Empirical and Molecular Formula Lecture Review

A comprehensive vocabulary flashcard set covering empirical and molecular formulas, step-by-step determination methods, chemical formulas, and calculated sample examples from the lecture.

Compound

A substance formed when two or more atoms of different elements are chemically combined.

Molecular Formula

The chemical formula for a compound that gives the total number of atoms of each element in a molecule.

Empirical Formula

A formula that gives the proportions of the elements present in a compound in the lowest whole-number ratio, but not the actual numbers or arrangement of atoms.

Chemical Formula

A written representation of the chemical proportions of atoms that comprise a particular compound.

Structural Formula

A chemical formula that shows the arrangement of atoms within a molecule or compound.

Percent Composition Formula

The formula used to calculate an element's percentage mass in a compound: $\% \text{ element} = \frac{\text{molar mass of element}}{\text{molar mass of compound}} \times 100\%$.

Element Mass Formula from Percent Composition

The formula used to calculate the mass of an element when total mass and percent composition are known: $\text{mass of element} = \frac{\text{total mass of compound} \times \% \text{ element}}{100}$.

100-g Total Mass Assumption

The procedural rule applied when total mass is not provided in an empirical formula problem, where a total compound mass of $100\,g$ is assumed so percent values directly equal mass in grams.

Mass-to-Mole Conversion Step

The second step in determining an empirical formula where each element's mass is divided by its molar mass or relative atomic mass to find its amount in moles.

Preliminary Subscripts

The calculated mole amounts of each constituent element placed initially as temporary subscripts in a chemical formula.

Whole-Number Ratio Conversion Step

The step where all calculated mole amounts of constituent elements are divided by the smallest number of moles obtained among them.

Empirical Formula of Hydrogen Peroxide ($H_2O_2$)

$HO$, obtained by reducing the molecular subscript ratio from 2:2 to the simplest ratio of 1:1.

Empirical Formula of Benzene ($C_6H_6$)

$CH$, obtained by dividing both subscripts in the molecular formula by 6.

Empirical Formula of Tetraphosphorus Decoxide ($P_4O_{10}$)

$P_2O_5$, obtained by dividing the molecular subscripts 4 and 10 by their greatest common divisor, 2.

Empirical Formula of Glucose ($C_6H_{12}O_6$)

$CH_2O$, obtained by dividing all subscripts in the molecular formula by 6.

Empirical Formula of Zinc-Phosphorus-Oxygen Compound Sample

$Zn_7P_4O_{18}$, determined from a sample containing $0.28\,mol$ Zn, $0.16\,mol$ P, and $0.72\,mol$ O.

Zinc Ratio in $Zn_7P_4O_{18}$ Step 1

$1.75$, calculated by dividing $0.28\,mol$ of zinc by the smallest mole value of $0.16\,mol$.

Oxygen Ratio in $Zn_7P_4O_{18}$ Step 1

$4.5$, calculated by dividing $0.72\,mol$ of oxygen by the smallest mole value of $0.16\,mol$.

Ratio Multiplier for Zinc-Phosphorus-Oxygen Sample

$4$, used to convert the decimal mole ratios of 1.75, 1, and 4.5 into whole numbers (7, 4, and 18).

Empirical Formula of Carbon-Hydrogen-Oxygen Sample

$C_5H_2O$, derived from a sample containing $1.25\,mol$ C, $0.50\,mol$ H, and $0.25\,mol$ O.

Carbon Mole Ratio in C-H-O Sample

$5$, obtained by dividing $1.25\,mol$ of carbon by the smallest mole amount of $0.25\,mol$.

Hydrogen Mole Ratio in C-H-O Sample

$2$, obtained by dividing $0.50\,mol$ of hydrogen by the smallest mole amount of $0.25\,mol$.

Smallest Mole Value in C-H-O Sample

$0.25\,mol$, corresponding to oxygen, used to simplify all element mole counts to small whole numbers.

Moles of Sodium in $2.82\,g$ Na

$0.1227\,mol$, calculated by dividing $2.82\,g$ by sodium's molar mass of $22.99\,g\,mol^{-1}$.

Moles of Chlorine in $4.35\,g$ Cl

$0.1227\,mol$, calculated by dividing $4.35\,g$ by chlorine's molar mass of $35.45\,g\,mol^{-1}$.

Moles of Oxygen in $7.83\,g$ O

$0.4894\,mol$, calculated by dividing $7.83\,g$ by oxygen's molar mass of $16.00\,g\,mol^{-1}$.

Mole Ratio of Na:Cl:O in Ionic Compound Sample

$1:1:4$, obtained by dividing $0.1227\,mol$ Na, $0.1227\,mol$ Cl, and $0.4894\,mol$ O each by $0.1227\,mol$.

Empirical Formula of Na-Cl-O Ionic Compound

$NaClO_4$, derived from the simplified whole-number ratio of 1 Na : 1 Cl : 4 O.

Relative Atomic Mass of Aluminium ($Al$)

$27$, as used in calculating empirical formulas for aluminium compounds.

Relative Atomic Mass of Chlorine ($Cl$)

$35.5$, as used in empirical formula calculations involving chlorine.

Moles of Aluminium in $0.135\,g$ Sample

$0.005\,mol$, calculated as $\frac{0.135\,g}{27\,g\,mol^{-1}}$.

Moles of Chlorine in $0.533\,g$ Sample

$0.015\,mol$, calculated as $\frac{0.533\,g}{35.5\,g\,mol^{-1}}$.

Empirical Formula of Aluminium Chloride Sample

$AlCl_3$, calculated from $0.135\,g$ Al and $0.533\,g$ Cl.

Smallest Mole Value in Aluminium Chloride Calculation

$0.005\,mol$, which is the mole value for aluminium used to divide both elemental mole amounts.

Relative Atomic Mass of Carbon ($C$)

$12$, as given in the carbon chloride calculation example.

Moles of Carbon in $0.36\,g$ Sample

$0.03\,mol$, calculated as $\frac{0.36\,g}{12\,g\,mol^{-1}}$.

Moles of Chlorine in $4.26\,g$ Sample

$0.12\,mol$, calculated as $\frac{4.26\,g}{35.5\,g\,mol^{-1}}$.

Smallest Mole Value in Carbon Chloride Calculation

$0.03\,mol$, which is the calculated molar amount of carbon used as the divisor for ratio calculation.

Empirical Formula of Carbon Chloride Sample

$CCl_4$, determined from $0.36\,g$ C and $4.26\,g$ Cl.

Mole Ratio of Carbon to Chlorine in Carbon Chloride Sample

$1:4$, derived by dividing $0.03\,mol$ C and $0.12\,mol$ Cl by $0.03\,mol$.

Simplest Whole-Number Ratio

The defining characteristic of an empirical formula's subscript representation.

Whole-Number Multiple

The relationship between a molecular formula's subscripts and its corresponding empirical formula's subscripts.

Formula Cards

A manipulative learning tool used by students to practice matching molecular formulas with their corresponding empirical formulas.

Empirical Formula Reducibility

An empirical formula cannot be reduced further because its subscripts are already in their lowest whole-number terms.

Molecular Formula Reducibility

A molecular formula can sometimes be reduced to a simpler whole-number ratio unless its subscripts already have a greatest common divisor of 1.

Empirical Formula as a Starting Point

The empirical formula serves as the foundational starting point for determining a compound's actual molecular formula.

Assumed Element Mass for $65\%$ Element A

$65\,g$, based on assuming a standard total compound mass of $100\,g$.

Assumed Element Mass for $25\%$ Element B

$25\,g$, based on assuming a standard total compound mass of $100\,g$.

Assumed Element Mass for $10\%$ Element C

$10\,g$, based on assuming a standard total compound mass of $100\,g$.

Values-Focused Reflection Qualities

Being careful, disciplined, and cooperative to achieve accurate determinations of empirical and molecular formulas.