Atomic Structure

1. Atomic Structure

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

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

5. 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

6. 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 #

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

8. Practice 1

9. Practice 2

10. 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?

11. 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

12. 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

13. 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.

14. 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

15. 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

16. 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

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

19. Magnesium Example

21. 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?

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