The Arrangement of Electrons in Atoms

1. The Arrangement of Electrons in Atoms Ch. 4 Modern Chemistry

2. A New Atomic Model • Cathode Ray Tube Summary • As of 1911, Rutherford’s Gold Foil Experiment had shown scientists the nucleus of an atom was very small, dense, and positively charged. • Why was Rutherford’s model of the atom incomplete? • Rutherford’s model did not explain how electrons were distributed around the nucleus. • Just assumed electrons located around nucleus

3. Wave-like Properties of Light • Over the next decade, scientists began to discover the dual nature of electrons - particles and waves. • Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. • Together, all the forms of electromagnetic radiation form the electromagnetic spectrum. • Visible spectrum: visible portion of spectrum • But there are many other forms of energy

4. Electromagnetic Spectrum

5. Properties of Light Wavelength (λ): distance between corresponding points on adjacent waves Frequency (ν): number of waves that pass a given point per unit time time (usually 1 sec)

6. Properties of Light • Frequency and wavelength are mathematically related to each other: c = λν • c: speed of light (m/s) • λ: wavelength (m) • ν: frequency of wave (s−1 or 1/s).

7. Photoelectric EffectParticle Description of Light Electromagnetic radiation strikes the surface of the metal Electrons are ejected from the metal causing an electric current

8. Photoelectric Effect • The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. • Not all forms of electromagnetic radiation were “powerful” enough to eject electrons from the surface • A minimum amount of energy was needed • A quantum of energy - the minimum quantity of energy that can be lost or gained by an atom

9. Photoelectric Effect • German physicist Max Planck proposed the following relationship: E = hν • E: quantum of energy (J – joules) • ν : frequency of the radiation emitted; s−1 • h :fundamental constant called Planck’s constant; h = 6.626 × 10−34 J• s

10. Particle Description of Light photon : particle of electromagnetic radiation having zero mass and carrying a quantum of energy. The energy of a particular photon depends on the frequency of the radiation. Ephoton = hν

11. The Hydrogen Line Emission Spectrum

12. The Hydrogen Line Emission Spectrum When investigators passed electric current through a vacuum tube containing hydrogen gas at low pressure, they observed the emission of a characteristic pinkish glow. When a narrow beam of the emitted light was shined through a prism, it was separated into four specific colors of the visible spectrum. The four bands of light were part of what is known as hydrogen’sline-emission spectrum.

13. The Hydrogen Line Emission Spectrum • What had scientists expected to observe? • A continuous range of frequency (all bands of ER)

14. The Hydrogen Line Emission Spectrum The lowest energy state of an atom is its ground state. A state in which an atom has a higher potential energy than it has in its ground state is an excited state.

15. Bohr Model of the Atom • By 1922, NielsBohr proposed a hydrogen-atom model that linked the atom’s electron to photon emission. • electrons can circle the nucleus only in allowed paths, called orbits. • The energy of the electron is higher when it is in an orbit farther from the nucleus.

16. Bohr Model of the Atom • When electrons move from an excited state to a ground state or to a lower energy state, a photon of energy is emitted from an atom • Process called emission (high to low orbit) • absorption (low to high orbit): energy is added to an atom to move an electron from lower energy level to higher energy level 3) The energy of a photon can be directly calculated if you know the frequency of radiation being emitted

17. Photo Absorption & Emission

18. Section 4.2 Quantum Model of the Atom

19. Electrons as Waves • French scientist Louis de Brogliesuggested that electrons be considered waves confined to the space around an atomic nucleus. • electron waves could exist only at specific frequencies. • It followed that if each electron has a specific frequency, it also has a corresponding energy. • Recall equation E = hv

20. Heisenberg Uncertainty Principle German physicist Werner Heisenberg proposed that any attempt to locate a specific electron with a photon knocks the electron off its course. The Heisenberg uncertainty principle: it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.

21. The Schrödinger Wave Equation In 1926, Austrian physicist Erwin Schrödinger developed an equation that treated electrons as waves in an atom. Heisenberg & Schrödinger ideas laid the foundation for modern quantum theory. Quantum theory describes mathematically the wave properties of electrons and other very small particles.

22. The Schrödinger Wave Equation • Electrons do not travel around the nucleus in neat orbits, as Bohr had postulated, but rather in certain regions called orbitals. • An orbital: three-dimensional region around the nucleus - indicates the probable location of an electron. • Quantum numbers specify the properties of atomic orbitals and the properties (location & spin) of electrons in orbitals. • 4 quantum numbers

23. How should we look at the atom? Models of the Atom – An Atomic Rant