3.6
Peso Atómico
Los protones y neutrones tienen aproximadamente la misma masa, alrededor de 1.67 × 10^-24 gramos. Los científicos definen arbitrariamente esta cantida…
El peso atómico de un elemento, a veces llamado masa atómica, es aproximadamente igual a la suma del número de protones y neutrones. Esto se expresa en unidades de masa atómica, abreviado como amu, o Daltons, y se representa cerca de su símbolo químico en la tabla periódica.
Pero la mayoría de los elementos tienen varios isótopos que difieren en el número de neutrones y, por lo tanto, en sus masas.
Para tales elementos, el peso atómico es la masa promedio de todos los isótopos de un elemento, ponderada de acuerdo con su abundancia relativa o el porcentaje de cada isótopo que se encuentra en la naturaleza.
Por ejemplo, el cloro tiene dos isótopos naturales. La masa de cloro-35 es de 34,97 uma y la de cloro-37 es de 36,97 uma.
Cuando se multiplica por sus abundancias relativas, 75,77% para el cloro-35 y 24,23% para el cloro-37, y se suman, se obtiene el peso atómico final de 35,45 uma.
Dado que el cloro-35 es más abundante en la naturaleza que el cloro-37, el peso atómico del cloro es cercano a 35 uma.
View the full transcript and gain access to JoVE Core videos
Q1: What is the difference between atomic weight and atomic mass?
Atomic weight and atomic mass are closely related terms. Atomic mass refers to the sum of protons and neutrons in a single atom, expressed in atomic mass units (amu) or Daltons. Atomic weight is the weighted average mass of all naturally occurring isotopes of an element, accounting for their relative abundance in nature. For example, chlorine's atomic weight of 35.45 amu reflects the average of its two isotopes.
Q2: Why don't electrons contribute significantly to an element's atomic weight?
Electrons are much smaller than protons and neutrons, weighing only about 1/1800 of an atomic mass unit. Since protons and neutrons each weigh approximately 1.67 × 10⁻²⁴ grams, their combined mass dominates the atom's total mass. Therefore, scientists ignore electron mass when calculating atomic weight and base it solely on the number of protons and neutrons.
Q3: How is atomic weight calculated for elements with multiple isotopes?
Atomic weight is calculated by summing the mass of each isotope multiplied by its natural fractional abundance. For chlorine, chlorine-35 (34.97 amu at 75.77% abundance) and chlorine-37 (36.97 amu at 24.23% abundance) combine to yield an atomic weight of 35.45 amu. This weighted average reflects the isotope composition found in nature.
Q4: What are atomic mass units and why are they used?
Atomic mass units (amu), also called Daltons, are standardized units for measuring atomic and subatomic particle masses. Scientists arbitrarily defined one amu as the mass of a proton or neutron, approximately 1.67 × 10⁻²⁴ grams. This convenient scale allows chemists to express atomic weights as simple whole numbers or decimals rather than using unwieldy scientific notation.
Q5: How does isotope abundance affect an element's atomic weight on the periodic table?
The periodic table displays atomic weight as a weighted average based on isotope abundance. Since chlorine-35 is more abundant than chlorine-37 in nature, chlorine's listed atomic weight of 35.45 amu is closer to 35 than to 37. Elements with more evenly distributed isotopes show atomic weights further from whole numbers, reflecting their natural isotopic composition.
Q6: What is the relationship between atomic weight and molar mass?
A mole of an element has a mass in grams equal to its atomic weight. For compounds, molecular weight is the sum of atomic weights of all component atoms. Glucose (C₆H₁₂O₆) has a molecular weight of 180.156 grams per mole, calculated by summing six carbons (72.066 g), twelve hydrogens (12.108 g), and six oxygens (95.982 g).
Q7: Why is understanding isotope composition important for determining atomic weight?
Most naturally occurring elements consist of multiple isotopes with different masses and abundances. Ignoring isotope composition would yield inaccurate atomic weights. By considering both the mass and natural fractional abundance of each isotope, chemists obtain precise atomic weights that reflect real elemental composition, essential for accurate stoichiometric calculations in chemical reactions.