AP Chemistry Unit 1: Quantitative Atomic Foundations

mole

The amount of substance containing the same number of chemical units as there are atoms in exactly 12 grams of pure carbon-12.

Avogadro's Number (N_A)

6.022 x 10^23 particles/mol, the number of particles in one mole of a substance.

molar mass

The mass of one mole of a substance, expressed in grams per mole (g/mol).

atomic mass unit (amu)

A unit of mass used to express atomic and molecular weights; 1 amu is defined as one twelfth of the mass of an unbound neutral atom of carbon-12.

conversion between mass, moles, and particles

Using the formula n = m/M to convert between the amount of substance in moles (n), mass (m), and molar mass (M).

mass spectrometry

An analytical technique used to determine the relative abundance of isotopes in a sample by separating ions based on their mass-to-charge ratio.

mass spectrum

A graph that shows the isotopes of an element, with the x-axis representing the mass-to-charge ratio (m/z) and the y-axis representing relative intensity.

average atomic mass

A weighted average of the masses of an element's isotopes, calculated using the formula Avg Atomic Mass = ∑(Isotopic Massi × Fractional Abundancei).

Law of Definite Proportions

A chemical law stating that a given compound always contains its components in fixed ratio by mass.

percent composition by mass

The percentage of a specific element in a compound, calculated with % Element = (Mass of Element in Formula / Molar Mass of Compound) × 100.

empirical formula

The simplest whole-number ratio of atoms in a compound.

molecular formula

The actual number of atoms of each element in a molecule, which is often a whole-number multiple of the empirical formula.

determining empirical formula mnemonic

Percent to mass, mass to mole, divide by small, multiply 'til whole.

mixture

A physical combination of two or more substances where each retains its own properties.

purity analysis

The process of analyzing a mixture to determine the purity of a specific substance.

stoichiometry

The part of chemistry involving the calculation of reactants and products in chemical reactions.

isotopic mass

The mass of a specific isotope of an element, usually expressed in atomic mass units (amu).

relative intensity

The measure of how common a particular isotope is, often represented as a percentage in a mass spectrum.

percent to mass conversion

Assuming a 100 g sample allows you to convert percentages directly to grams for calculations in empirical formula determination.

diatomic elements

Elements that naturally exist as molecules composed of two atoms, such as O2 (oxygen gas).

fractional abundance

The proportion of a particular isotope of an element relative to the total amount of that element in a sample.

molar mass of Cu

63.55 g/mol, the molar mass of copper used in calculations involving copper.

mass-to-charge ratio (m/z)

A value used in mass spectrometry that refers to the mass of an ion divided by its charge.

common mistakes in mass spectrometry

Including not summing percentages correctly and confusing peak heights for mass.

calculating average atomic mass example

For isotopes X-10 and X-11, Avg Mass = (10.0 x 0.20) + (11.0 x 0.80) = 10.8 amu.

conversion between grams and moles formula

n = m/M, used for calculating the number of moles from mass and molar mass.

compressive stoichiometry

Using stoichiometric coefficients to relate the amounts of reactants and products in a chemical reaction.