CHAPTER 5: ELECTRONS IN ATOMS Rutherford’s model of the atom did not explain how the electrons occupy the space surrounding the nucleus. Recall: Opposites attract….right?? In the early 1900’s, investigations using the relationships that exist between light and electrons will reveal a more accurate model of the atom. Model of The Atom

Emission Spectrum atoms can exist in low energy states called ground states or at high energy states called excited states. when excited atoms return to ground states, it gives off energy in the form of electromagnetic radiation. all elements have a characteristic emission spectra that corresponds to the amount of energy given off. Ground & Excited States Emission Spectra

Probability of Finding an Electron in a Hydrogen Atom.

The Bohr Model ---it’s really not boring Niels Bohr linked photon emission with an atom’s electrons. electrons exist in fixed orbits and never in between. his model only explains the emission of Hydrogen atoms. Niels Bohr Flame Test Emission Flame Test Explanation The Quantum Model of the Atom Louis deBroglie suggested that electrons be considered waves confined to the space around the nucleus.

electrons, like waves, are capable of interference. investigations also confirmed that electrons can be bent or diffracted. electrons, like waves, are capable of interference. Wave Properties Of Electrons Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity of an electron. Heisenberg Erwin Schrodinger’s wave equation and Heisenberg’s Uncertainty Principle laid the foundation for the quantum model of the atom. electrons do not travel in specified orbits. instead electrons exist in regions called orbitals. Quantum Model of the Atom

Electron Configurations “The arrangement of electrons in an atom” Rules Governing Electron Configurations Aufbau Principle: an electron occupies the lowest-energy orbital that can receive it. 2. Pauli Exclusion Principle: no two electrons in the same atom can have the same set of four quantum numbers. Quantum Subshells

3. Hund’s Rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. all electrons in singly occupied orbitals must have the same spin. Orbital Notation a way to illustrate what orbitals the electrons occupy in an atom. ie: Nitrogen – 7 electrons

Electron Configurations electrons configurations (shorthand method) are a method of noting how many electrons are in an atom and where they are located. Order of Filling Maximum amount of electrons: s-orbital  2 electrons p-orbital  6 electrons d-orbital  10 electrons f-orbital  14 electrons Quantum Subshells

Example: Determine the electron configuration for the following elements: Lithium  3 electrons 1s22s1 Nitrogen  7 electrons 1s22s22p3 Neon  10 electrons 1s22s22p6

Noble Gas Notation in order to write long configurations much easier we can use a noble gas configuration. In order to write a noble gas notation simply include the previous noble gas (in brackets) and list the electron configuration from the next element from the noble gas to the element you are determining the configuration for. Write the noble gas notation for Sodium [Ne]3s¹ Write the noble gas notation for Sulfur [Ne]3s²3p⁴

Light behaves both as a wave and a particle. Properties of Light Light behaves both as a wave and a particle. Light’s Wavelike Properties visible light is one kind of electromagnetic radiation. Light as a Wave Electromagnetic Spectrum Prism

all electromagnetic radiation moves at the speed of light (3 all electromagnetic radiation moves at the speed of light (3.00 x 108 m/s) Properties of Waves Wavelength: (l) the distance between corresponding points on adjacent waves. 2. Frequency: (n) the number of waves that pass a given point in one second. Frequency is measured in Hertz (Hz).

Speed of Light Equation What is the wavelength of radiation with a frequency of 1.5 x 1013 Hz? Wavelength x 1.50 x 1013Hz = 3.00 x 108m/s Wavelength = 3.00 x 108m/s / 1.50 x 1013Hz Wavelength = 2.00 x 10-5m

The Photoelectric Effect a phenomena involved with the emission of electrons from a metal when light shines on a metal. light of only a certain frequency enabled this event to occur  light behaves like streams of particles! Photoelectric Effect Max Planck suggested that energy could only exist in discrete lumps called quantum. E = h n h = 6.626 x 10-34 Js Planck and Blackbody Radiation Einstein also proposed that light delivers its energy in chunks called photons. Therefore, in order for an electron to be ejected (photoelectric effect) , the electron must be struck by a single photon, of certain energy, to knock the electron loose.

One more question……….Are You Ready? The threshold photoelectric effect in tungsten is produced by light of a wavelength 260 nm. Give the energy of a photon of this light in joules. Find frequency! λ x ν = C 260 nm x 1m/109 nm = 2.6 x 10-7 m 2.6 x 10-7 m x ν = 3.0 x 108 m/s v = 1.15 x 1015/sec E = h x ν E = 6.626 x 10-34Js x 1.15 x 1015/sec E = 7.6 x 10-19 J