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isotopes are elements of the atomic number but have different mass numbers.


No. If they had different atomic numbers, they would not be the same atom and would be different elements. Different isotopes have different mass numbers.


Each isotope of an element has an atomic mass. The average of all the atomic masses of an element's isotopes gives an elements atomic weight. For this reason, atomic weights are decimal numbers.


The isotopes of the same element has the same atomic number. But the mass number (atomic mass) is different.


The atomic numbers of an element is the average mass of all that element's isotopes. Some elements have A LOT of isotopes. Sometimes, that average isn't always a whole number.


Atomic masses of elements are usually not whole numbers because of isotopes. Elements tend to exist as more than one isotope, so the atomic mass in the corner is a weighted average of all of the isotopes an element exists as.


The atomic number of the isotopes of an element is identical; the mass number is different.


Isotopes were then explained as elements with the same number of protons, --- In the 1950s, the development of improved particle accelerators and --- All known isotopes of elements with atomic numbers greater than 82 are radioactive.Evaa_I_WILL_H3LP_YOU


Isotopes of an element have nuclei with the same number of protons (the same atomic number) but different numbers of neutrons Neutrons (:


Isotopes have different numbers of electrons, but not different atomic numbers (numbers of protons) or they'd be different elements.



They have different numbers of neutrons, which changes the atomic mass and nuclear properties.


Isotopes are those element which have same atomic numbers but different atomic masses. Example- Hydrogen has three isotopes, namely, protium, deuterium and tritium Isobars are those elements which have same atomic mass but different atomic number. Example- Calcium and Argon


Yes, there are elements which have isotopes, i.e same atomic number & different atomic weight, so atoms f different elements can have same atomic weights.


That it is not hydrogen-1. All other elements and their isotopes have differing atomic masses and numbers.


Elements also possess isotopes. So their average atomic mass is rarely whole number.


Elements can exist in the form of different isotopes. Isotopes of the same element have the same number of protons in their nuclei but have different numbers of neutrons. The first gives them the same atomic number and chemical properties while the second gives them different atomic weights.


Depends on what version you are looking at. The integers going from 1 to around 100 are the atomic numbers of the elements and these are the numbers of protons (= numbers of electrons) in the atom. Numbers which run from 1 to around 250 and are integers are the atomic weights of isotopes, and numbers which are not integers are atomic weights of elements found in nature, which are mixtures of isiotopes.


No. Isotopes have the same atomic number, protons and electrons. They have different neutrons.


Generally isotopes of transfermium family elements (elements with atomic numbers between 101 and 118) are extremely unstable. But many isotopes have very short half lives.


No: All atoms of the same element have the same atomic number, but isotopes have different mass numbers.


Yes, it must be used as all elements have one or more isotopes and all elements have an atomic mass.


Oxygen. In fact, the atomic number of oxygen is more accurately represented by 8.0, since many elements exist as isotopes and the atomic number is actually an average value. The atomic numbers of all elements can be found in the Periodic Table of Elements.


The periodic table give the atomic numbers and the atomic weights of chemical elements.


Atomic mass numbers are not properties of elements overall, but only of particular isotopes of elements. The only stable element with an isotope with mass number 11 is boron. Beryllium and carbon also have isotopes with mass number 11, but these are radioactive.



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