First three quantum numbers uniquely describe orbitals 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f level: contains orbitals with the same n

Transcription

1 1.1 Orbital Energies 1.2 Orbital Energies First three quantum numbers uniquely describe orbitals 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f level: contains orbitals with the same n sublevel: contains orbitals with the same n and l What determines the relative energies of these orbitals? Which are lower in energy? Which are higher in energy? What general principle explains orbital energies? Which orbital has higher energy, 1s or 2s? Why? Which orbital has higher energy, 2s or 2p? Why? Which orbital has higher energy, 2p x, 2p y or 2p z? Why? 1.3 Orbital Energies 1.4 Orbital Energies Which orbitals have higher energy, 3s, 3p or 3d? Why? (what do you need to know?) Will the five 3d orbitals have the same energy? Why or why not? 1

2 1.5 Energy of Orbitals: Summary For the same type of orbital (same ), energy increases as n increases (1s < 2s < 3s < 4s ) For the same n, energy increases s < p < d < f (3s < 3p < 3d) All orbitals of the same sublevel have the same energy (they are ) (3p x = 3p y = 3p z ) 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p 2.1 Magnetic Properties: Electron A physical phenomenon: spinning, charged particles produce magnetic fields Spinning electrons produce tiny magnetic fields Electrons can spin in one of two directions 2.2 Magnetic Properties: Electron 2.3 The 4 th Quantum Number Diamagnetic Electron spin, m s : m s = ½ or -½ Paramagnetic Pauli exclusion principle: 2

3 3.1 Electron Configuration Rules 3.2 Electron Configurations Electrons fill the lowest energy orbital first (Aufbau principle) Two notations for the arrangement of electrons in atoms Two electrons (max) per orbital Maximize parallel spins when filling a sublevel spdf notation Fill orbitals of equal energy with one electron each before pairing up Why don t we double-up first? noble gas notation 3.3 Electron Configurations 3.4 More Examples Hydrogen Provide the electron configurations (in spdf and noble gas notation) phosphorus Lithium vanadium Oxygen iodine 3

4 3.5 More Examples: Ions 4.1 Periodic Table Organization (a) S 2 So does S 2 = Ar? (b) Br (c) Al Periodic Table Organization 4.3 Periodic Properties s-block p-block d-block Valence electrons Core electrons atoms where an s sublevel is being filled atoms where a p sublevel is being filled atoms where a d sublevel is being filled Same group = same number and type of valence electrons You will need to know the following: 1. Definitions and chemical equations where appropriate 2. Periodic trends moving up and down and left to right across the periodic table 3. Explanations of the trends 4. How the atomic properties affect chemical properties 4

5 4.4 Effective Nuclear Charge 4.5 Effective Nuclear Charge Valence electrons don t feel the full charge of the nucleus Valence electrons are shielded from nuclear charge (Z) Take the case of Li 1s 2 2s 1 Z eff : the positive charge felt by a valence electron An crude approximation: Z eff = Z core electrons where Z = atomic number Z eff increases across the periodic table Lithium: Z eff = 3 2 = 1 Carbon: Z eff =6 2=4 Fluorine: Z eff = 9 2 = 7 Sodium: Z eff = Silicon: Z eff = 4.6 Atomic Size 4.7 Atomic Size The distance from the nucleus to the edge of the outermost electron Periodic trend: Explanation: 5

6 5.1 Ionization Energy (IE) 5.2 Sign Conventions The energy required to remove an electron from a gaseous atom Energy input required A(g) + energy A + (g) + e - Energy absorbed (in) = a positive value kj Energy required (input, raw material) Energy released (out) = a negative value kj Energy produced (output, product) The sign tells us which way energy is going The magnitude tells us how much energy is involved 5.3 Ionization Energies 5.4 Ionization Energies: Summary 6

7 5.5 Successive Ionizations 5.6 Successive Ionizations IE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Na Mg Al Si P S Example: Na(g) + IE 1 Na + (g) + e - Na + (g) + IE 2 Na 2+ (g) + e - For Mg, 2 nd IE > 1 st IE For Al, 3 rd IE > 2 nd IE > 1 st IE Why? For Mg, 3 rd IE >>> 2 nd IE For Al, 4 th IE >>> 3 rd IE Why? 5.7 Electron Affinity 5.8 Electron Affinity Predictions The energy released when an electron is added to a gaseous atom A(g) + e - A - (g) + energy A free electron is not a stable beast. It would rather be associated with an atom. A(g) + e - A - (g) + energy Across a period (left to right): Should it get easier or harder to add an electron? Down a group: Should it get easier or harder to add an electron? If it s easy to add an electron, is the EA a large negative number or a small negative number? Deviations from the general trends 7

8 5.9 Electron Affinity Trends 5.10 Electron Affinity Summary 6.1 Ionization: Change in Size 6.2 Ionization: Change in Size Why does the size decrease? Why does the size increase? 8