Unit 3: Atomic Structure

1. Unit 3: Atomic Structure Chapter 3

2. This tutorial is designed to help students understand scientific measurements. • Objectives for this unit appear on the next slide. • Each objective is linked to its description. • Select the number at the front of the slide to go directly to its description. • Throughout the tutorial, key words will be defined. • Select the word to see its definition.

3. Objectives 3.1 understand the historical progression of structure of the atom including such models created by Democritus, Dalton, Thomson, Rutherford, Bohr, and Schrodinger 3.2 define and identify examples of the laws of conservation of mass, definite proportions, and multiple proportions 3.3 identify the structure of an atom • state the relative masses of a proton, neutron, and electron, their relative charges, and locations in the atom • define and identify isotopes and ions • define and distinguish between atomic number and mass number • identify the parts of the nuclear symbol 3.4 define and calculate average atomic mass in amu’s 3.5understand and record the arrangement of electrons in an atom • state and use Hund’s rule, the Pauli exclusion principle, and the Aufbau principle to describe the order in which electrons fill energy levels and their sublevels • write the orbital notation for specific elements • write the electron configuration for specific elements

4. 3.1 History of the atom

5. 3.2 Define and Identify Three Laws • Law of Conservation of Mass • The law of conservation of mass states that mass can be mass can neither be created or destroyed. • This means that the products of a chemical reaction will have the same mass as the reactants.

6. 3.2 Define and Identify Three Laws • Law of Definite Proportions • This law states that any sample of a compound has the same composition. • This means that water will always be H2O whether it is found in Iowa or somewhere else.

7. 3.2 Define and Identify Three Laws • Law of Multiple Proportions • This law states that the mass ratio for one of the elements in a compound that combines with a fixed mass of another element can be expressed in small whole numbers.

8. 3.3 Identify the structure of the atom • The atom is constructed of three basic subatomic particles. • In the last 20 years, it has been discovered that quarks make up both protons and neutrons.

9. Structure of the Atom-Bohr Model

10. Isotopes and Ions • Each atom can have different variations. • All atoms are identified by the number of protons they contain. • Oxygen will always have 8 protons. If you added an additional proton, the atom would no longer be oxygen (it would be fluorine). • The number of neutrons and electrons can vary slightly.

11. Isotopes-changes in neutrons • Protons contain a positive charge. When an atom becomes large, it contains several protons. That much positive charge in one location is unstable. • Neutrons, which have no charge, act as spacers in between the protons. • An isotope is the name of an atom with different amounts of neutrons. • Oxygen for example has three common isotopes. • Oxygen-16 has 8 neutrons • Oxygen-17 has 9 neutrons • Oxygen-18 has 10 neutrons

12. Ions-changes in electrons • Electrons fill the orbitals surrounding the nucleus. • This is known as the electron cloud. • In a neutral atom, each proton has an electron in the electron cloud. • When an atom becomes charged, the number of electrons no longer matches the protons. • Positive charges indicate a loss of an electron. • Negative charges indicate a gain of an electron.

13. Nuclear Symbols • Nuclear symbols are useful for determining the amount of electrons, protons, and neutrons in an element. Mass number: Shows the number of neutrons and the number of protons. Charge: Indicates the difference between the electrons and protons. Negative charges indicate more electrons while positive charges indicate more protons. Atomic Number: Shows the number of protons.

14. 3.4 Average atomic mass • On the Periodic Table, the mass listed is the average atomic mass. • This is an average of all the naturally occurring isotopes of an atom. • Calculating an average for a large amount of particles can be challenging.

15. Average Atomic Mass • Carbon has three common isotopes. • Carbon-12 • Carbon-13 • Carbon-14 • These three common isotopes do not come in equal amounts though: • 98.89% is carbon-12 • 1.10% is carbon-13 • 0.01% is carbon-14

16. Average atomic mass • To determine the average atomic mass, the following formula should be used: (Atomic Mass x Percent Abundance) + (Atomic Mass x Percent Abundance) +…..=average atomic mass • So in the case of carbon: (12.00 amu x .9889) + (13.00 amu x 0.0111) + (14.00 amu x 0.0001) = 12.01 amu

17. 3.5 Arrangement of Electrons • According to Bohr and Schrodinger, electrons surround the atom in energy levels that take the form of orbitals. • The Periodic Table is broken into four sections that correlate to the orbitals • The placement of electrons within these orbitals follow set rules.

18. Arrangement of Electrons • Hund’s Rule • Electrons will fill the available orbitals at a certain energy level before pairing. • Pauli Exclusion Principle • Only two electrons can occupy each orbital and their spins will be opposite. • Aufbau Principle • The lowest-energy orbitals will be filled first.

19. Below is a depiction of how the Periodic Table shows the orbitals (SPDF). S P F D

20. Energies of Each Level • The orbitals correspond to different levels of energies. • In general, each row on the Periodic Table represents a different level of energy. • It gets more complicated farther down the Periodic Table.

21. Energies of Each Level • 7p ___ ___ ___ • 6d ___ ___ ___ ___ ___ • 5f ___ ___ ___ ___ ___ ___ ___ • 7s ___ • 6p ___ ___ ___ • 5d ___ ___ ___ ___ ___ • 4f ___ ___ ___ ___ ___ ___ ___ • 6s ___ • 5p ___ ___ ___ • 4d ___ ___ ___ ___ ___ • 5s ___ • 4p ___ • 3d ___ ___ ___ ___ ___ • 4s ___ • 3p ___ ___ ___ • 3s ___ • 2p ___ ___ ___ • 2s ___ • 1s ___ As you progress from the 1s to the 7p, you increase the amount of energy. Notice how the d-sublevel is always after the s-sublevel of the previous energy level (example 3d follows 4s). Notice the f-sublevel follows the s-sublevel of two energy levels before it (example: 4f follows 6s). Energy

22. Orbital Notation • To show orbital notation, the three rules must be followed. • Therefore, start at the lowest energy level. • Designate an electron by drawing an arrow • The arrow indicates the spin • Place one electron in each orbital until they each have one on that level. Then go back and pair them. • Only two electrons fit in each orbital. • The arrows must point in opposite directions to show opposite spins.

23. Orbital Notation • Look at the element sulfur-32. • First determine the number of electrons • Since sulfur has an atomic number of 16, it has 16 protons, 16 electrons, and 16 neutrons. 3p ___ ___ ___ 3s ___ 2p ___ ___ ___ 2s ___ 1s ___ Place electrons in the first energy level and continue up. When 2p is hit, fill each orbital and then go back and pair. Finish by repeating the process used for 2p but stop when you hit 16 arrows. This is the orbital notation.

24. Electron Configuration • Orbital notation can take up a lot of space. • It does a nice job of giving a visual of the location of each electron. • Electron configuration is a shorthand notation for determining the location of the electrons. • It follows the same rules but is slightly easier to write.

25. Electron Configuration • As you read across the Periodic Table, you can pick out the electron configuration. • The electron configuration for oxygen-16 is: O:1s2 2s2 2p4 • Red represents the s-sublevel and yellow is the p-sublevel. • Each row is an energy level. • Since oxygen has eight electrons, we count eight boxes. The superscripts on the sublevels indicate the number of electrons.

26. Electron Configuration • The same rules apply to the d and f sublevels. • Example: Gold-196 (79 electrons) Au: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d9

27. This concludes the tutorial on measurements. • To try some practice problems, click here. • To return to the objective page, click here. • To exit the tutorial, hit escape.

28. Dalton’s Atomic Theory • Elements are made of tiny particles called atoms. • Atoms of an element are different from other elements and can be distinguished by their atomic masses. • All atoms of a given element must be identical in properties. • Atoms of an element can combine with atoms of a different element to for compounds. • Chemical reactions rearrange the atoms but cannot create or destroy atoms. Return

29. Thomson’s Plum Pudding Model • As that Thomson discovered the electron, that meant the atom contained smaller parts. • This would change the model used for the atom. • The model used to be a solid sphere • The plum pudding model used electrons as the “plums” and the rest of the atom as the “pudding”. • The plums were negative and the pudding was positive. Return

30. Gold Foil Experiment • The gold foil experiment was conducted by Ernest Rutherford and his graduate students, Hans Geiger and Robert Marsden. • By shooting alpha particles at a gold foil, they noticed the particles essentially went straight through. • This led them to conclude the atom was mostly empty space with a dense positive core (protons. Return

31. Bohr Model of the Atom • Niels Bohr created an atomic model after doing work with the color spectra emitted from a hydrogen atom. • His model is sometimes called the solar system model. • The following link provides a more detailed description (the first 15 slides covers Bohr): http://science.sbcc.edu/physics/solar/sciencesegment/bohratom.swf Return

32. Schrodinger’s Orbitals • The current model of the atom uses the orbitals discovered by Schrodinger. • Within each energy level, there exists four kinds of orbitals: s, p, d, and f. • Each can hold a certain number of electrons. • The shape of each orbital is shown on the next slide. • The image was taken from: http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electronic_Orbitals on July 28th, 2011.

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