Nuclear Notation

Need help preparing for the General Chemistry section of the MCAT? MedSchoolCoach expert, Ken Tao, will teach everything you need to know about Nuclear Notation for Atomic Structure. Watch this video to get all the MCAT study tips you need to do well on this section of the exam! In chemistry, it is essential that scientists can quickly and reliably identify different atomic compounds and molecules. To facilitate this, chemists have constructed a specific method of displaying atoms and molecules known as nuclear notation. Nuclear notation is a convenient and accessible way to quickly visualize an atom’s identity, including its charge and the structure of its nucleus. Chemical Identity In this notation, X is the chemical symbol, representing its elemental identity as shown in the periodic table of elements. This is often a single letter, such as C for carbon, but can also consist of two letters, such as the He used for helium. While the chemical symbol for most elements is fairly intuitive, some symbols are less clear, like Pb for lead, typically because these symbols were derived from the Latin names for those elements. Memorizing every less intuitive chemical symbol would be both time consuming and unnecessary, as the periodic table will be provided on the MCAT, but it is helpful to be familiar with some of the key elements that are prevalent in organic chemistry and biochemistry, particularly sodium (Na) and potassium (K). Atomic Number The Z in the notation above stands for the atomic number of the element, or the number of protons present in its nucleus. Each proton has a positive charge of 1.6 x 10-19 coulombs and a mass of approximately 1 atomic mass unit (amu). This charge is the equal and opposite charge of the electron, and so is commonly known as the elementary charge. A proton, then, can be said to have a charge of +1e (one positive elementary charge), or simply +1. As the proton is the only charged nucleon (neutrons having a neutral charge), the atomic number is equal to the nuclear charge. The atomic number is an essential value due to its identification role: each element is named for the number of protons, or its atomic number. That is, all atoms of carbon have 6 protons, all atoms of nitrogen have 7 protons, etc. Any change to the number of protons in an atom’s nucleus, such as by nuclear fusion and fission, or radioactive decay, will result in the creation of an entirely new element with a correspondingly higher or lower atomic number. Mass Number A, in the nuclear notation, is the mass number, or the total number of protons and neutrons. Remember that neutrons have no net charge, but are approximately the same size (one amu) as protons. While every element of the same kind has the same number of protons, the element can exist in different states, known as isotopes, each of which have the same number of protons but a different number of neutrons, and thus a different mass number. Consider carbon as an example. Most carbon on earth exists as carbon-12, the primary isotope that consists of 6 protons and 6 neutrons. However, a lesser amount of carbon also exists as carbon-13 and carbon-14, which have 7 and 8 neutrons, respectively. Isotopes tend to have significantly different stabilities and half-lives, which play a major role in our understanding of radiation and nuclear chemistry, and helps explain why many elements predominantly exist as a single isotope in nature. Now, because most elements can exist in a variety of isotopic states, this posed a question for the creators of the periodic table: which isotopes should be used to represent each element in the table? A rather elegant solution was devised, by which each element would be displayed with its chemical symbol, its atomic number, and a mass number that was a weighted average of its isotopes by their abundance in nature. Let us consider carbon, for example. Approximately 98.9% of all the carbon on earth is found as carbon-12, which has a total atomic weight of approximately 12 amu, the most abundant and most stable isotope. Most of the remaining 1.1% of carbon is carbon-13, which has an atomic mass of approximately 13 amu. By averaging these relative to their abundance via the basic equation of mass of isotope 1 x percent abundance of isotope 1 + mass of isotope 2 x percent abundance of isotope 2, and then dividing the answer by 2, we are left with an average of approximately 12.011, the weighted average atomic mass of carbon. MEDSCHOOLCOACH To watch more MCAT video tutorials like this and have access to study scheduling, progress tracking, flashcard and question bank, download MCAT Prep by MedSchoolCoach IOS Link: https://play.google.com/store/apps/details?id=com.htd.medschoolcoach&hl=en_US Apple Link: https://apps.apple.com/us/app/mcat-prep-by-medschoolcoach/id1503000883 #medschoolcoach #MCATprep #MCATstudytools

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