![• Shorthand Configuration
S 16e-
Valence ElectronsCore Electrons
S 16e-
[Ne] 3s2
3p4
1s2
2s2
2p6
3s2
3p4
B. Notation
• Longhand Configuration](https://image.slidesharecdn.com/chap5pptchemistry2016-161003042929/85/Chap-5-ppt-chemistry-2016-28-320.jpg)
Chap 5 ppt chemistry 2016
- 1. Chapter 5: Electrons in Atoms
- 2. » Rutherford’s model fails to explain why objects change color when heated. •Rutherford’s atomic model could not explain the chemical properties of elements.
- 3. Models of The Atom 1863- John Dalton pictures atoms as tiny, indestructible particles, with no internal structure.
- 4. Models of the Atom • 1897- J.J. Thomson, a British scientist, discovers the electron. The later leads to his “Plum Pudding” model. He pictures electrons embedded in a sphere of positive electrical charge.
- 5. Models Of the Atom •1911- Ernest Rutherford finds that an atom has a small, dense, positively charged nucleus. • Most of the atom is empty space in which the electrons move.
- 6. The Bohr Model • 1913- Bohr proposed that an electron is found only in specific circular paths or orbits around the nucleus.
- 7. The Bohr Model • Each possible electron orbit has a fixed energy called a fixed energy level. The Bohr Model
- 8. •These ladder steps are somewhat like energy levels. •The higher an electron is on the energy ladder, the farther it is from the nucleus. • quantum of energy is the amount of energy required to move an electron from one energy level to another energy level. Nucleus
- 9. • To move from one energy level to another an electron must gain or lose just the right amount of energy. • The higher an electron is on the energy ladder, the farther from the nucleus.
- 10. Nucleus
- 11. The Bohr Model • The Bohr model establishes the concept of definite electron energy levels within atoms. But Bohr's model was rather simplistic and as scientists made more discoveries about more complex atoms, Bohr's model was modified and eventually was replaced by more sophisticated models.
- 12. Quantum Mechanical Model (QMM) • In 1926, the quantum mechanical model was proposed as a mathematical model to describe the energies and probable locations of electrons in atoms. • nucleus electron cloud
- 13. Quantum Mechanical Model • In 1926, the quantum mechanical model was proposed as a mathematical model to describe the energies of electrons in atoms. • nucleus electron cloud
- 14. Key Concept!!! • The Quantum Mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
- 15. The Quantum Mechanical Model • The modern description of the electrons in atoms, the Quantum Mechanical Model, comes from the mathematical solutions to the Schrodinger equation. • Like the Bohr model, the quantum mechanical model of the atom restricts the energy of electrons to certain values. • Unlike the Bohr model, however, the QMM does not involve an exact path the electron takes around the nucleus.
- 16. • THINGS TO NOTICE • The nucleus is not actually shown, but is located at the center. • You can't tell exactly where the electron is, just where it is most likely to be. • The individual dots are not electrons. They are meant to be used in the context of how dense, or heavy an area of dots appears. • The more crowded (or heavier packed) the dots are in an area the better chance you have to finding your electron there. • Hydrogen Atom Quantum Mechanical Model (QMM)
- 17. Atomic Orbitals • An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. • ORBITALS are not the same thing as the Bohr’s Orbits.
- 18. • Here is a great website that shows the different orbitals!!! The Orbitron: a gallery o
- 19. Atomic Orbitals Cont… • The numbers and kinds of atomic orbitals depend on the energy sublevel. • The lowest principal energy level (n=1) has only one sublevel, called 1s. • The second level (n=2) has two sublevels, 2s and 2p. Thus, the second energy level has four orbitals: 2s, 2px,2py,2pz. • The third level (n=3) has three sublevels, 3s,3p and 3d. Thus the third energy level has nine orbitals: one 3s, three 3p, and five 3d. • The fourth level (n=4) has four sublevels, 4s,4p,4d, and 4f. Thus the fourth energy level has 16 orbitals: one 4s, three 4p, five 4d, and seven 4f orbitals.
- 20. Maximum # of Electrons • Energy level n Max. # of electrons • 1 2 • 2 8 • 3 18 • 4 32
- 21. 5.2 Electron Arrangement in Atoms • The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations. • Three Rules- the Aufbau Principle, the Pauli exclusion principle, and Hund’s rule- tell you how to find the electron configurations of atoms.
- 22. Aufbau Principle • Electrons occupy the orbitals of lowest energy first.
- 23. Pauli Exclusion Principle • An atomic orbital may describe at most two electrons. – For example, either one or two electrons can occupy an s or p orbital. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.
- 24. A. General Rules • Pauli Exclusion Principle – Each orbital can hold TWO electrons with opposite spins.
- 25. Hund’s Rule • Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.
- 26. RIGHTWRONG A. General Rules • Hund’s Rule – Within a sublevel, place one e- per orbital before pairing them. – “Empty Bus Seat Rule”
- 27. O 8e- • Orbital Diagram • Electron Configuration 1s2 2s2 2p4 B. Notation 1s 2s 2p
- 28. • Shorthand Configuration S 16e- Valence ElectronsCore Electrons S 16e- [Ne] 3s2 3p4 1s2 2s2 2p6 3s2 3p4 B. Notation • Longhand Configuration
- 29. s-block1st Period 1s1 1st column of s-block C. Periodic Patterns • Example - Hydrogen
- 30. Practice, Practice, Practice!!! • Write the electron configurations for each atom. • Carbon • Argon • Nickel
- 31. 5.3 Physics and the QMM
- 32. Light • The Quantum mechanical model grew out of the study of light. • Isaac Newton tried to explain what was known about the behavior of light by assuming that light consists of particles. • By the year 1900, there was enough evidence to say that light consists of waves.
- 35. The wavelength of light are inversely proportional to each other.
- 36. Calculating the wavelength of light • Practice Problems on Page 140. • When atoms absorb energy, electrons move into higher energy levels, and these electrons lose energy by emitting light when they return to lower energy levels.
- 37. Atomic Emission Spectrum • Each specific frequency of visible light emitted corresponds to a particular color. When the light passes through a prism, the frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element. • The emission spectrum of each element is like a person’s fingerprint.
- 39. Explanation of Atomic Spectra • The lowest possible energy of the electron is the ground state. The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electrons. E = h x v h= 6.626 x 10^ -34 J.s
- 41. The Quantum Mechanical Model • After Max Planck determined that energy is released and absorbed by atoms in certain fixed amounts known as quanta, Albert Einstein took his work a step further,determining that radiant energy is also quantized—he called the discrete energy packets photons. Einstein’s theory was that electromagnetic radiation (light, for example) has characteristics of both a wave and a stream of particles.
- 42. Quanta of Light
- 43. Quantum Mechanics • Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves. • The Heisenberg Uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time.