Isotopes and Average Atomic Mass

Flashcards covering the definitions of isotopes, atomic notation, subatomic particles, relative abundance, and weighted average mass calculations.

Isotopes

Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different atomic masses.

$Z$ (Atomic Number)

The number of protons in an atom, which determines the element's identity.

$A$ (Mass Number)

The sum of the number of protons and the number of neutrons in an atom ($A = p + n$).

$$X$$

The element symbol or abbreviation used in atomic notation.

Relative Abundance

The percentage of a specific isotope of an element found in nature.

Fractional Abundance

The decimal representation of the relative abundance percentage (e.g., $0.760$ for $76.0\%$) used in weighted average calculations.

Average Atomic Mass

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

Protium

An isotope of hydrogen ($^1H$) consisting of $1$ proton and $0$ neutrons.

Deuterium

An isotope of hydrogen ($^2H$) consisting of $1$ proton and $1$ neutron.

Tritium

An isotope of hydrogen ($^3H$) consisting of $1$ proton and $2$ neutrons.

Ion

The result of changing the number of electrons in an atom, which changes the charge but not the element identity or isotope status.

Isotope Hydrology

A field of study that uses stable isotope ratios (such as $^2H$ and $^{18}O$) to investigate water sources, movement, and evaporation patterns.

$u$ ($amu$)

The atomic mass unit used to express the mass of isotopes and average atomic mass values.

Neutron Calculation

Formula used to determine the count of neutrons in a specific isotope: $\text{Neutrons} = A - Z$.