Electrons in Atoms Chapter 5

Objectives Be able to discuss how the Bohr model of the atom was able to explain the existence of atomic emission spectra. Be able to explain how emission spectra are formed. Be able to explain why the emission spectrum of each element is unique.

Models of the Atom John Dalton (1803): solid sphere - - - - J.J. Thomson (1897): plum-pudding - - - - Ernest Rutherford (1911): dense nucleus + Niels Bohr (1913): energy levels + Erwin Schröedinger (1926): electron clouds

Emission Spectra Each element had a unique “emission spectrum,” but what causes it?

Observing of Emission Spectra What we see with our eyes… = A prism or diffraction grating inside a spectroscope separates the individual colors.

The Bohr Model Niels Bohr (1913): Electrons orbit the nucleus at specific energy levels; colors occur when electrons change levels. ultraviolet, UV infrared, IR higher energy visible light lower energy

Production of Emission Spectra photon (1) A quantum (small, exact amount) of energy is absorbed by the atom. (2) An electron in its ground state… (3) … absorbs the energy and makes a quantum leap to an excited state. (4) The excited electron soon drops back down… (5) … releasing energy as a photon (particle of light).

Production of Emission Spectra Each color depends on the energy difference (distance) between the levels: small = red, orange, medium = green, yellow large = blue, violet. very small = infrared (IR) very large = ultraviolet (UV) Each atom has several energy levels, so it can produce several colors in its emission spectrum

Production of Emission Spectra Each element has a unique spacing of energy levels, so it will produce a unique emission spectrum.

Hydrogen’s Spectrum Bohr could only calculate the colors for H, not other “complex” elements We’ll do the calculations in Physics next year.

Objectives Understand the Heisenberg Uncertainty Principle and how it applies to atomic structure. Be able to discuss the concepts of the atomic sublevels and orbitals.

Energy Levels and Sublevels Bohr’s model was too simple, more levels must exist orbits can’t be circular

Heisenberg Uncertainty Principle Heisenberg uncertainty principle: it is not possible to precisely determine an electron’s orbital path. observing the path alters the path The photon changes the path of the electron. Werner Heisenberg

Electrons and Probability 90% Erwin Schröedinger: determined the shapes of orbitals (where an e- is 90% of the time).

Quantum Mechanical Model Electrons exist in orbitals of various shapes (s, p, d, f):

Energy Levels, Sublevels, and Orbitals 4f 4d Energy Levels, Sublevels, and Orbitals n = 4 4p 4s 3d n = 3 3p sublevels (rows) 3s 2p n = 2 2s n = 1 1s orbitals (represented by boxes)

Objectives Understand how electrons fill orbitals in atoms. Be able to write a complete electron configuration for an element.

Orbital Filling Rules Aufbau principle: fill upwards from bottom Pauli exclusion principle: up to 2 e- per orbital, must have opposite spin (↑ or ↓) due to magnetism Hund’s rule: place one e- into each orbital before putting in the second e- What are the electron configurations of H and He? 2p 2p 2s 2s H = 1s1 He = 1s2 1s 1s ↑ ↑↓

Orbital Filling Rules What are the electron 4p What are the electron configurations of O, P, Ti, and Se? 3d 4s 3p 3s 2p 2s 1s

Orbital Filling Rules What is the electron configuration of silicon and what would the atom look like? 3d 3p ↑ ↑ 3s ↑↓ 2p ↑↓ ↑↓ ↑↓ 2s ↑↓ Si = 1s2 2s2 2p6 3s2 3p2 This is why we don’t draw them! 1s ↑↓

Objectives Be able to write a shorthand “noble gas notation” electron configuration. Be able to use the periodic table to write an electron configuration.

Noble Gas Notation A shorthand noble gas notation can be used to represent an electron configuration. 3d ↑ ↑ 3p ↑ 3s ↑↓ 2p ↑↓ ↑↓ ↑↓ same as neon (Ne) 2s ↑↓ P = 1s2 2s2 2p6 3s2 3p3 1s ↑↓ P = [Ne] 3s2 3p3

Overlapping Energy Levels 6d □ □ □ □ □ 5f □ □ □ □ □ □ □ 7s □ 6p □ □ □ 5d □ □ □ □ □ 4f □ □ □ □ □ □ □ 6s □ 5p □ □ □ 4d □ □ □ □ □ 5s □ 4p □ □ □ 3d □ □ □ □ □ 4s □ 3p □ □ □ 3s □ 2p □ □ □ 2s □ 1s □ energy levels and sublevels overlap the periodic table shows this!

Sublevels and the Periodic Table 4f 5f

Writing Electron Configurations 3d 4p 5s 4d What is the electron configuration for Se? Se = [Ar] 4s2 3d10 4p4 What is the electron configuration for Tc? Se = [Kr] 5s2 4d5

Objectives Be able to explain why some elements have exceptional electron configurations. Be able to write the most likely electron configuration of an exceptional element like silver or gold.

Exceptional Electron Configurations Cr and Ag are less corrosive due to this stability! Half-filled and/or completely-filled sublevels are more stable expected: Cr = [Ar] 4s2 3d4 actual: Cr = [Ar] 4s1 3d5 expected: Ag = [Kr] 5s2 4d9 actual: Ag = [Kr] 5s1 4d10 Usually, a single electron gets placed in a d-sublevel before the following electrons get placed in the f-sublevel. actual: U = [Rn] 7s2 6d1 5f 3