Atomic Orbitals By PresenterMedia.com PresenterMedia.com

Objectives Discuss Valence Electrons Item 1 DIscuss Item 2 Item 3 Item 4 Discuss # and type of orbitals in each energy level

Aufbau’s Principle When filling energy levels, fill lowest energy levels first (exceptions to this rule we will talk about later) Pauli Exclusion Principle Each orbital can hold at most 2 electrons Hund’s Rule All unfilled orbitals of the same shape on an energy level will receive one electron before any orbital will take a 2 nd electron (with an opposite spin) Principles for filling energy shells Fill empty orbitals first before adding a 2 nd electron Start closest to nucleus and fill out.

Filling in Electrons S=2 electrons P=6 electrons D=10 electrons F=14 electrons G=18 electrons

Electron Configuration Problems 1. Ne 2. Na 3. S 4. Ar 5. Fe 6. Br

The Shortcut Notice the difference between Ne and Na is only 3s 1 electron. We can write this in a shorthand way by using the noble gas in the period above and then filling in the electrons for the period we are in. For example: [Ne] 3s 1 = Na [Ar] 4s 1 = K [Ar] 4s 1 4s 2 = Ca [Ar] 4s 1 4s 2 3d 1 = Sc

Valence Electrons Only the electrons in the outermost energy level ( should be the s and p orbitals) Never more than 8 Compare this number to the group number (family) the element is in Lewis Dot Structure Shows Element surrounded by up to 4 electron pairs based on valence electrons.

Objectives Discuss Valence Electrons Summarize and apply trends in the periodic table Item 1 Discuss Lewis Electron Structure Item 2 Item 3 Item 4 Apply Lewis Electron Structure to bonding

Even though technically Carbon should have a filled s orbital and 2 half filled p orbitals, it behaves as if it has 4 half filled orbitals

Lewis Dot Structure Easy, All you need to know are two things. 1. # of valence electrons 2. Hund’s rule 1. Electrons fill unoccupied orbitals before filling an occupied orbital

Combining Lewis Electron Structure We will get into how atoms bond to each other in the future, but for now I want you to see how Lewis Dot Structure can help us see how bonding works. What you need to now is that when atoms bond, they need to fill up the empty spaces. They do this by giving, taking, or even sharing electrons.

Octet Rule Atoms are generally most stable when the valence shell is filled with 8 electrons. This means that atoms with less than 8 electrons in its outer shell will either give, take or share electrons with other atoms until the outer electron shell is full. In truth, there is some stability in being half filled (every valence orbital having only 1 electron, but this is nowhere near as stable as being full)

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Objectives Discuss Valence Electrons Summarize and apply trends in the periodic table Item 1 Discuss Lewis Electron Structure Item 2 Item 3 Item 4 Apply Lewis Electron Structure to bonding

Periodic Trends We know that there are trends in the periodic table. Metals, non metals, metalloids, gases, etc. But there are other trends that are based on how the atoms look. One note, transition metals have a more complicated arrangement of electrons which can affect some of the trends.

Atomic Radii How big the atom is Based on size of the nucleus (greater nuclear attraction) and Based on # of energy levels Greater distance from nucleus

Electronegativity Ability for an atom to steal an electron from another atom. Based on how many electrons needed to fill the shell and Based on # of energy levels Electrons between the nucleus and the valence shell can shield those electrons from the attractive forces

Bellringer – Chemistry Agenda Bellringer Finish Periodic Trends Website Activity Covalent Bonding (multiple days) Clear Targets I can use the periodic table to detect trends in elements Atomic radii Electronegativity Electron affinity Ionization energies I can compare/contrast covalent bonding I can relate single-double-triple bonds to bond strength and bond length Paper Due 15 th.

Electron Affinity Very similar to electronegativity. It is how much energy is released when an atom absorbs an electron to be come more stable Makes a negative ion - anion Easier when the atom is small (less shielding) Easier when the atom has a nearly full valence shell (closer to full than empty) Some fluctuation

Bellringer – Chemistry Agenda Bellringer Finish Periodic Trends Website Activity Covalent Bonding (multiple days) Clear Targets I can use the periodic table to detect trends in elements Atomic radii Electronegativity Electron affinity Ionization energies I can compare/contrast covalent bonding I can relate single-double-triple bonds to bond strength and bond length Paper Due 15 th.

Ionization energy Basically the opposite of electronegativity. It is how easy it is to take an electron away from an atom. (make positive ion – cation) Energy required to take electron away Easier when the atom is large (shielding) Easier when the atom has only a couple of electrons (closer to empty than full)

Bellringer – Chemistry Agenda Bellringer Finish Periodic Trends Website Activity Covalent Bonding (multiple days) Clear Targets I can use the periodic table to detect trends in elements Atomic radii Electronegativity Electron affinity Ionization energies I can compare/contrast covalent bonding I can relate single-double-triple bonds to bond strength and bond length Paper Due 15 th.

Check your understanding 1.) Based on the periodic trends for ionization energy, which do you expect to have the highest ionization energy? A.) Fluorine (F) B.) Nitrogen (N) C.) Helium (He) 2.) Nitrogen has a larger atomic radius than Oxygen. A.) True B.) False 5.) Which element do you expect to be more electronegative, sulfur (S) or selenium (Se)? 6) Why is the electronegativity value of most noble gases equal to zero? 8) Rewrite the following list in order of decreasing electronegativity: Fluorine (F), Phosphorous (P), Sulfur (S), Boron (B). 9) An atom with an atomic radius smaller than that of Sulfur (S) is __________. A.) Oxygen (O) B.) Chlorine (Cl) C.) Calcium (Ca) D.) Lithium (Li) E.) None of the above 10) A nonmetal will have a smaller ionic radius when compared to a metal of the same period. A.) True B.) False

Objectives Discuss Valence Electrons Summarize and apply trends in the periodic table Item 1 Discuss Lewis Electron Structure Item 2 Item 3 Item 4 Apply Lewis Electron Structure to bonding