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1.3Atomic Structure and Properties

Elements & Mixtures — AP CHEMISTRY Revision Notes & Topical Questions

5 min read 9 parts Lesson 10 of 16 1 quiz

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Lesson summary

1.3 Elements & Mixtures — AP Notes — AP CHEMISTRY notes: Explain how elemental composition by mass relates to the empirical formula of a pure substance and to the composition of a mixture.Element = a pure substance made of one type of atom, which cannot be broken down chemically.…

  • Mass percent of an element
  • Molecular compound
  • Ionic compound
  • Molecular formula
  • Percent purity
  • Combustion analysis

Part 1 of 7Pure substances and mixtures

Element
Element a pure substance made of one type of atom, which cannot be broken down chemically.
Compound
Compound a pure substance of two or more elements chemically bonded in a fixed mass ratio.
Pure substance
Pure substance fixed composition throughout, with sharp melting and boiling points.
Mixture
Mixture substances physically combined in variable proportions, with components keeping their properties.
Homogeneous mixtures (solutions, alloys, air) are uniform; heterogeneous mixtures are not.
PropertyElementCompoundMixture
CompositionOne type of atomFixed ratio of elementsVariable
SeparationCannot be broken down chemicallyChemical reaction onlyPhysical methods
Melting and boiling pointSharpSharpOver a range
ExampleO2, FeH2O, NaClAir, brass, sea water

Part 2 of 7Law of definite proportions

Law of definite proportions
A given compound always has the same elements in the same mass proportions, wherever it comes from.
Pure water is always 11.19% hydrogen and 88.81% oxygen by mass.
Mass percent data can therefore identify a pure compound.
Mass percent of an element
(mass of the element in one mole of compound ÷ molar mass) × 100.
The mass percents of all elements in a pure compound sum to 100%.

Part 3 of 7Molecular and ionic formulae

Molecular compound
  • Molecular compound discrete covalently bonded molecules, such as H2O.
Ionic compound
  • Ionic compound no molecules;
    • the formula unit gives the fixed ion ratio in the lattice, so ionic formulae are empirical.
Empirical formula
Empirical formula the simplest whole-number mole ratio of the elements.
Molecular formula
Molecular formula the actual numbers of atoms, a whole-number multiple of the empirical formula.
Several compounds can share an empirical formula, such as C2H2 and C6H6 (both CH).
CompoundMolecular formulaEmpirical formula
GlucoseC6H12O6CH2O
Hydrogen peroxideH2O2HO
BenzeneC6H6CH

Part 4 of 7From mass percent to molecular formula

Assume a 100 g sample so each percentage becomes grams.
Convert each mass to moles by dividing by the molar mass.
Divide every mole value by the smallest.
Multiply all by 2, 3, 4 or 5 if ratios are decimals such as 0.5, 0.33, 0.25 or 0.2.
Write the empirical formula from the whole-number ratio.
Find n
Find n molar mass ÷ empirical formula mass and multiply every subscript by n.

Part 5 of 7Elemental analysis, purity and combustion analysis

Impurities change the measured mass percent of an element.
  • A sample of NaCl with a higher Cl percent than 60.66% must contain an impurity richer in Cl, such as MgCl2 (74.5% Cl).
  • An impurity with a lower Cl percent, such as KCl (47.6%), lowers it.
Percent purity
(mass of pure substance ÷ mass of sample) × 100.
Combustion analysis
  • Combustion analysis all C becomes CO2 and all H becomes H2O;
    • the masses give moles of C and H, and oxygen is found by difference.

Part 6 of 7Composition and separation of mixtures

Composition and separation of mixtures
Use an element found in only one component to find that component's mass.
Moles of the element equal moles of the component when the formula has one atom of the element.
Mixtures are separated by physical differences.
  • Filtration separates by particle size or solubility.
  • Distillation separates by boiling point.
  • Chromatography separates by attraction to the stationary versus mobile phase.

Part 7 of 7More from the specification

More from the specification
Chemical formula units versus localized covalent bonds
  • Chemical formula units versus localized covalent bonds network compounds contain repeating stoichiometric formula units,
    • while discrete molecules feature directional covalent bonds.
Chloride
  • Salt of hydrochloric acid.
Matter
  • Everything you can see and feel; it exists as solid, liquid or gas.
Sodium chloride
  • Compound of sodium and chlorine; common salt.
Covalent bond
  • A shared pair of electrons between two atoms, giving noble gas configurations.

Key terms

TermMeaning
ElementPure substance made of one type of atom.
CompoundPure substance of elements chemically bonded in a fixed ratio.
Law of definite proportionsA compound always has the same mass ratio of its elements.
Empirical formulaSimplest whole-number ratio of atoms.
Molecular formulaActual number of atoms of each element in a molecule.
Percent purityMass of pure substance as a percentage of the sample mass.

Avoid these mistakes

1
Treating the empirical formula as the actual molecule — the molecular formula is a whole-number multiple
2
Rounding a mole ratio such as 1.5 to 2 — multiply all ratios to reach whole numbers
3
Calling a mixture a compound — mixture components are not chemically bonded
4
Using mass ratios directly in formulae — convert to moles first

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