3.6
Masse atomique
Les protons et les neutrons ont approximativement la même masse, soit environ 1,67 × 10-24 grammes. Les scientifiques définissent arbitrairement cette…
La masse atomique d'un élément, parfois appelée masse atomique, est approximativement égale à la somme du nombre de protons et de neutrons. Celle-ci est exprimée en unités de masse atomique, abrégée en amu, ou Daltons, et est représentée près de son symbole chimique sur le tableau périodique.
Mais, la plupart des éléments ont plusieurs isotopes qui diffèrent par le nombre de neutrons et donc par leurs masses.
Pour de tels éléments, le poids atomique est la masse moyenne de tous les isotopes d’un élément, pondérée en fonction de leur abondance relative ou du pourcentage de chaque isotope présent dans la nature.
Par exemple, le chlore a deux isotopes naturels. La masse de chlore-35 est de 34,97 amu et celle de chlore-37 est de 36,97 amu.
Lorsqu’on multiplie par leurs abondances relatives, 75,77 % pour le chlore 35 et 24,23 % pour le chlore 37, et qu’on additionne, on obtient la masse atomique finale de 35,45 amu.
Étant donné que le chlore 35 est plus abondant dans la nature que le chlore 37, le poids atomique du chlore est proche de 35 amu.
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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.