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Atomic Structure

This resource provides a comprehensive overview of atomic structure, focusing on the atomic number, isotopes, and atomic mass. The atomic number represents the number of protons in an atom and helps identify the element. Isotopes are defined as atoms with the same number of protons but different neutron counts. The resource also explores mass number, atomic mass calculations, and concepts such as mass defect and relative abundance. Practical exercises enhance understanding by applying concepts related to isotopes and atomic composition.

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Atomic Structure

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  1. Atomic Structure

  2. Atomic Number • Definition: Equals the number of protons • Identifies the element • # protons = # electrons in a neutral atom (an atom without a charge)

  3. Isotopes & Mass Number • Isotopes = atoms with the same # of protons, but different numbers of neutrons

  4. Ex: Isotopes 3 isotopes of hydrogen: • Hydrogen-1; mass # = 1; 0 neutrons • Hydrogen-2 (deuterium); mass # = 2; 1 neutron • Hydrogen-3 (tritium); mass # = 3; 2 neutrons

  5. There are two different ways to write chemical symbols for isotopes: • Write the mass number after the element’s name (gold-197) • Use the symbol, a superscript for mass #, and a subscript for atomic #

  6. Mass Number = protons + neutrons • To calculate # neutrons = mass # - atomic # • Use shorthand writing of atom’s composition • Superscript = mass # • Subscript = atomic #

  7. Practice 1

  8. Practice 2

  9. Practice 3 • Three isotopes of sulfur are sulfur-32, sulfur-33, and sulfur-34. Write the complete symbol for each isotope, including the atomic number and the mass number. • How many neutrons, protons, and electrons are in Na+ with a mass number of 24? What is its atomic number?

  10. Atomic Mass • Weighted atomic mass of all the isotopes of that element – that is why it is a decimal, not an integer • It reflects mass and relative abundance of the isotopes as they occur naturally

  11. AMU – the mass of an atom • Amu’s are used to define the weight of a single atom or isotope • How measured: Mass spectrometer – determines masses of atoms • Carbon-12 was set to 12 atomic mass units (amu’s) • Used to define masses of atoms

  12. Mass Number and Atomic Mass Carbon-12 = 6 protons and 6 neutrons, • mass # = 12, • atomic mass = 12 amu, (atomic # = 6) • 1 proton = 1 amu • 1 neutron = 1 amu • Mass number = Atomic Mass FOR A SINGLE ATOM ONLY • The defined mass of each elements is actually a weighted average so it is NOT a whole number.

  13. Average Atomic Mass: Uses a weighted average to account for mass and relative amounts • Most elements are present as a mixture of 2+ isotopes • Not all isotopes are present in equal amount (percent abundance) • Ex – hydrogen

  14. Mass defect: the mass of any nucleus is less than the sum of the separate masses of its protons and neutrons. • How can the actual mass be less than the mass number? Can atoms lose mass? • E=mc2... Some of the mass of the protons and neutrons get converted to energy when they come together in the nucleus

  15. Calculating average atomic mass • Atomic mass = (Mass isotope A * natural abundance of A) + (Mass isotope B * natural abundance of B) + … • You must know the following in order to calculate the atomic mass of an element: • # of stable isotopes • mass of each isotope • % abundance of each isotope

  16. Chlorine: Example 1 • Use the data from the previous slide on chlorine to determine the average atomic mass.

  17. Magnesium Example

  18. Calculating Relative Abundance from Average Atomic Mass • The element Quahog has two isotopes: Quahog-35 and Quahog-40. • If the atomic mass of Quahog is 39.00 amu, what is the relative abundance of each isotope?

  19. The sum of the relative abundance must equal 1. • So assign X and 1-X as the relative abundance of each and solve.

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