7 min read•Last Updated on June 18, 2024
Dalia Savy
Jeremy Kiggundu
Dalia Savy
Jeremy Kiggundu
When taking a look at the periodic table, you may wonder how scientists discovered all these numbers! As we discussed in the last section of this unit, an atom is made up of protons, neutrons, and electrons. Understanding these three subatomic particles and how they contribute to the makeup of the periodic table is significant.
On the periodic table, you can find the following information for each element:
Let's take a look at carbon on the periodic table that will be given to you during the AP exam:
Here is what we know so far about carbon based on the periodic table:
This is all because the atomic masses that you're given on the periodic table are actually the average atomic masses of these elements. The average atomic mass of an element is the weighted average of the masses of the naturally occurring isotopes of that element, based on their relative abundances.
🎥 Watch Jacob Jeffries discuss the parts of the atom and the experiments scientists use to study them.
Isotopes are variants of an element. They have the same number of protons and electrons, but a different number of neutrons. This means that isotopes have the same atomic number (number of protons) but a different atomic mass (total number of protons and neutrons).
Therefore, the average atomic mass represents all of the isotopes of an atom, and how often they occur naturally in the environment. Let's break this down further. 🤔
There are three naturally occurring isotopes of carbon: carbon-12, carbon-13, and carbon-14.
So how do we calculate the average atomic mass of 12.01 knowing this information about carbon's naturally occurring isotopes? Here is what we know so far:
Average atomic mass = (abundance of isotope 1 x mass of isotope 1) + (abundance of isotope 2 x mass of isotope 2) + ... +
To calculate the average atomic mass of carbon from this data:
AAM = 0.989(12) + 0.011(13) = 12.01, which is the same as the number on the periodic table.
Let's see what would happen if we changed the numbers. Let's say now the chemist finds that only 75% of the carbon in nature was carbon-12 and 25% was carbon-13. Now, the AAM is:
AAM = 0.75(12) + 0.25(13) = 12.25
As you can see, the abundance of isotopes can have a significant effect on the average atomic mass of an element. If you are only given the mass numbers 12 and 13, as well as the average atomic mass, you can easily tell which isotope exists in greater amounts in nature. Since 12.01 is much closer to 12 than 13 is, Carbon-12 is clearly more abundant.
** Always convert the % given to a decimal when calculating the average atomic mass.**
Mass spectrometry is a technique used to measure the mass and relative abundance of ions in a sample. This technique produces a graph, called the mass spectrum, that allows us to identify different isotopes of an element and the relative abundance of each isotope in nature.
Looking at the mass spectrum of the element carbon, we can identify the isotopes of carbon as carbon-12 and carbon-13.
The fact that carbon-12 gives a significantly greater signal tells us that carbon-12 is much more abundant than carbon-13 in nature.
Let's put all of this information together and try to identify an element just from its mass spectrum. Here is an example:
The peaks below show a mass spectrum of element X. What is element X based on this spectrum? What is its average atomic mass?
Without a calculator, you can easily tell that the average atomic mass is going to be somewhere between 24 and 25 since X-24 has the greatest % natural abundance. However, let's do some work! Here is what you should do when approaching this problem:
AAM = 0.828(24) + 0.081(25) + 0.091(26) = 24.263. This tells us that element X has an average atomic mass of approximately 24.3. When looking at the periodic table, we can identify this element as Magnesium:
Part a of question 2 on the AP Chemistry 2007 Exam (Form B) includes basic stoichiometry and advanced mass spectrum calculations. Here it is:
Instead of giving you the graph and asking you to solve for the average atomic mass, they are giving you the average atomic mass and asking you to solve for the percent abundance.
Don't worry! As long as you remember the basic formula, you can solve this problem:
AAM = (natural abundance)(mass) + ... + (natural abundance)(mass)
Now, just plug in what you know!
20.18 = (natural abundance)(19.99) + (natural abundance)(21.99)
One crucial piece of information that you have to remember for this question is that the percent abundances add up to 100%. This allows us to have x as one natural abundance and 1-x as the other, as you have probably done in previous algebra classes. Therefore, the setup for this problem is:
20.18 = (x)(19.99) + (1-x)(21.99)
Once you solve for x, you would get x = 0.905. Therefore, the percent abundance for Ne-20 is 90.5% and the percent abundance for Ne-22 is 9.5% (100%-90.5%).
To recall from the last guide, you always want to put the value you know first when doing dimensional analysis. Well, how do you even know you have to do stoichiometry? This question is asking you to convert from grams to atoms which should give you a signal that dimensional analysis should be performed. Try this problem on your own first but here is the setup:
In the first conversion step, you are simply using the molar mass of neon on the periodic table. In the second conversion step, you are using the ratio of Ne to Ne-22, which is the percent abundance you found in part i. Then, you use Avogadro's number to convert to atoms and you get the answer of 3.558 x 10^22 Ne-22 atoms.
Mass Spectroscopy is more of a lesson that is geared towards the laboratory aspect of Chemistry🧪, but nonetheless, here’s an excellent lesson on Khan Academy which reviews mass spectroscopy:
Abundance of isotopes refers to how much an isotope contributes relative to all the isotopes of that element in a given sample.
Term 1 of 15
Abundance of isotopes refers to how much an isotope contributes relative to all the isotopes of that element in a given sample.
Term 1 of 15
Abundance of isotopes refers to how much an isotope contributes relative to all the isotopes of that element in a given sample.
Term 1 of 15
The periodic table is a tabular arrangement of chemical elements, organized based on their atomic number, electron configuration, and recurring chemical properties. Elements are listed in order of increasing atomic number.
Atomic Number: This is the number of protons found in the nucleus of an atom. It defines the identity of an element.
Element Symbol: This is a one or two-letter abbreviation that represents an element on the periodic table. For example, "H" stands for Hydrogen and "O" stands for Oxygen.
Group (Periodic Table): A column in the periodic table that contains elements with similar properties due to having the same number of valence electrons.
An element symbol is a one or two-letter code used to represent a chemical element. It's usually derived from either current English name or Latin name of an element.
Chemical Equation: A representation using symbols and formulas to illustrate what happens during a chemical reaction.
Compound Formula: A formula that shows which atoms and how many atoms are present in a molecule of compound.
Isotope Symbols: These are representations that indicate both atomic number and mass number for specific isotopes.
The atomic number is equal to the number protons found in an atom's nucleus. It defines what element an atom is and its place on the periodic table.
Protons: Positively charged particles found in the nucleus of an atom.
Electrons: Negatively charged particles that orbit around the nucleus of an atom.
Neutrons: Neutral particles (no charge) found in the nucleus of an atom.
The atomic mass is the total weight of protons, neutrons, and electrons in a single atom when at rest.
Proton: A subatomic particle with a positive charge that resides in an atom's nucleus.
Neutron: A subatomic particle with no charge (neutral) that also resides in an atom's nucleus.
Electron: A subatomic particle with a negative charge that orbits around the nucleus of an atom.
The average atomic mass is the weighted average of all naturally occurring isotopes of an element. It takes into account both the mass and abundance (% occurrence) of each isotope.
Weighted Average: An average where each value has a different level of importance or frequency.
Abundance: The relative proportion or quantity of a particular species present in a sample.
Element: A substance consisting entirely from atoms with the same number of protons.
Isotopes are versions of an element that have different numbers of neutrons but the same number of protons. They have identical chemical properties but different physical properties due to their varying masses.
Radioisotopes: Isotopes that are unstable and undergo radioactive decay.
Stable Isotope: An isotope that remains unchanged over time because it does not undergo radioactive decay.
Isotope Abundance: The percentage of a particular isotope in a sample of an element.
Carbon-12 is the most common isotope of carbon, with 6 protons and 6 neutrons in its nucleus.
Isotopes: Variants of a particular chemical element which differ in neutron number, and consequently in nucleon number.
Atomic Number: The number of protons found in the nucleus of an atom, which determines the chemical properties of an element and its place in the periodic table.
Nucleus: The central part of an atom that contains protons and neutrons.
Carbon-13 is a naturally occurring isotope of carbon, with 6 protons and 7 neutrons. It makes up about 1.1% of all natural carbon on Earth.
Stable Isotope: An isotope that has not observed to decay into other elements.
Natural Abundance: The measure of the average amount of a given isotope naturally occurring on Earth.
Neutron Number: The number of neutrons in an atomic nucleus.
Carbon-14 (C14) is a radioactive isotope of carbon with 6 protons and 8 neutrons. It's used for radiocarbon dating to determine the age of old artifacts or ancient remains.
Radioactive Decay: The process by which an unstable atomic nucleus loses energy by radiation.
Radiocarbon Dating: A method for determining the age of an object containing organic material by using the properties of radiocarbon, a radioactive isotope of carbon.
Half-Life: The time required for half the atoms in a sample of a radioactive isotope to decay.
Abundance of isotopes refers to how much an isotope contributes relative to all the isotopes of that element in a given sample.
Relative Atomic Mass: This term refers to the calculation weighted by the abundances of naturally occurring isotopes that form a particular element.
Isotopic Ratio: The ratio between the amounts of two isotopes in a given sample.
Natural Abundance: Refers to the occurrence, percentage wise, of isotopes in nature.
A mass spectrum is a plot of ion abundance versus mass-to-charge ratio generated from a mass spectrometer.
Spectrometer: A device used for measuring spectra, especially as produced by dispersion of light and other radiation into components according to frequency or wavelength.
Fragmentation Pattern: In mass spectrometry, fragmentation patterns can provide clues about the structure and identity of a molecule in your sample.
Base Peak: In a mass spectrum, the base peak is the peak representing the most abundant ion.
Avogadro's number, also known as Avogadro's constant, represents the number of atoms or molecules in one mole of any substance. It’s approximately 6.022 x 10^23 particles per mole.
Mole: A unit in chemistry representing Avogadro's number worth of particles (atoms, molecules, ions etc.)
Atoms/Molecules/Ions: These are basic units that make up matter; atoms form molecules and ions are charged atoms or groups of atoms.
Stoichiometry: The part of chemistry dealing with quantitative relationships between reactants and products in chemical reactions; often involves calculations using moles and Avogadro’s number.
Mass spectroscopy is an analytical technique that measures the mass-to-charge ratio of charged particles. It is used to identify the amount and type of chemicals present in a sample.
Isotope: An isotope is a variant of a particular chemical element, which differs in neutron number.
Ionization: Ionization refers to any process by which electrically neutral atoms or molecules are converted to electrically charged atoms or molecules (ions).
Mass-to-Charge Ratio: This term refers to the physical quantity that is most directly accessible in mass spectrometry. It represents the strength of interaction between an ion and an electromagnetic field.