Unit 3: Electrons in the Atom

Copyright © Houghton Mifflin Company 1 The Rutherford atom model. A positive nucleus surrounded by electrons like our solar system. However, this model did not properly explain chemical reactivity and certain light phenomena.

In 1665 Sir Isaac Newton noticed that white (sun) light could be split into a multicolored band of light just like a rainbow. The multicolored band of light is called a color spectrum. Brief History of Light

Light as a wave In the 19 th century the works of Michael Faraday and (later) James Maxwell showed that electricity and magnetism are simply two parts of a single phenomenon, electromagnetism. This phenomenon would produce waves which travel at the speed of light, having light waves other than those that produced the light that we could see. We now refer to this collection of different waves of electromagnetic radiation (light) as the electromagnetic spectrum (EMS)

Waves There are 2 types of waves, transverse and longitudinal.

Waves are characterized by three properties: 1. wavelength, 2. frequency, 3. amplitude

Wavelength and Frequency An important feature of wave motion is the inverse relationship between wavelength and frequency. That is, as one increases the other decreases. They are related in the following way: c = λ ν Speed of light (c) = wavelength ( λ) x frequency (ν)

c has a constant value of 3.00 x 10 8 m/s. Wavelength is in meters (m). Frequency is in cycles/s, in s -1, in 1/s, or in Hertz (Hz).

Example Calculate the frequency of light with a wavelength of 5.22 x m. c = λ ν 3.00 x 10 8 m/s = 5.22 x m x ν so ν = (3.00 x 10 8 m/s) / (5.22 x m) = 5.75 x /s = 5.75 x Hz

Practice Calculate the wavelength of a radio station signal with a frequency of 99.7 MHz.

Electromagnetic Spectrum In increasing frequency, ROY G BIV

Electromagnetic Spectrum Long wavelength --> small frequency Short wavelength --> high frequency increasing frequency increasing wavelength

Light as particles In 1900 German scientist Max Planck found that light is given off in discrete units (quanta). He also found light energy (E) is proportional to its frequency ( ν ). The relationship is: E = h ν, where h = Planck’s constant, 6.63 x J. S.

Wave particle nature In 1905 Albert Einstein confirmed Planck’s findings and he called the quanta “photons” (packets of energy).

Example Calculate the energy of a photon of light with a frequency of 5.45 x Hz. E = h ν = 6.63 x Js x 5.45 x s -1 = 3.61 x J. You try: Calculate the energy of a photon having a wavelength of 4.5 x m.

Bohr’s observations In 1913 Niels Bohr used the observations of Planck to explain the specific lines observed in the hydrogen emission spectrum. These lines resulted from some whole number transition.

Bohr’s model Bohr suggested that a transition corresponded to an electron jumping from one possible orbit to another and emitting a photon of light energy. In Bohr’s model of the atom, the electron can only exist in these specific orbits, known as energy levels, in an atom. Normally the electron would be in its lowest available energy level, this is called its ground state.

If the atom is exposed to an energy source the electron can absorb a quantum of energy (photon) and the electron will make a quantum leap to a higher energy level. Then the electron will drop back down to a lower energy level, thereby emitting a photon of light. The energy of this photon would correspond exactly to the energy difference between the two levels.

Light emitted produces a unique emission spectrum.

n=1 n=2 n=3 n=4 Spectrum UV IR VisibleVisible Ground State Excited State Excited State unstable and drops back down Energy released as a photon Frequency proportional to energy drop Excited State But only as far as n = 2 this time

Emission Spectrum Animation

Line Spectra of Other Elements

Wave – particle duality Louis De Broglie (1924) proposed that ALL matter has wave and particle properties, not just electrons. Heisenberg (1927) said that because of size and speed it is impossible to know both exact position and momentum of an electron at the same time. –This is referred to as “Heisenberg Uncertainty Principle”

Quantum mechanical model Schroedinger ( ) developed the “quantum mechanical model” of the atom. He calculated the probability where to find electrons, thereby creating “electron clouds”: areas with a great chance (90 %) to find electrons. The region in space in which there is a high probability of finding an electron is now known as an “orbital”.

Orbitals Every element has discrete energy levels called principal energy levels (given with letter n). The principal levels are divided into sublevels. Sublevels contain spaces for the electron called orbitals.

Orbital types s-orbital = spherical shape, only 1 of them p-orbital = gumdrop or dumbell shape, 3 of them – one on each axis (x,y,z) d-orbital = donut shape, 5 of them f-orbital = cigar shape, 7 of them Each orbital contains a max of 2 electrons

s and p orbitals

d orbitals

Filling orbitals with electrons We have 3 general rules for “distributing” these electrons. 1.Pauli Exclusion Principal: Orbitals contain no more than two electrons. –Each electron has a spin: up (↑) or down (↓) –Two electrons must have opposite spins to occupy an orbital

2. Hund Rule: When filling orbitals, assign one electron to each orbital (of that type) before doubling up with two electrons per orbital. 3. Aufbau: Electrons fill lowest orbitals first, then proceed to higher energy levels.

First 4 energy levels

Filling orbitals Energy level Orbital type # orbitals # of types# electrons n = 1s11 s2 n = 2s, p41 s, 3 p8 n = 3s, p, d91 s, 3 p, 5 d18 n = 4s, p, d, f161 s, 3 p, 5 d, 7 f 32 Energy level = the number of orbital types Total number of orbitals in an energy level = n 2 Total number of electrons in any energy level = 2n 2

Electron configuration The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. The sublevel is written followed by a superscript with the number of electrons in the sublevel.

Electron configuration H1s 1 He1s 2 Li1s 2 2s 1 Be1s 2 2s 2 B1s 2 2s 2 2p 1 C1s 2 2s 2 2p 2 N1s 2 2s 2 2p 3 O1s 2 2s 2 2p 4 F1s 2 2s 2 2p 5 Ne1s 2 2s 2 2p 6

Filling Diagram for Sublevels

Order of filling orbitals 1s (with 2 electrons) 2s (2), 2p (6) 3s (2), 3p (6) 4s (2), 3d (10), 4p (6) 5s (2), 4d (10), 5p (6) 6s (2), 4f (14), 5d (10), 6p (6)

Practice: Give the electron configuration for: P 1s 2 2s 2 2p 6 3s 2 3p 3 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Br Al

Electron configuration and the Periodic Table

Abbreviated notation When an energy level is completely filled we often use an abbreviated notation with the noble gas configuration of the last filled period representing the inner electrons. Example: Na 1s 2 2s 2 2p 6 3s 1 or [Ne]3s 1

Practice Give the abbreviated electron configuration of the following elements: S Co I

Electron configuration of Cu Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 However, Cu is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 It is energetically slightly favorable for Cu to completely fill the 3d orbital, so one electron is moved from the 4s to the 3d orbital.

Electron configuration of Cr Cr shows a similar electron configuration effect as Cu. Cr is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 rather than 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Due to the fact that a half-filled 3d orbital is energetically favorable over a filled 4s orbital.

Valence electrons These electrons are in the outermost principal energy level of an atom: the s and p electrons beyond the noble gas core. These electrons are involved in forming bonds with other atoms Inner electrons (core electrons) are NOT involved in bonding

Electron dot structure Elements (except helium) have the same # of valence electrons as their group #. Electron dot structures are used to show valence electrons. We use one dot for each valence electron. Consider phosphorus, P, which has 5 valence electrons. Here is the method for writing the electron dot formula.