Enriched Chemistry Chapter 4 – Arrangement of Electrons in Atoms

1. Enriched ChemistryChapter 4 – Arrangement of Electrons in Atoms Section One – The Development of a New Atomic Model

2. The Wave Description of Light • Visible light is a kind of electromagnetic radiation; • Together, all forms electromagnetic radiation form the electromagnetic spectrum; • All forms e.r. move at a constant speed of 3.00 x 108 m/s through a vacuum and at slightly slower speeds through matter;

3. The Electromagnetic Spectrum HIGH ENERGY LOW ENERGY

4. Describing Waves • Wavelength () - length of one complete wave • Frequency () - # of waves that pass a point during a certain time period • hertz (Hz) = 1/s • Amplitude (A) - distance from the origin to the trough or crest

5.  crest A origin trough

6. Frequency & wavelength are inversely proportional c =  c: speed of light (3.00  108 m/s) : wavelength (m, nm, etc.) : frequency (Hz)

7. Calculate the wavelength (in meters) of radiation a frequency of 5.00 x 1014 s¯1.

8. When certain frequencies of light strike a metal, electrons are emitted. • Photoelectric effect – refers to the emission of electrons from a metal when light shines on the metal. • Scientists observed that for a specific metal, no electrons were emitted if the light’s frequency was below a certain minimum, regardless of the intensity. • This puzzled scientists because it was not predicted by the wave theory of light. https://www.youtube.com/watch?v=v-1zjdUTu0o

9. The Particle Description of Light • In 1900, German physicist Max Planck suggested that hot objects emit energy in small, specific packets called quanta. • A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom.

10. The energy of a photon is proportional to its frequency. E = h E: energy (J, joules) h: Planck’s constant (6.6262  10-34 J·s) : frequency (Hz) In 1905, Einstein introduced the idea of photons, which are particles of e.r. having zero mass and carrying a quantum of energy.

11. Electrons exist only in very specific energy states for every atom of each element. • Ground state – the lowest energy state of an atom. • Excited state – the atom has a higher potential energy than it has in its ground state. • Basically, when an atom absorbs energy it moves to an excited state; • When that atom returns to its ground state, it releases energy in the form of e.r. • Neon signs are an example.

12. Hydrogen’s Line-Emission Spectrum • 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’s line-emission spectrum. • Scientists had expected to observe the emission of a continuous range of frequencies of electromagnetic radiation, a continuous spectrum.

13. Hydrogen’s Line-Emission Spectrum

14. Bohr Model of the Hydrogen Atom • Niels Bohr proposed a hydrogen-atom model that linked the atom’s electron to photon emission. • According to the model, the electron can circle the nucleus only in allowed paths, or orbits. • The energy of the electron is higher when the electron is in orbits that are successively farther from the nucleus. • When an electron falls to a lower energy level, a photon is emitted, and the process is called emission. • Energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level. This process is called absorption. • video

15. Photon Emission and Absorption

16. Section Two – The Quantum Model of the Atom • In the same way that no two houses have the same address, no two electrons in an atom have the same set of four quantum numbers. • In this section, you will learn how to use the quantum-number code to describe the properties and locations of electrons in atoms.

17. Electrons have wave-like properties. • French scientist Louis de Broglie suggested in 1924 that electrons be considered waves confined to the space around an atomic nucleus. • It followed that the electron waves could exist only at specific frequencies. • According to the relationship E = hν, these frequencies corresponded to specific energies—the quantized energies of Bohr’s orbits.

18. Heisenberg’s Uncertainty Principle • German physicist Werner Heisenberg proposed that any attempt to locate a specific electron with a photon knocks the electron off its course. • Electrons are detected by their interactions with photons. • The Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. http://www.teachertube.com/video/electrons-and-observer-heisenberg-copenhagen-131499

19. Atomic Orbitals and Quantum Numbers • https://www.youtube.com/watch?v=9E3QaRxqXZc

20. Atomic Orbitals and Quantum Numbers • Quantum numbers specify the properties of atomic orbitals and the properties of electrons in orbitals. • The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron.

21. Quantum numbers (cont…) • The angular momentum quantum number, symbolized by l, indicates the shape of the orbital. • We will learn about the s, p, d and f orbital shapes. • N = 1 has one sublevel (s) • N = 2 has two sublevels (s, p) • N = 3 has three sublevels (s, p, d) • N = 4 has four sublevels (s, p, d and f)

22. Atomic Orbitals and Quantum Numbers, continued • The magnetic quantum number, symbolized by m, indicates the orientation of an orbital around the nucleus. • The spin quantum number has only two possible values—(+1/2 , −1/2)—which indicate the two fundamental spin states of an electron in an orbital.

23. Let’s review two terms. • Orbital – a single allowed location for electrons capable of holding two electrons of opposite spin states. • Sublevel – includes all of the similarly shaped orbitals in an energy level.

24. Shapes of s, p, and d Orbitals

25. Electrons Accommodated in Energy Levels and Sublevels

26. Electron Configurations • The arrangement of electrons in an atom is known as the atom’s electron configuration. • The lowest-energy arrangement of the electrons for each element is called the element’s ground-state electron configuration.

27. Rules Governing Electron Configurations • According to the Aufbau principle, an electron occupies the lowest-energy orbital that can receive it. • According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers.