Isotopes

Isotopes are different forms of the same atom with different numbers of neutrons. For example, normal carbon is carbon-12, which has 6 protons, 6 neutrons. Carbon-14 has 2 extra neutrons making it heavier. But they have the same numbers of electrons and so react and form compounds in the same way - so the formula of methane CH4 or carbon dioxide CO2 is the same. If you want to be very specific you can indicate the labeled position with a superscript, thus: 13CH4 or CH313CH2OH indicate methane with the carbon-13 isotope or ethane labeled with carbon-13 at the alpha position, respectively.

Isotopes are variations of an element. Isotopes of one element always have the same number of protons because it's the number of protons that define the element. Isotopes of an element have the same number of electrons which will equal the number of protons. (That ignores atoms that are in the form of ions.)

Different isotopes of the same element have different numbers of neutrons. The variations in the number of neutrons make the mass of atoms different but they do not change the element nor do they change the chemical properties of the element.

Some isotopes of an element can be radio-active, that is, they are unstable and can emit an alpha or beta particle or a gamma ray. As an example, carbon has 6 protons and most atoms have 6 neutrons. A few carbon atoms have 8 neutrons with a mass of 14. It is known as "Carbon 14" and is radio active. That is because the nucleus can emit a beta particle and in doing so, the atom actually changes to a nitrogen atom.

The stability of an isotope nucleus depends on the balance between the strong nuclear force, which holds protons and neutrons together, and the electromagnetic force, which repels protons due to their positive charge. Isotopes with too many or too few neutrons compared to protons may be unstable and undergo radioactive decay to achieve a more stable configuration.

The term isotope is used to indicate the different varieties of a single element, based upon variations in the number of neutrons in the nucleus. Every atom can be described as an isotope if we specify the number of neutrons. And every atom can lose electrons (a process officially called ionization). There is no relation between the number of neutrons and the loss of electrons.

The isotopes of an atom are defined by the number of neutrons their nuclei have for their fixed number of protons. However, I wouldn't say that neutrons are solely responsible for the presence of isotopes because isotopes also depend on the existence of many other particles such as protons, quarks and gluons. Neutrons certainly are responsible for the way we label isotopes, though.

Well a chemistry class for a freshman year would basically be chemical problems and lots of algebra work (since most 9th graders haven't taken geometry yet) If you can avoid taking chemistry freshman year, I would recommend it. But

depositing a metal to another

The isotopes of nickel share the same number of protons (28) but differ in the number of neutrons in their nuclei, resulting in different mass numbers. Each isotope of nickel has a different abundance in nature, with nickel-58 being the most common.

Making any change in the half-life of an isotope of any element is generally something that lies outside our abilities. A very few radioactive materials have demonstrated a change in their half-lives when bathed in intense magnetic fields. Generally, however, the half-life on a given radionuclide is not something that can be changed. A number of experiments have been conducted wherein investigators have deliberately sought to influence radioactive half-life, but in all but the rarest cases, radionuclides are sublimely resistant to having their half-lives changed.

The age of rocks and minerals through radiometric dating, the presence of tumors in the body through imaging techniques like PET scans, and the effectiveness of treatments for cancer through radiation therapy.

Strontium-88

It is the closest to the Strontium atomic mass.

Nuclear Isotopes are different types of atoms of the same chemical element having a different number of neutrons, which that chemical element is involved in the nuclear industry (e.g.: Uranium-235, Astatine-211, Americium-241). Most are very unstable, but a handful of them naturally occur on Earth (e.g.: Iodine-131, Carbon-14, Caesium-137)

Nuclear Isotopes are radioactive and should be treated with extreme care! No joke here when handling even the safest of nuclear isotopes (just in case you were wondering, Uranium-235 or Depleted Uranium, is the safest nuclear isotope).

Yes because they look at the ammount of radioactive decay and they can determine the age of the granite. As the radioactive isotope decreases the non radioactive element increases. The less radioactive decay ammount there is the older it is.

The most abundant isotope of xenon is xenon-132. It makes up about 26.9% of naturally occurring xenon.

No, 1-hexyne is not an isotope.

Isotopes are same elements that have the same number of protons (and therefore the same chemical properties) but different numbers of neutrons. They have slightly different atomic masses due to the varying number of neutrons in their nuclei.

1-hexyne, on the other hand, is a specific chemical compound. It is an alkyne with the molecular formula C6H10 and a carbon-carbon triple bond at the first position in a hexane chain. Isotopes are not specific chemical compounds, but they are variations of elements.

Henri Bacquerel first experimented with phosphorescent materials and discovered radioactivity. The Curies and Ernest Rutherford later experimented with it.

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The first hypothesis on isotopes is from Frederick Soddy (1912); the practical confirmation is attributed to J. J. Thomson (1913).

Henri Becquerel discovered the phenomenon of radioactivity studying uranium salts.

Since the radioactive decay of 14C is first order, one can use the equation of
fraction remaining = 0.5^n where n is the number of half lives. 28650yr/5730 yr = 5 half lives. Thus..
fraction remaining = 0.5^5 = 0.03125 or 3.125% would be remaining.

Isotopes of the same element have the same number of protons and electrons, but varied numbers of neutrons.

Isotopes are atoms of the same element with the same number of protons but different number of neutrons. Some examples of atoms with isotopes are hydrogen (protium, deuterium, tritium), carbon (carbon-12, carbon-13, carbon-14), and uranium (uranium-235, uranium-238).

Neon 16 through to Neon 34 have been synthesised/discovered.

Neon 9 through to Neon 11 are the naturally occurring isotopes,

all of which are stable. Neon 9 comprises over 90% of natural Neon in the air.

Neon 15 is the most recently reported and has not been verified (as of may 2014). Other isotopes, both heavier and lighter, could be synthesised in the future, increasing the isotope count.

Xenon is an element on the periodic table with the symbol Xe and atomic number 54. Xenon has several isotopes, including xenon-129 and xenon-131, which have different numbers of neutrons in their nuclei. Xenon is a colorless, odorless noble gas that is commonly used in various applications such as lighting and medical imaging.

The atomic mass minus the atomic number equals the number of neutrons. Thus in the case above the number of neutrons would calculate out to be 11. HOWEVER PLEASE NOTE Rhenium (Re) does not have an atomic mass of 86, it has two isotopes one of atomic mass 185 and another of atomic mass 187. Thus the real number of neutrons is 110 or 112.

You can use a mass spectrometer, that is it detects the mass of the particle based on inertia. As isotopes differ from each other in the number of neutrons they have, a difference in charge cannot be detected.

A stable isotope does not decay and therefore, maintains a constant concentration on Earth. An unstable isotope, also known as a radioactive isotope, decays at a predictable

and measurable rate on Earth. An unstable isotope may decay by the ejection of an electron or positron, known as beta decay, or by the ejection of two protons and two neutrons, known as alpha decay.