Quantum numbers to Periodic Table

1. Quantum numbers to Periodic Table

2. Aufbau Principle • As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to orbitals • The lowest energy level is the 1 s orbital • From this we can begin to diagram the orbitals in an atom + nucleus electrons • 1 s __ • 2 s __ p __ __ __ • 3 s __ p __ __ __ d __ __ __ __ __

3. Hund’s Rule • The lowest energy configuration (ground state) for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli Exclusion Principle (two electrons per orbital). • The convention is to represent the spins of electrons as up or down arrows

4. Periodic Table • The modern design of the periodic table is built off of these orbital diagrams

5. You can determine the orbital by the placement For these elements the last electrons are filling s orbitals for these elements the last electrons are filling p orbitals For these elements the last electrons are filling d orbitals For the elements the last electrons are filling f orbitals

7. Electrons configurations are written to shorten orbital diagrams • The electron configuration for potassium is 1s22s2p63s2p64s1 • This can be abbreviated to [Ar]4s1 • This is the same as saying everything that is in Ar, 1s22s2p63s2p6 +4s1

8. Total order • 1 s __ • 2 s __ p __ __ __ • 3 s __ p __ __ __ d __ __ __ __ __ • 4 s __ p __ __ __ d __ __ __ __ __ f__ __ __ __ __ __ __ • 5 s __ p __ __ __ d __ __ __ __ __ f__ __ __ __ __ __ __ • 6 s __ p __ __ __ d __ __ __ __ __ • 7 s __ p __ __ __ 1 3 2 7 4 5 6 8 13 10 9 11 14 17 12 15 18 16 19

9. Aufbau Principle is the order of filling the orbitals • The (n + 1)s orbitals always fill before the nd orbitals, explained by the penetration effect. • After lanthanum, which has the configuration [Xe]6s25d1, a series of 14 elements called the lanthanide series, or the lanthanides, occurs. • After actinium, which has the configuration [Rn]7s26d1, a series of 14 elements called the actinide series, or the actinides, occurs.

10. Lone electron in the 5d or 6 d • Odd exception • Lanthanum’s last electron fills the 5d orbital. • Ce has its last electron fill in the 4f, and the one from 5d (La) jumps to 4f. • The next elements all fill up the 4f until Gadolinium which fills up the 5d again. • A similar “jumping” of one electron occurs with the actinide series

11. Different Periodic Tables • The electron “jumping” is why there are two forms of the periodic table in common usage. • One has lanthanum and actinium in with the d filling elements. • The other has lanthanum and actinium with the f filling elements

13. To Determine the Condensed electron configuration • Start at the previous noble gas (lower atomic number) • Write out every element in order of atomic number to your given • Determine which shell all elements are filling. The number of elements in that shell is your superscript number.

14. For Iron • Previous Noble Gas----Ar • Write all elements • K CaSc Ti V Cr Mn Fe • These fill 4 s • These fill 3 d • So it is • [Ar]3d64s2

15. Practice • Write out the condensed electron configuration and orbital diagram for… • Nitrogen • Tin • Molybdenum • Rutherfordium

16. Chemical Bonds • Bonds are forces that hold groups of atoms together and make them function as a unit. • Bond energy, the energy required to break a bond, provides information regarding the strength of bonding interactions, which in turn indicates the “type” of bond. • *It always requires energy to break a bond, activation energy, It energy is always released when bonds are formed.

17. Bonds form to achieve the lowest possible energy state • A bond will form if the energy of the compound is lower in energy (more stable) than that of the separated atoms. • The system will act to minimize the sum of the positive (repulsive) energy terms and the negative (attractive) energy term. • The distance where the energy is minimum is called the bond length.

18. Bond Forces • simultaneous attraction of each electron by the protons generates a force that pulls the protons toward each other and that just balances the proton-proton and electron-electron repulsive forces at the distance corresponding to the bond length.

19. Types of Bonding • Metallic Bonding • Ionic Bonding • Covalent Bonding

20. Major differences • Covalent bonding is a sharing of electrons, Ionic bonding is a transfer of electrons, and in metallic bonds the electrons become community property. • Covalent bonds are stronger than ionic bonds because there is actually an electron going between them, it is a “true” bond. • Therefore, it is harder to break a covalent bond than it is to break an ionic bond. • Metallic and ionic bonds bond every atom in a substance to each other.

21. Shortcut to determining type of bond • When a metal and nonmetal bond you get an ionic bond • ~ something from the left excluding H bonds with something from the right = ionic bond. • When two nonmetals bond you get a covalent bond • ~things from the right bond with each other =covalent bond. • Metallic bonds only occur with the same metal atom

22. Metallic Bonding • All metal atoms in a substance share their valence electrons to form a “sea” of electrons, which is delocalized throughout the substance. • The electrostatic attractions between electrons and the metal ions account for the low melting points, but higher boiling points. • The metal is malleable as the lattice is shifted, but not repelled. • Metals are good conductors of electricity because of this electron movement.

23. Ionic Bonding • Electrons are donated from a cation to an anion. • Electrons lost are equal to electrons gained according to the octet rule. • Ionic bonding is not a “true bond” but a very strong attraction for opposing charges. • Ionic compounds tend to be hard, brittle, and rigid.

24. Covalent Bond • Electrons are shared between atoms. • Electrons are localized (fixed). • Nonpolar covalent bonding results from equal sharing of electrons. • Polar covalent bonding results form unequal sharing of electrons.

25. Electronegativity • Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. • Electronegativity difference is useful in determining bond type. • Electronegativity generally increases up and to the right, excluding the noble gases on the periodic table. • Fluorine (4.0) is the most electronegative element, meaning that in a molecule, fluorine always attracts electrons toward itself.

26. Electronegativity Chart

27. Electronegativity Difference • If the difference is between 0.0 and < 0.4, the compound is nonpolar covalent. • If the difference is between 0.4 and 1.9, the compound is polar covalent. • If the difference is greater than 1.9, the compound is ionic. • Order the following bonds according to polarity: H-H, O-H, Cl-H, S-H, and F-H.

28. Bonds

29. Why it is called polar polar implies different ends have different charges similar to a magnet. Water has 2 polar covalent bonds, meaning the electrons stay around oxygen more than H H Oxygen and this side positive That makes this side negative H

30. Denoting positive and negative Neither side is completely positive or negative, they are only partially positive and partially negative.  +  + H H Oxygen The symbol  (lower case delta) means partial  2-

31. Dipole Moment • Dipole moment- property of a molecule where the charge distribution can be represented by a center of a positive charge and a center of negative charge. • It is represented by this symbol • Positive center Negative center

32. So the dipole moment for water… is represented like this. Note the center of the positive charge is in between the two hydrogen atoms. H H Oxygen

33. Practice • For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl2, SO3, CH4, and H2S.

34. Molecular polarity • Not to be confused with bond polarity. • If the centers of both the positive and negative charges of the bond dipoles “averages” to the same point, the molecule is nonpolar. • A molecule can be nonpolar even though it is made up of polar bonds. • If the center of both the positive and negative charges of the bond dipoles “averages” to different points, the molecule is polar.