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1. The way the real world works 63% +3% 46% 80% 97% 29%

2. http://slc.tamu.edu/ • Tutoring • Supplemental Instruction • Courses • Texas Success Initiative • About UsContact Us • 118 Hotard North of Sbisa, between • Neeley Hall and the Northside Post Office • (979) 845-2724 • The Student Learning Center has won the 2008 National College Learning Center Association • Frank L. Christ Outstanding Learning Center Award! • The award recognizes the center's commitment to supporting and strengthening the • Academic experience of students at Texas A&M University by providing a variety of programs and • services that promote retention and success. Read more... •   The Student Learning Center provides Supplemental Instruction and tutoring free of charge to all • Texas A&M University students. The SLC oversees the STLC courses (formerly CAEN), which teach • students how to improve their study skills and prepare for the job market. The SLC manages • Developmental programs for students who have not yet passed the assessment tests required by the state. • Study Tips • General • Time Management • Reading Textbooks • Setting Goals • Preparing for Exams • Success Tips from Fellow Aggies • Tutoring • During the Fall 2008 semester, drop-in tutoring will be offered Sunday nights 5-8pm and Monday through Thursday nights from 5-10pm. Tutoring will begin on Monday, September 1st. Tutor Zones are currently planned for Studio 12 of The Commons. Look for our table and tutors and just ask for help! See the schedule... Drop-in tutoring is available for most lower level math and science courses on the first floor of Hotard Hall. Tutors are also available to help out with many other courses. If you need help in a particular course and would like to check to see if a tutor is available for that course, you can contact our tutor coordinator, Linda Callen, at 845-2724.   SI Sessions: Monday, Wednesday, Sunday 8 PM Room 110 Harrington

3. Excited Atoms& Atomic Structure

4. The Quantum Mechanical Picture of the Atom Basic Postulates of Quantum Theory • Atoms and molecules can exist only in certain energy states. In each energy state, the atom or molecule has a definite energy. When an atom or molecule changes its energy state, it must emit or absorb just enough energy to bring it to the new energy state (the quantum condition). • Atoms or molecules emit or absorb radiation (light) as they change their energies. The frequency of the light emitted or absorbed is related to the energy change by a simple equation.

5. The Quantum Mechanical Picture of the Atom The allowed energy states of atoms and molecules can be described by sets of numbers called quantum numbers. • Quantum numbers are the solutions of the Schrodinger, Heisenberg & Dirac equations. • Four quantum numbers are necessary to describe energy states of electrons in atoms – n, , m, ms Schroedinger 3-dimensional time independent equation Heisenberg’s uncertainty Equation Dirac’s quantum mechanical model E. Schrodinger 1887-1961 W. Heisenberg 1901-1976

6. n n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7 • The principal quantum number has the symbol – n. • n = 1, 2, 3, 4, ...... “shells” • n = K, L, M, N, ...... • The electron’s energy depends principallyon n .

7.  • 2. The angular momentum quantum • number has the symbol. tells us the shape of the orbitals • and when linked with n defines the energy of the electron • = 0, 1, 2, 3, 4, 5, .......(n-1) • = s, p, d, f, g, h, .......(n-1) • = 0 • = s • = 1 • = p n = 1 n = 2 • = 2 • = d n = 3 n = 4 n = 5 n = 6 n = 7 • = 3 • = f

8. Quantum Numbers • The symbol for the magnetic quantum number is m. m = -  , (-  + 1), (-  +2), .....0, ......., ( -2), ( -1),  • If = 0 (or an s orbital), then m = 0. • Notice that there is only 1 value of m. This implies that there is one s orbital per n value. n  1 • If  = 1 (or a p orbital), then m = -1,0,+1. • There are 3 values of m. Thus there are three p orbitals per n value. n  2 • If  = 2 (or a d orbital), then m = -2,-1,0,+1,+2. • There are 5 values of m. Thus there are five d orbitals per n value. n  3 • If  = 3 (or an f orbital), then m = -3,-2,-1,0,+1,+2, +3. • There are 7 values of m. Thus there are seven f orbitals per n value, n • Theoretically, this series continues on to g,h,i, etc. orbitals. • Practically speaking atoms that have been discovered or made up to this point in time only have electrons in s, p, d, or f orbitals in their ground state configurations.

9. Atomic Orbitals • Atomic orbitals are regions of space where the probability of finding an electron about an atom is highest. • s orbital properties: • There is one s orbital per n level. •  = 0 • 1 value of m • = 0 • = s n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7

10. Atomic Orbitals • p orbitals are peanut or dumbbell shaped. • They are directed along the axes of a Cartesian coordinate system. • The first p orbitals appear in the n = 2 shell. • There are 3 p orbitals per n level. • The three orbitals are named px, py, pz. • They have an= 1. • m= -1,0,+1 3 values of m

11. Atomic Orbitals • d orbital properties: • The first d orbitals appear in the n = 3 shell. • The five d orbitals have two different shapes: • 4 are clover leaf shaped. • 1 is peanut shaped with a doughnut around it. • The orbitals lie directly on the Cartesian axes or are rotated 45o from the axes. • There are 5 d orbitals per n level. • The five orbitals are named • They have an  = 2. • m = -2,-1,0,+1,+2 5 values of m

12. Atomic Orbitals • d orbital shapes

13. Atomic Orbitals • f orbital properties: • The first f orbitals appear in the n = 4 shell. • The f orbitals have the most complex shapes. • There are seven f orbitals per n level. • The f orbitals have complicated names. • They have an= 3 • m = -3,-2,-1,0,+1,+2, +3 7 values of m • The f orbitals have important effects in the lanthanide and actinide elements.

14. Atomic Orbitals • f orbital shapes

15. Quantum Numbers • The last quantum number is the spin quantum number which has the symbol ms. • The spin quantum number only has two possible values. • ms = +1/2 or -1/2 • ms = ± 1/2 • This quantum number tells us the spin and orientation of the magnetic field of the electrons. • Wolfgang Pauli in 1925 discovered the Exclusion Principle. • No two electrons in an atom can have the same set of 4 quantum numbers. • Spin quantum number effects: • Every orbital can hold up to two electrons. • Consequence of the Pauli Exclusion Principle. • The two electrons are designated as having • one spin up ↑ and one spin down↓ • Spin describes the direction of the electron’s magnetic fields.

16. c (ms-1) l (m) n (s-1)x l (m) = c (ms-1) Frequency times wavelength equals the speed of light Calculate the frequency, the energy, and determine the color of light of an excited Li+ ion emitting radiation at 670.8nm. n (s-1) = E = n

17. The Periodic Table and Electron Configurations Note that the 3d subshell is higher in energy than the 4s subshell so appears in the 4th period

18. Quantum Numbers • The principal quantum number has the symbol ~ n which defines the energy of the shell • n = 1, 2, 3, 4, ...... “shells” The angular momentum quantum number has the symbol ~  which defines the subshells. •  = 0, 1, 2, 3, 4, 5, .......(n-1) •  = s, p, d, f, g, h, .......(n-1) n and l define the energy of the electron • The symbol for the magnetic quantum number is m which defines the orbital. • m = -  , (-  + 1), (-  +2), .....0, ......., ( -2), ( -1),  • The last quantum number is the spin quantum number which has the symbol ms which characterizes the single electron. • The spin quantum number only has two possible values. • ms = +1/2 or -1/2 one spin up ↑ and one spin down↓

20. The Nucleus: Build by adding the required number of protons (the atomic number) and neutrons (the mass of the atom) Electrons: Hund’s Rule states that each orbital will be filled singly before pairing begins. The singly filled orbitals will have a parallel spin. Pauli’s Exclusion Principle states that paired electrons in an orbital will have opposite spins. Fill the electrons in starting with the lowest energy level adhering to Hund’s and Pauli’s rules. Neon -

21. The Periodic Table and Electron Configurations • Now we can write a complete set of quantum numbers for all of the electrons in these three elements as examples. • Na • First for 11Na. • When completed there must be one set of 4 quantum numbers for each of the 11 electrons in Na (remember Ne has 10 electrons) [Ne] = 1s22s22p6

22. The Periodic Table and Electron Configurations 11Na 1s22s22p63s1

27. Electronegativity Increases Increases Increases

28. The Continuous Range of Bonding Types • Covalent and ionic bonding represent two extremes. • In pure covalent bonds electrons are equally shared by the atoms. • In pure ionic bonds electrons are completely lost or gained by one of the atoms. • Most compounds fall somewhere between these two extremes. • All bonds have some ionic and some covalent character. • For example, HI is about 17% ionic • The greater the electronegativity differences the more polar the bond.

29. Guidelines for assigning oxidation numbers. • The oxidation number of any free, uncombined element is zero. • The oxidation number of an element in a simple (monatomic) ion is the charge on the ion. • In the formula for any compound, the sum of the oxidation numbers of all elements in the compound is zero. • In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion. • Fluorine has an oxidation number of –1 in its compounds. • Hydrogen, H, has an oxidation number of +1 unless it is combined with metals, where it has the oxidation number -1. • Examples – LiH, BaH2 • Oxygen usually has the oxidation number -2. • Exceptions: • In peroxides O has oxidation number of –1. • Examples - H2O2, CaO2, Na2O2 • In OF2, O has oxidation number of +2. Use the periodic table to help with assigning oxidation numbers of other elements. • IA metals have oxidation numbers of +1. • IIA metals have oxidation numbers of +2. • IIIA metals have oxidation numbers of +3. • VA elements have oxidation numbers of –3 in binary compounds with H, metals or NH4+. • VIA elements below O have oxidation numbers of –2 in binary compounds with H, metals or NH4+.

30. OXIDATION NUMBERS NH3 ClO- H3PO4 MnO4- Cr2O72- C3H8

31. Formal Charges • The most likely formula for a molecule or ion is usually the one in which the formal charge on each atom is zero or as near zero as possible • Negative formal charges are more likely to occur on the more electronegative elements • Lewis dot formulas in which adjacent atoms have formula charges of the same sign are usually not accurate Cl=N-O Cl=N-O

32. H H + + Cl Cl Al H H N B H H Cl H H ∙∙ ∙∙ ∙∙ + O N H H H S H H H ∙∙ H H

33. Covalent Bonding • Covalent bonds are formed when atoms share electrons. • If the atoms share 2 electrons a single covalent bond is formed. • If the atoms share 4 electrons a double covalent bond is formed. • If the atoms share 6 electrons a triple covalent bond is formed. • The attraction between the electrons is electrostatic in nature • The atoms have a lower potential energy when bound. • Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds. • Nonpolar covalent bonds have a symmetrical charge distribution. • To be nonpolar the two atoms involved in the bond must be the same element to share equally.

34. Polar and Nonpolar Covalent Bonds • Covalent bonds in which the electrons are not shared equally are designated as polar covalent bonds • Polar covalent bonds have an asymmetrical charge distribution • To be a polar covalent bond the two atoms involved in the bond must have different electronegativities. • Some examples of polar covalent bonds. • HF 2.1 4.0

35. Polar and Nonpolar Covalent Bonds • Compare HF to HI. 2.1 2.5

36. Polar and Nonpolar Covalent Bonds • Polar molecules can be attracted by magnetic and electric fields.

37. Dipole Moments • Molecules whose centers of positive and negative charge do not coincide, have an asymmetric charge distribution, and are polar. • These molecules have a dipole moment. • The dipole moment has the symbol . •  is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q. • There are some nonpolar molecules that have polar bonds. • There are two conditions that must be true for a molecule to be polar. • There must be at least one polar bond present or one lone pair of electrons. • The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

38. Ionic Polar Covalent Covalent Determine Inductive effect Count the number of electrons the element should have Determine how equally electrons are shared (DEN) >1.7 consider it ionic Oxidation number Formal charge