There atomic mass is the sum of the masses of all protons, neutrons and electrons present in a single atom or molecule. The mass of an electron is so small that it is considered negligible and therefore is not included in the calculation. The term is also often used to refer to the average atomic mass of all isotopes of an element, although this use is technically incorrect. This second definition actually refers to the relative atomic mass, also called atomic weight of an element. The atomic weight takes into account the average of the masses of the natural isotopes of an element. Chemists must distinguish these two concepts during their activity because, for example, an incorrect value of the atomic mass can lead to errors in calculating the yield of an experiment.
Steps
Method 1 of 3: Finding the Atomic Mass on the Periodic Table
Step 1. Learn how atomic mass is represented
This can be expressed in the standard units of the International System (grams, kilograms, and so on), regardless of whether it refers to a single atom or a molecule. However, when denoted with these units, atomic mass values are extremely small and therefore atomic mass units (generally abbreviated to "uma") are preferred. An atomic mass unit corresponds to 1/12 of the standard atomic mass of the isotope 12 of carbon.
Atomic mass units indicate the mass expressed in grams of a mole of a given element or molecule. This is a very useful property when calculations are to be made, since it allows a simple conversion between mass and moles of a given quantity of atoms or molecules of the same type
Step 2. Find the atomic mass on the periodic table
Most periodic tables list the relative atomic masses (atomic weights) of all elements. The value is written at the bottom of the box that encloses the chemical symbol consisting of one or two letters. Generally it is a decimal number, more rarely an integer.
- Remember that the relative atomic masses you find in the periodic table are "average" values for each element. Elements have different "isotopes" - atoms with different masses because they have more or less neutrons in their nuclei. Therefore the relative atomic mass reported in the periodic table is an acceptable average value of the atoms of a given element, but Not is the mass of a single atom of the element itself.
- The relative atomic masses indicated on the periodic table are used for the calculation of the molar masses of atoms and molecules. The atomic masses, when they are expressed in uma as it happens on the periodic table, are technically numbers without units of measure. However, it is sufficient to multiply them by 1 g / mol to obtain a usable value of the molar mass, that is, the mass expressed in grams of a mole of atoms of the given element.
Step 3. Remember that the values shown on the periodic table are the average of the atomic mass for the particular element
As previously stated, the relative atomic masses that are placed in the box of each element of the periodic table represent the average value of all the atomic masses of the isotopes of that element. The average value is useful for many practical calculations, for example to find the molar mass of a molecule made up of several atoms. However, when you have to consider single atoms, this number is often not enough.
- Since it is the average of different types of isotopes, the figure expressed on the periodic table is not exactly the atomic mass of a single atom.
- The atomic mass of each atom must be calculated taking into account the precise number of protons and neutrons that make up its nucleus.
Method 2 of 3: Calculate the Atomic Mass of a Single Atom
Step 1. Find the atomic number of the element or isotope
This corresponds to the number of protons found in the element and never varies. For example, all hydrogen atoms and only hydrogen atoms have a proton in their nucleus. Sodium has an atomic number of 11 because there are eleven protons in its nucleus, while the atomic number of oxygen is 8 because its nucleus is made up of 8 protons. You can find this data in almost all the standard periodic tables: you see it above the chemical symbol of the element. This value is always a positive integer.
- Consider the carbon atom. This always has six protons, so you know its atomic number is 6. On the periodic table you can also read a small number "6" above the element symbol inside the carbon box (C); this indicates its atomic number.
- Remember that the element's atomic number has no direct bearing on the relative atomic mass value indicated on the periodic table. Despite this, you may get the impression that the atomic mass is double the atomic number, especially for the elements found at the top of the periodic table, but be aware that the atomic mass is never calculated by doubling the atomic number.
Step 2. Find the number of neutrons that make up the nucleus
This can vary between the atoms of a given element. Although two atoms with the same number of protons and a different number of neutrons are always the same "element", they are actually two different isotopes. Unlike the number of protons, which is constant, the number of neutrons in a given atom can change to such an extent that the average atomic mass must be expressed as a decimal value between two integers.
- The number of neutrons is determined by how the isotope has been designated. For example, carbon-14 is a naturally occurring radioactive isotope of carbon-12. Often the isotope is indicated with a superscript number preceding the element symbol: 14C. The number of neutrons is calculated by subtracting the number of protons from the isotope number: 14 - 6 = 8 neutrons.
- Suppose the carbon atom you are considering has six neutrons (12C). This is the most common isotope of carbon and accounts for 99% of the existing carbon atoms. However, about 1% of carbon atoms have 7 neutrons (13C). The other types of carbon atoms with less than 6 or 7 neutrons represent a very small amount.
Step 3. Add the number of protons and neutrons together
This is the atomic mass of the atom. Don't worry about the number of electrons orbiting the nucleus, the mass they generate is really very, very small so, in most practical cases, it doesn't interfere with the result.
- Your carbon atom has 6 protons + 6 neutrons = 12. The atomic mass of this specific atom is equal to 12. If you had considered the isotope carbon-13, then you should have calculated 6 protons + 7 neutrons = 13.
- The real atomic weight of carbon-13 is 13, 003355 and is obtained more precisely by experiment.
- The atomic mass is a value very close to the isotope number of an element. For basic calculations, the isotope number is assumed to be equal to the atomic mass. When a calculation is made experimentally, the atomic mass figure is slightly greater than the isotope number, due to the minimal contribution made by the electron mass.
Method 3 of 3: Calculate the Relative Atomic Mass (Atomic Weight) of an Element
Step 1. Determine which isotopes make up the sample
Chemists often determine the proportions between the various isotopes that make up a sample using a special instrument called a spectrometer. However, for a chemistry student, this information is mostly provided by the text of the problem or can be found as fixed data in textbooks.
For your purpose, consider a sample composed of the isotopes carbon-13 and carbon-12
Step 2. Determine the relative abundance of each isotope in the sample
For each element, isotopes are present with different proportions which are usually expressed as a percentage. Some isotopes are very common, while others are very rare, so much so that they can hardly be identified. You can find this through mass spectrometry or by consulting a chemistry book.
Suppose the abundance for carbon-12 is 99% and that of carbon-13 is 1%. Of course, there are other carbon isotopes, but in such small quantities that they can be ignored in this experiment
Step 3. Multiply the atomic mass of each isotope by the value of its proportion in the sample expressed as a decimal value
To convert a percentage to decimals, simply divide the number by 100. The sum of the proportions expressed in decimals of the various isotopes that make up a sample should always be equal to 1.
- Your sample contains carbon-12 and carbon-13. If carbon-12 represents 99% of the sample and carbon-13 represents 1%, multiply 12 (the atomic mass of carbon-12) by 0, 99 and 13 (the atomic mass of carbon-13) by 0, 01.
- A reference text will give you the percentage proportions of all isotopes of an element. You can usually find this data in the tables on the back pages of each chemistry book. Alternatively, you can use a mass spectrometer to test the sample directly.
Step 4. Add the results together
Add up the products of the multiplications you did earlier. The resulting value is the relative atomic mass of the element, i.e. the average value of the atomic masses of the element's isotopes. When we talk about an element in general without taking into consideration a particular isotope, this data is used.
