Standard measurements in science are very important so that we can compare experimental data from one lab to another and make sure we all are talking about the same thing. Masses of individual atoms are very, very small. Using a modern device called a mass spectrometer, it is possible to measure such minuscule masses. An atom of oxygen, for example, has a mass of 2.
While comparisons of masses measured in grams would have some usefulness, it is far more practical to have a system that will allow us to more easily compare relative atomic masses. Scientists decided on using the carbon nuclide as the reference standard by which all other masses would be compared. By definition, one atom of carbon is assigned a mass of 12 atomic mass units amu.
An atomic mass unit is defined as a mass equal to one twelfth the mass of an atom of carbon The mass of any isotope of any element is expressed in relation to the carbon standard. For example, one atom of helium-4 has a mass of 4. An atom of sulfur has a mass of The carbon atom has six protons and six neutrons in its nucleus for a mass number of This technique separates the different isotopes of atoms to allow determination of the percent abundance or isotopic composition of the element in the given sample.
Each isotope appears as a peak in the mass spectrum. Mass spectrometry is used in a diverse range of applications, such as accurate determinations of molecular masses, drug testing, determining the age of archaelogical artifacts 14 C dating and for studying the chemistry of DNA see the Advanced Application below. Consider the mass spectrum of silicon, shown below.
The abundances are the same as those in Example 2. As you can see, there are three isotopes. Each peak represents one of the isotopes. The most abundant isotope has the highest peak intensity and the least abundant isotope has the smallest intensity. Since the peak intensities heights are proportional to the isotopic abundances, analysis of the data allows for the determination of the relative abundances of each isotope in the sample.
There are only two naturally occurring isotopes of Boron, 10 B and 11 B. If 10 B has a mass of Now we have two unknowns and only one equation.
We need another equation related to the fractional abundances. We know that the fractional abundances must add up to 1. Notice that there are no units associated with the fractional abundance. So naturally occuring Boron is In , the Italian scientist Amedeo Avogadro realized that the volume of a gas at a given pressure and temperature is proportional to the number of atoms or molecules composing it, regardless of which gas it was. This useful fact allowed chemists to compare the relative weights of equal volumes of different gases to determine the relative masses of the atoms composing them.
They measured atomic weights in terms of atomic mass units amu , where 1 amu was equal to one-twelfth of the mass of a carbon atom. When in the second half of the 19th century, chemists used other means to approximate the number of atoms in a given volume of gas — that famous constant known as Avogadro's number — they began producing rough estimates of the mass of a single atom by weighing the volume of the whole gas, and dividing by the number.
Many people use the terms weight and mass interchangeably, and even most scales offer options in units such as pounds and kilograms. And while mass and weight are related, they are not the same thing. When discussing atoms, many people use atomic weight and atomic mass interchangeably, even though they aren't quite the same thing either.
Atomic mass is defined as the number of protons and neutrons in an atom, where each proton and neutron has a mass of approximately 1 amu 1. The electrons within an atom are so miniscule compared to protons and neutrons that their mass is negligible.
The carbon atom, which is still used as the standard today, contains six protons and six neutrons for an atomic mass of twelve amu. Different isotopes of the same element same element with different amounts of neutrons do not have the same atomic mass. Carbon has an atomic mass of 13 amu. Atomic weight, unlike the weight of an object, has nothing to do with the pull of gravity. It is a unitless value that is a ratio of the atomic masses of naturally occurring isotopes of an element compared with that of one-twelfth the mass of carbon For elements such as beryllium or fluorine that only have one naturally occurring isotope, the atomic mass is equal to the atomic weight.
Carbon has two naturally occurring isotopes — carbon and carbon The atomic masses of each are The atomic weight is very similar to the mass of carbon due to the majority of carbon in nature being made of the carbon isotope.
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