).

Avogadro's Number ($N_a$)

The constant $6.022 \times 10^{23}$, indicating the total number of elementary particles present in exactly one mole of any chemical substance.

Mass of One Carbon-12 Atom

$1.993 \times 10^{-23}\,\text{g}$, obtained by dividing the molar mass ($12.00\,\text{g}$) by Avogadro's number ($6.022 \times 10^{23}$ atoms).

Mass Spectrometer

A scientific instrument that determines the masses of atoms or molecules and identifies substances by separating gas-phase ions based on their mass-to-charge ratio ($m/z$).

Ion Source

The component of a mass spectrometer where gaseous atoms or molecules are converted into positively charged ions.

Accelerating Plates

Parallel plates in a mass spectrometer that produce an electric field ($\Delta V$) to accelerate ions into a beam of uniform velocity.

Magnetic Field ($B$) in Mass Spectrometry

The magnetic region that bends the trajectory of moving charged ions into circular paths with radii depending on their mass-to-charge ratio ($m/z$).

Mass Spectrum of Neon

A plot of peak intensity versus atomic mass showing three isotopes: $^{20}_{10}\text{Ne}$ ($90.92\%$), $^{21}_{10}\text{Ne}$ ($0.26\%$), and $^{22}_{10}\text{Ne}$ ($8.82\%$).

Mass-to-Charge Ratio ($m/z$)

The ratio of the mass of an ion to its electric charge, governing the extent of deflection in a mass spectrometer's magnetic field.

Molecular Mass

The sum of the atomic masses (in $\text{amu}$) of all atoms represented in a single molecule of a substance.

Percent Composition by Mass

The percentage of the total mass of a compound contributed by each individual constituent element.

Percent Composition Formula

$\text{percent composition of an element} = \frac{n \times \text{molar mass of element}}{\text{molar mass of compound}} \times 100\%$, where $n$ is the number of moles of the element in one mole of compound.

Percent Composition of Hydrogen Peroxide ($H_2O_2$)

$5.926\%\,\text{H}$ and $94.06\%\,\text{O}$, calculated from a molar mass of $34.02\,\text{g}$, yielding a total of $99.99\%$ due to rounding.

Empirical Formula

A chemical formula showing the elements present in the simplest whole-number ratio of their atoms.

Empirical Formula Mass (EFM)

The sum of the average atomic masses of all atoms present in the empirical formula unit of a compound.

Molecular Formula

A chemical formula showing the exact, actual number of atoms of each element present in one molecule of a compound.

Molecular Formula Multiplier ($n$)

The integer ratio calculated by dividing the experimental molar mass by the empirical formula mass: $n = \frac{\text{molar mass}}{\text{empirical formula mass}}$.

Chemical Reaction

A process in which one or more starting substances (reactants) break bonds and form new bonds to yield new substances (products) with distinct properties.

Reactants

The starting materials undergoing chemical change, written on the left side of a chemical equation.

Products

The newly generated chemical substances produced by a reaction, written on the right side of a chemical equation.

Chemical Equation

A symbolic representation of a chemical reaction utilizing chemical formulas and stoichiometric coefficients.

Plus Sign ($$+$) in Chemical Equations

A mathematical symbol meaning 'reacts with' when placed between reactants, or 'and' when placed between products.

Reaction Arrow ($\rightarrow$)

A directed symbol indicating 'yields' or 'forms' that points from reactants toward products.

Observable Characteristics of Chemical Reactions

Macroscopic indicators including color changes, gas evolution, precipitate formation, temperature changes (heat/light), or physical state changes.

Law of Conservation of Mass

The fundamental scientific law stating that matter is neither created nor destroyed during a chemical reaction.

Balanced Chemical Equation

An equation with stoichiometric coefficients adjusted such that the total number of atoms of each element is identical on both sides.

Skeletal Equation

An initial, unbalanced qualitative chemical equation displaying only formulas without stoichiometric balance.

Stoichiometry

The quantitative study and calculation of relative amounts of reactants and products involved in a chemical reaction.

Stoichiometry Road Map

A systemic diagram charting conversion pathways among masses, volumes at STP, particle counts, and moles via mole ratios.

Molar Volume at STP

The conversion factor of $22.4\,L\,mol^{-1}$ representing the volume occupied by one mole of any ideal gas at standard temperature and pressure.

Mole Ratio

A conversion factor derived directly from the stoichiometric coefficients of a balanced chemical equation.

Limiting Reagent

The reactant that is completely consumed first in a reaction, thereby limiting and deciding the maximum amount of product formed.

Excess Reagent

The reactant present in a quantity greater than required to react with the limiting reagent, leaving unreacted material behind.

Theoretical Yield

The maximum amount of product that would result if all of the limiting reagent were converted to product, as predicted by stoichiometry.

Actual Yield

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

Percent Yield

The ratio of actual yield to theoretical yield multiplied by 100, measuring reaction efficiency: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$.

Reversible Reaction (Yield Impact)

A reaction that proceeds in both directions simultaneously and does not go $100\%$ to completion, lowering the actual yield.

Solution

A homogeneous mixture composed of two or more substances uniformly dispersed at the molecular or ionic level.

Solute

The minor component in a solution, present in a smaller amount relative to the solvent.

Solvent

The substance present in the largest amount in a solution, which dissolves the solute.

Aqueous Solution

A homogeneous liquid mixture in which liquid water acts as the solvent.

Electrolyte

A substance that dissolves in water to yield mobile ions, resulting in a solution that conducts electricity.

Nonelectrolyte

A substance that dissolves in water as intact neutral molecules, resulting in a solution that does not conduct electricity.

Strong Electrolyte

A solute that dissociates or ionizes completely into free cations and anions when dissolved in aqueous solution.

Weak Electrolyte

A solute that only partially ionizes in aqueous solution, leaving the majority of dissolved solute as neutral molecules.

Examples of Strong Electrolytes

Inorganic strong acids and bases such as $HCl$, $HNO_3$, $HClO_4$, $H_2SO_4$, $NaOH$, $Ba(OH)_2$, and soluble ionic salts.

Examples of Weak Electrolytes

Weak molecular acids and bases including $CH_3COOH$, $HF$, $HNO_2$, $NH_3$, and pure $H_2O$.

Examples of Nonelectrolytes

Water-soluble molecular compounds including urea ($(NH_2)_2CO$), methanol ($CH_3OH$), ethanol ($C_2H_5OH$), glucose ($C_6H_{12}O_6$), and sucrose ($C_{12}H_{22}O_{11}$).

Solubility

The maximum amount of solute that can dissolve in a given volume of solvent at a specific temperature to produce a saturated solution.

Saturated Solution

A solution that contains the maximum concentration of dissolved solute possible in dynamic equilibrium with undissolved solute at a given temperature.

Solubility Rule: Alkali Metals and Ammonium

All ionic compounds containing alkali metal cations ($Li^+$, $Na^+$, $K^+$, $Rb^+$, $Cs^+$) or the ammonium cation ($NH_4^+$) are soluble in water with no exceptions.

Solubility Rule: Nitrates, Bicarbonates, and Chlorates

All ionic compounds containing nitrate ($NO_3^-$), bicarbonate ($HCO_3^-$), or chlorate ($ClO_3^-$) anions are completely soluble in water.

Solubility Rule: Halides

Ionic compounds containing chloride ($Cl^-$), bromide ($Br^-$), or iodide ($I^-$) are soluble, except when paired with $Ag^+$, $Hg_2^{2+}$, or $Pb^{2+}$.

Solubility Rule: Sulfates

Most sulfate compounds ($SO_4^{2-}$) are water-soluble, except those of silver ($Ag^+$), calcium ($Ca^{2+}$), strontium ($Sr^{2+}$), barium ($Ba^{2+}$), mercury(I) ($Hg_2^{2+}$), and lead ($Pb^{2+}$).

Solubility Rule: Insoluble Anions

Carbonates ($CO_3^{2-}$), phosphates ($PO_4^{3-}$), chromates ($CrO_4^{2-}$), and sulfides ($S^{2-}$) are insoluble, except compounds containing alkali metal ions or the ammonium ion.

Solubility Rule: Hydroxides

Hydroxides ($OH^-$) are generally insoluble in water, with the exceptions of compounds containing alkali metal ions or the barium ion ($Ba^{2+}$).

Precipitation Reaction

An aqueous chemical reaction characterized by the combining of dissolved ions to form an insoluble solid product called a precipitate.

Precipitate

An insoluble solid substance that separates from a solution during a chemical reaction.

Metathesis Reaction

A reaction (also termed double-displacement) involving an exchange of ionic partners (cations swap anions) between two reacting compounds.

Molecular Equation

A chemical equation in which the formulas of all reacting and produced compounds are written as whole, neutral formula units.

Ionic Equation

A chemical equation that explicitly shows all dissolved, completely dissociated electrolytes as separate free hydrated ions.

Spectator Ions

Ions in an aqueous reaction mixture that do not participate directly in chemical bonding changes and remain unchanged on both sides of the equation.

Net Ionic Equation

An equation derived by removing all spectator ions from an ionic equation, showing only the chemical species that directly participate in the reaction.

Neutralization Reaction

An acid-base reaction between an acid and a base that typically produces liquid water and a dissolved ionic salt.

Salt

An ionic compound formed in a neutralization reaction consisting of a cation other than $H^+$ and an anion other than $OH^-$ or $O^{2-}$.

Gas-Forming Acid-Base Reaction: Carbonates and Bicarbonates

The reaction of an acid with a carbonate ($CO_3^{2-}$) or bicarbonate ($HCO_3^-$) producing a salt, water, and gaseous carbon dioxide ($CO_2(g)$).

Carbonic Acid Decomposition

The breakdown of unstable intermediate carbonic acid ($H_2CO_3(aq)$) in aqueous solution into liquid water ($H_2O(l)$) and carbon dioxide gas ($CO_2(g)$).

Gas-Forming Reaction: Sulfites ($SO_3^{2-}$)

The reaction of sulfite salts with acids to yield a salt, water, and sulfur dioxide gas ($SO_2(g)$).

Gas-Forming Reaction: Sulfides ($S^{2-}$)

The reaction of metal sulfides with acids directly forming a salt and gaseous hydrogen sulfide ($H_2S(g)$).

Oxidation-Reduction (Redox) Reaction

A coupled chemical process defined by the direct or formal transfer of electrons from a reducing agent to an oxidizing agent.

Oxidation

The loss of one or more electrons by an atom, molecule, or ion, corresponding to an algebraic increase in its oxidation number.

Reduction

The gain of one or more electrons by an atom, molecule, or ion, corresponding to an algebraic decrease in its oxidation number.

Oxidizing Agent (Oxidant)

The electron acceptor in a redox reaction that causes another species to be oxidized while itself becoming reduced.

Reducing Agent (Reductant)

The electron donor in a redox reaction that causes another species to be reduced while itself becoming oxidized.

Half-Reaction

An explicit stoichiometric equation that isolates and displays either the oxidation process or the reduction process with its involved electrons.

Disproportionation Reaction

A specialized redox process in which an element in a single oxidation state simultaneously undergoes both oxidation and reduction.

Oxidation Number (Oxidation State)

The formal charge an atom would carry in a compound if all chemical bonds were completely ionic.

Oxidation Number Rule: Free Elements

Each atom in an uncombined elemental state (such as $H_2$, $Br_2$, $Na$, $Be$, $K$, $O_2$, or $P_4$) has an oxidation number of zero.

Oxidation Number Rule: Monatomic Ions

The oxidation number of any single-atom ion is exactly equal to the electrical charge on that ion.

Oxidation Number Rule: Oxygen

Oxygen is assigned $-2$ in most compounds, $-1$ in peroxides ($O_2^{2-}$), and $-\frac{1}{2}$ in superoxides ($O_2^-$).

Oxidation Number Rule: Hydrogen

Hydrogen has an oxidation state of $+1$ in nonmetal compounds, but $-1$ in binary metallic hydrides such as $LiH$, $NaH$, or $CaH_2$.

Oxidation Number Rule: Fluorine

Fluorine carries an oxidation state of $-1$ in all of its chemical compounds.

Oxidation Number Rule: Algebraic Sums

The sum of the oxidation numbers of all atoms in a neutral molecule must equal zero, while in a polyatomic ion it must equal the net ionic charge.

Concentration of a Solution

A quantitative measurement specifying the amount of dissolved solute present in a given amount of solvent or total solution.

Molarity ($M$)

The number of moles of dissolved solute per liter of solution: $M = \frac{n}{V}$.

Stock Solution

A concentrated solution stored in a laboratory stockroom from which less concentrated working solutions are prepared.

Dilution

The laboratory procedure of preparing a less concentrated solution by adding more solvent to a measured volume of stock solution.

Dilution Principle

The physical rule stating that the total moles of solute remain constant during dilution: $\text{moles before dilution} = \text{moles after dilution}$.