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Exploring Quantum Mechanics: Electrons in Atoms

Learn about quantum mechanics, wave-particle duality, and the uncertainty principle as they relate to electrons in atoms. Discover Bohr's model and its limitations, and delve into the quantum mechanical model and its orbital concept.

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Exploring Quantum Mechanics: Electrons in Atoms

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  1. Section 2: Quantum Theory and the Atom Chapter 5: Electrons in Atoms and the Periodic Table

  2. the branch of mechanics that deals with: a. the mathematical description of the motion and interaction of subatomic particlesB. incorporating the concepts of quantization of energyC. wave-particle dualityD. the uncertainty principleE. the correspondence principle.In laymans terms- The minimum amount of energy that can be gained or lost by an atom. Quantum Mechanics

  3. Bohr’s Model of the Atom • In 1913, Niels Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question. • This model correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.

  4. Bohr’s Model of the Atom • The lowest allowable energy state of an atom is called its ground state. • When an atom gains energy, it is in an excited state.

  5. Bohr’s Model of the Atom • Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.

  6. Bohr’s Model of the Atom • Each orbit was given a number, called the quantum number. • Bohr orbits are like steps of a ladder, each at a specific distance from the nucleus and each at a specific energy.

  7. Bohr’s Model of the Atom • Hydrogen’s single electron is in the n= 1 orbit when it is in the ground state. • When energy is added, the electron moves to the n = 2 orbit.

  8. Bohr’s Model of the Atom • The electron releases energy as it fallsback towards the ground state.

  9. Bohr’s Model of the Atom • Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines. • Spectral Lines- The light or color emitted from the atom when excited. • For this and other reasons, the Bohr model was replaced with a more sophisticated model called the quantum-mechanical or wave-mechanical model.

  10. Quantum Mechanical Model • Louis de Broglie ( Bro- le- ere) (1892–1987) hypothesized that particles, including electrons, could also have wavelike behaviors. • Electrons do not behave like particles flying through space. • We cannot, describe their exact paths.

  11. Quantum Mechanical Model • Heisenberg showed it is impossible to take any measurement of an object without disturbing it. • The Heisenberg uncertainty principlestates that it is fundamentally impossible to know preciselyboth the velocity and position of a particle at the same time.

  12. Quantum Mechanical Model • The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

  13. Quantum Mechanical Model • Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom. • Schrödinger’s equation applied equally well to elements other than hydrogen (unlike Bohr’s model).

  14. Quantum Mechanical Model • The quantum mechanical model makes no attempt to predict the path of an electron around the nucleus. • Bohr orbits were replaced with quantum-mechanical orbitals.

  15. Quantum Mechanical Model • Orbitals are different from orbits in that they represent probability maps that show a statistical distribution of where the electron is likely to be found.

  16. Quantum Mechanical Model • In the quantum-mechanical model, a number and a letter specify an orbital. • The lowest-energy orbital is called the 1s orbital. • It is specified by the number 1 and the letter s.

  17. Hydrogen’s Atomic Orbitals • The number is called the Principal quantum number (n) and it indicatesthe relative size and energy of atomic orbitals. • n specifies the atom’s major energy levels, called the principal energy levels.

  18. Hydrogen’s Atomic Orbitals • Energy sublevels are contained within the principal energy levels.

  19. Hydrogen’s Atomic Orbitals • Each energy sublevel relates to orbitals of different shape. s, p, d, f s, p, d s, p s

  20. Hydrogen’s Atomic Orbitals • s sublevel:

  21. Hydrogen’s Atomic Orbitals • p sublevel:

  22. Hydrogen’s Atomic Orbitals • d sublevel:

  23. Hydrogen’s Atomic Orbitals • f sublevel:

  24. Hydrogen’s Atomic Orbitals • Orbitals are sometimes represented by dots, where the dot density is proportional to the probability of finding the electron. • The dot density for the 1s orbital is greatest near the nucleus and decreases farther away from the nucleus. • The electron is more likely to be found close to the nucleus than far away from it.

  25. Hydrogen’s Atomic Orbitals

  26. Hydrogen’s Atomic Orbitals • At any given time, hydrogen’s electron can occupy just one orbital. • When hydrogen is in the ground state, the electron occupies the 1s orbital. • When the atom gains a quantum of energy, the electron is excited to one of the unoccupied orbitals. • In other words, it will share it’s electron.

  27. https://youtu.be/9E3QaRxqXZc

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