CHEM 1465 1st Edition Lecture 11Outline of Last Lecture 1. Terms to know2. Formulas to know3. Classic mechanicsA. The photoelectric effectB. Atomic spectra 4. Wave-like properties of matterOutline of Current Lecture 1. Quantum mechanics and the Heisenberg uncertainty principle2. Quantum numbers3. Electron configurationA. RulesB. Terms to knowC. Using periodic table4. Periodic trends 5. Atomic and ionic radii6. Ionization energy7. Electron affinity Current Lecture1. Quantum mechanics and the Heisenberg uncertainty principle- Niels Bohr suggested that electrons exists in orbitals.- Quantum mechanics treats electrons differently. Quantum mechanics says that an electron has both particle like and wave like properties - The Heisenberg uncertainty principle say it is impossible to know both the position and the momentum of a particle simultaneously - Thus we cannot assign fixed paths for electrons, and the most we can hope to know is the probability of finding an electron in a given region of space.These 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.- Since an electron has wave like properties, quantum mechanics is also called wave mechanics. - Quantum mechanics is a mathematical description of an electron- Schrodinger’s equation: Hψ=Eψ where H is Hamiltonian operator, ψ is wave function, and E is energy- The physical meaning of ψ is probability density- the probability of finding an electron in a given volume in space2. Quantum numbersletter designation Name What is describes Possible valuesn Principal QN Shell of electron 1,2,3,4…l Angular momentumQNShape of orbitalL=0 is s orbitalL=1 is p orbitalL=2 is d orbitalL=3 is f orbital 0,1,2,3MlMagnetic QN Orientation -L to +LMsSpin QN Spin of electron -1/2 or +1/23. Electron configuration A. Rules for determining the ground state electron configuration- Aufbau procedure: start with lowest energy orbital and fill up from there- Pauli Exclusion Principle: no two electrons in an atom can have the same 4 quantum numbers. You can have a maximum of 2 electrons per orbital as long as they have opposite spins- Hund’s rule: when filling degenerate orbital (orbitals that are equal in energy)electrons always occupy singly with parallel spin before paired. (remember: parallel before paired)B. Terms to know- ground state: lowest energy state of an electron or molecule- excited state: any state that is not the ground state- degenerate: equal in energy- node: region of space with zero probability of finding an electron- shell: n= 1 shell n =2 shell- subshell: nlC. Using periodic table- Remember the s block starts with 1s and the p block starts with 2p. The d block starts 3d and the f block starts with 4f.- There is a special stability associated with filled outer most shell. This explains stability of the noble gases- There is a special stability associated with half hilled subshells. - When determining the electron configuration of cation, the first electrons removed are from the outer most shell. This is not always the outer most subshell. - Isoelectronic: having the same electron configuration- Paramagnetic: attracted to magnetic field (has an unpaired electron)- Diamagnetic: slightly repelled by magnetic field (all electrons paired)- Valence electrons: outer shell electrons- Core electrons: all other electrons, called inner core, inner electrons4. Periodic trendsA. Terms to know- Shielding: describes decrease in attraction of the nucleus for outer shell electrons due to inner core- Effective nuclear charge: nuclear change that an electron feelsB. To predict periodic trends, consider two things:- electron configuration- nuclear charge5. Atomic and ionic radii- Atomic size decreases going from left to right across the periodic table because we are increasing the nuclear charge.- Atomic size increase going down the periodic table because outer most electronsare in a higher shell which is farther from the nucleus- Ionic size: cations are always smaller than parent ion. Anions are always bigger than parent ion. In an isoelectronic series, ionic size decreases with increasing nuclear charge. 6. Ionization energy- Ionization energy: energy required to remove an electron from gaseous atom- 1E1 is the first ionization energy- 1E2 is 2nd electron removal- 1E3 is 3rdelectron removal- Notice the huge jump from 1E1 to 1E2. This is because it is harder to remove inner core electrons.7. Electron affinity: the energy that occurs when electrons are added to gaseous atoms- Some electrons are positive and some are negative. If the affinity is negative thenthe atoms wants an electron. When the affinity increases that means the electron affinity is increasingly negative- Group 2 and 18 have positive or near positive electron affinities. The halogens have the most exothermic or negative electron affinities. - You can make individual predictions based on electron configuration and nuclear