10/22/2012 ACT QotD

10/22/2012ACT QotD • As we have seen, the Hydrogen Spectrum contains three visible spectral lines; red, blue/green, and violet. Arrange these spectral lines according to decreasing energy of their photons. • a) red < violet < blue • b) violet < blue < red • c) blue < red < violet • d) red < blue < violet

Drawback • Bohr’s theory did not explain or show the shape or the path traveled by the electrons. • His theory could only explain hydrogen and not the more complex atoms

The Quantum Mechanical Model • Energy is quantized. It comes in chunks. • A quanta is the amount of energy needed to move from one energy level to another. • Since the energy of an atom is never “in between” there must be a quantum leap in energy. • Schrödinger derived an equation that described the energy and position of the electrons in an atom

Atomic Orbitals • Principal Quantum Number (n) = the energy level of the electron. • Within each energy level the complex math of Schrödinger's equation describes several shapes. • These are called atomic orbitals • Regions where there is a high probability of finding an electron

S orbitals • 1 s orbital for every energy level 1s 2s 3s • Spherical shaped • Each s orbital can hold 2 electrons • Called the 1s, 2s, 3s, etc.. orbitals

P orbitals • Start at the second energy level • 3 different directions • 3 different shapes • Each orbital can hold 2 electrons

The p Sublevel has 3 p orbitals

The D sublevel contains 5 D orbitals • The D sublevel starts in the 3rd energy level • 5 different shapes (orbitals) • Each orbital can hold 2 electrons

The F sublevel has 7 F orbitals • The F sublevel starts in the fourth energy level • The F sublevel has seven different shapes (orbitals) • 2 electrons per orbital

10/23/2012ACT QotD • A(An) __________ is a region of space in which there is a high probability of finding an electron in an atom. • a. shell • b. atomic orbital • c. core • d. major energy level • e. nucleus

Summary Starts at energy level

Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

Electron Configurations First Energy Level • only s sublevel (1 s orbital) • only 2 electrons • 1s2 Second Energy Level • s and p sublevels (s and p orbitals are available) • 2 in s, 6 in p • 2s22p6 • 8 total electrons

Third energy level • s, p, and d orbitals • 2 in s, 6 in p, and 10 in d • 3s23p63d10 • 18 total electrons Fourth energy level • s,p,d, and f orbitals • 2 in s, 6 in p, 10 in d, and 14 in f • 4s24p64d104f14 • 32 total electrons

7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

Electron Configuration • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.

Orbitals fill in order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. • Half filled orbitals have a lower energy. • Makes them more stable. • Changes the filling order

Write these electron configurations • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 is expected • But this is wrong!!

Chromium is actually • 1s22s22p63s23p64s13d5 • Why? • This gives us two half filled orbitals. • Slightly lower in energy. • The same principal applies to copper.

Copper’s electron configuration • Copper has 29 electrons so we expect • 1s22s22p63s23p64s23d9 • But the actual configuration is • 1s22s22p63s23p64s13d10 • This gives one filled orbital and one half filled orbital. • Remember these exceptions

Great site to practice and instantly see results for electron configuration.

Practice • Time to practice on your own filling up electron configurations. • Do electron configurations for the first 20 elements on the periodic table.

10/22/2012 ACT QotD