Chem 121 1st Edition Lecture 8Outline of Last Lecture I. Quantum NumbersII. Orbital ShapesOutline of Current Lecture II. Periodic Law and Periodic TableIII. Properties of a GroupIV. Classification of ElementsV. Electron- Spin and the Pauli PrincipleVI. Multi- Electron AtomsVII. Sublevel Energy SplittingVIII. Aufbau Principle and Electron ConfigurationCurrent LecturePeriodic Law and Periodic Table- Dimitri Mendeleev- Russian Chemist (1834-1907)o Put the elements as a deck of cards and arranged them in increasing atomic mass & their oxideso Properties of the elements repeated every 8 elements of first 20 elementso Predicted eka-silicon Germanium 1886o Periodic Law: “When the elements are arranged in order of increasing mass (atomic #), certain sets of properties reoccur periodically”Properties of a Group -Period: horizontal rows-Group (family): vertical columnsoMain Groups (Representative groups)1A(1), 2A(2), 3A(13)~8A(18)-Alkali metals (Except H)-1A; Alkaline earth metals-2A; Noble gases-8A; Halogens-7A; Chalcogens-6ATransition elements-3B(3)~2B(12); 8B: three columns; 1B (11)-Coinage metalsInner transition elements: between 3B and 4BThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.-Lanthanides (rare earth, Period 6) and Actinides (Period 7)Classification of Elements- Metals (Left of Periodic Table)o Good conductors of heat and electricityo Majority usually have ‘ium’ at the end- Nonmetals (Upper right of Periodic Table)o Poor conductors of heat and electricity- Metalloids (above and under zigzag)o Properties between metals and nonmetalso Zigzag is one of the most important lines to knowElectron- Spin and the Pauli Principle- Spin of electrons: spin quantum number (ms- Description of electron, direction of e- spin), can only be -1/2 or +1/2- Pauli Exclusion Principle: “No two electrons in the same atom can have the same set of quantum numbers”o Unique “identity” for each electron in an atomo Conclusion: “An atomic orbital can hold a maximum of two electrons with opposite spin” Capacity of an atomic orbital: 2 Total of 4 quantum numbers to identify e- Start electrons with lowest energy of atomic orbital possible, lowest quantum number possibleMulti- Electron Atoms- Attraction from nucleusVN ∞ - Ze2/ R Z: # of protonsR: distance between nucleus and electron- The closer electron to nucleus = lower energy level, the farther electron from nucleus= higher energy level- R determined by no Greater the n, greater the Ro Same distance R, same energy level- No energy splitting within a shellSublevel Energy Splitting- Penetration (PE): electron can be found in inner shell with small probability- PE effect lower energy of electrono “The more penetration, the closer the electron to nucleus, the more attraction the electron “feels”, the more stable the energy level”- Energy level when occupied: s<p<d<fo Smaller the l, lower the energyAufbau Principle and Electron Configuration- Aufbau Principle: “o “Electrons occupy orbitals starting with the lowest possible energy subshell available”o “Electrons must occupy all orbitals in the subshell before moving to next available one”o “Available for ground-state atom only”- Electron configurationo Arrangement of electrons in orbitals of an atomo Subshell with more than one orbitalso Hund’s Rule: “When filling a subshell having more than 1 orbitals (p,d,f,…), electron: Fills each orbital first With parallel