In Mendeleev's periodic table of elements, the relative atomic masses (now referred to as atomic weights) were sometimes found to be greater than their actual values due to the limited knowledge and understanding of atomic structure and isotopes at the time.
Mendeleev developed his periodic table based on the observed chemical and physical properties of elements. He arranged elements in order of increasing atomic mass, placing elements with similar properties in the same vertical columns, known as groups or families. However, the atomic masses available during Mendeleev's time were not always accurate due to various factors, such as impurities in the samples and the existence of isotopes.
Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei. They have the same atomic number (representing the number of protons) but different atomic masses. At the time of Mendeleev's work, the concept of isotopes was not well understood. Therefore, the atomic masses listed in his table often represented the average masses of the naturally occurring isotopes rather than the specific masses of a single isotope.
This led to some discrepancies between the observed atomic masses and the actual masses of individual isotopes. As scientific knowledge advanced and isotopes were discovered and characterized, the atomic masses in the periodic table were refined to reflect the accurate values.
In modern periodic tables, the atomic weights are expressed as weighted averages of the isotopic masses, taking into account the abundance of each isotope. This provides a more accurate representation of the atomic mass of an element, considering the variations caused by different isotopes.