Chapter 5 - Electron Configurations
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The document discusses the evolution of atomic models from Dalton to Bohr, including key discoveries by Thomson, Rutherford, and Bohr. It then explains electron configurations, describing the organization of electrons into energy levels, subshells, and orbitals according to quantum numbers and rules like the Aufbau principle and Hund's rule. The document provides examples of writing electron configurations using orbital, electron-configuration, and electron-dot notations.
Chapter 5 - Electron Configurations
- 1. Chapter 5 Notes By Kendon Smith Columbia Central High School, Brooklyn, MI Electron Configurations: Electrons move about the nucleus in a highly organized system!
- 2. Evolution of Atomic Models 1. Dalton Model (1803) – John Dalton pictured atoms as tiny and indestructible with no internal structure – tiny, solid spheres! DALTON MODEL
- 3. Evolution of Atomic Models 2. Thomson Model (1897) – J. J. Thomson discovered the electron with his Cathode ray tube experiment. – Thomson pictured atoms as spheres of positive charge embedded with negatively charged electrons. – Also called the Plum Pudding model, electrons are stuck in atoms “like raisins in plum pudding.”
- 4. Evolution of Atomic Models 2. Thomson Model (1897) Electrons (-) Positive Matrix (+)
- 5. Evolution of Atomic Models 3. Rutherford Model (1911) – Ernest Rutherford discovered a solid core called the nucleus with his gold foil experiment. – The nucleus is tiny, dense, and positively charged! – Electrons move around the nucleus in what is mostly empty space! * Rutherford’s atomic model could not explain
- 6. Evolution of Atomic Models 3. Rutherford Model (1911) Electron Cloud (-) e- e- e- Nucleus (+) + e- e- e-
- 7. Evolution of Atomic Models 4. Bohr Model (1913) - Neils Bohr proposed that electrons move around the nucleus in circular paths, or orbits, which are located at fixed distances from the nucleus. - Each electron orbits have fixed energies, so these possible energies are called energy levels. - Electrons can jump from one energy level to another by gaining or losing the right
- 8. Evolution of Atomic Models 4. Bohr Model (1913) Electrons (-) Nucleus (+) +
- 9. Evolution of Atomic Models - A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level. - In general, the higher energy levels are farther from the nucleus! - The amount of energy an electron loses or gains is not always the same. - Higher energy levels are closer together, requiring less energy to move electrons.
- 10. Ex. The Stair Step Analogy More Energetic Electrons Stairs = Energy Levels 6 Less Energetic 5 4 Increasing Electrons Energy! 3 2 1 *An electron needs less energy to jump from Level 4 to level 5, than from level 1 to level 2
- 11. Evolution of Atomic Models 5. Quantum Mechanical Model - The quantum mechanical model uses mathematical equations to describe the behavior of electrons within the electron cloud. - Does not define an exact path the electron takes around the nucleus, but rather, electron location is described as a fuzzy cloud where it spends most of its time. - It calculates the probability of finding an electron in a certain position.
- 12. Ex 1. The Dart Board Ex 2. Propeller Blades Where is the fan blade What are the chances my next when it’s moving? dart will hit a particular ring?
- 13. II. Electron Configurations A. Electron configurations are the ways in which electrons are arranged in various orbitals around the nucleus. - Electrons move about the electron cloud in a highly ORGANIZED system. Example: What are the levels of organization in your address? How does a letter find you when it is mailed?
- 14. II. Electron Configurations B. Levels of organization 1. Principle Quantum Number (Energy Level) a. Indicates the main ENERGY LEVEL surrounding a nucleus. b. Symbol = n Ex. Hydrogen – n = 1 Lithium – n = 2 3 Sodium – n =
- 15. 2. Orbital Quantum Number (Orbitals) a. Also called: SUBLEVELS or SUBSHELLS b. Symbols: s, p, d, f c. Indicates the SHAPE of an orbital. - s orbital = SPHERE - p orbital = PEANUT - d orbital = DOUBLE PEANUT - f orbital = FAR TOO COMPLEX
- 16. s orbital = sphere
- 17. p orbital = peanut
- 18. d orbital = double peanut
- 19. f orbital = “far too complex”
- 20. d. The n energy level has n subshells - 1st energy level has 1 subshells - 2nd energy level has 2 subshells - 3rd energy level has 3 subshells - 4th energy level has 4 subshells Principle Quantum Number Types of Orbitals 1 1s 2 2s 2p 3 3s 3p 3d 4 4s 4p 4d 4f
- 21. C. Spin Quantum Number a. Has only TWO possible values. b. Indicates direction of ELECTRON SPIN. c. Electrons spin on an imaginary axis, much like the earth, thus generating a MAGNETIC field. d. Symbols = ↑ ↓ e. Due to magnetism, electrons with OPPOSITE spins pair up in each orbital, or subshell. f. Each orbital is made up of an ELECTRON PAIR.
- 22. C. Rules governing electron configurations 1. Aufbau principle a. An electron occupies the LOWEST ENERGY orbital that is available to receive it.
- 23. C. Rules governing electron configurations 2. Hund’s rule a. Orbitals of equal energy are each occupied by one electron before any one orbital is occupied by a SECOND ELECTRON. - For example, you would put one electron in each p orbital, then come back and add the second electron. b. In orbitals that have only one electron, they must have the SAME SPIN.
- 24. C. Rules governing electron configurations 3. Pauli Exclusion Principle a. An atomic orbital may describe at most two electrons. b. Electrons that occupy the same orbital must have opposite spins.
- 25. Electron States 1. Electrons in their lowest energy level are said to be in the ground state. 2. By adding energy to an atom, an electron can jump into a higher energy level. This is called the excited state. 3. Almost immediately, the electrons fall back to their original ground state, giving off the added energy as visible light or radiation.
- 26. Electron States Atom becomes excited… Atom goes back to ground state. E3 E3 E2 E2 E1 E1 + + Nucleus Nucleus Energy LIGHT!
- 27. Some basic rules for electron configurations. 1. Atoms are in their ground state. 2. Atoms are neutral, having the same number of protons and electrons. 3. Lowest energy orbitals fill first. 4. Each orbital can hold a pair of electrons with opposite spin. 5. Completely fill all orbitals of the same energy before starting to fill the next level.
- 28. Orbital types and quantities: 1. There is one s orbital in each energy level. - holds a total of 2 electrons. 2. There are three p orbitals in each energy level. - holds a total of 6 electrons. 3. There are five d orbitals in each energy level. - holds a total of 10 electrons. 4. There are seven f orbitals in each energy level. - holds a total of 14 electrons.
- 29. Energy levels and their orbitals. - The n energy level contains n types of orbitals. Energy Level Orbitals 1 1s 2 2s 2p 3 3s 3p 3d 4 4s 4p 4d 4f 5 5s 5p 5d 5f 6 6s 6p 6d 6f 7 7s 7p 7d 7f
- 30. Question: In what order do we fill orbitals? Answer: Lowest energy orbitals first! The order: 7p Mr. Smith says, 6d “You MUST 7s 5f 6p memorize this 5d order!!!” 6s 4f 5p 4d 4p 5s You say, 4s 3d “How in the world 3s 3p am I supposed to 2p 1s 2s memorize this?”
- 31. Use this chart to memorize the order to fill the orbitals! Simply draw s p d f diagonal arrows starting with the 1 1s 1s orbital. 2 2s 2p Follow the arrows 3 3s 3p 3d to fill orbitals in each energy level. 4 4s 4p 4d 4f 5 5s 5p 5d 5f Jump from the head of one arrow 6 6s 6p 6d 6f to the tail of the 7 7s 7p 7d 7f next.
- 32. 3 Types of notations for configurations. 1. Orbital notation: a. An unoccupied orbital is represented by: b. An orbital occupied by a single electron: ↑ c. An orbital occupied by an electron pair: ↑↓ d. The lines are labeled with the proper principal quantum number and subshell letter. Ex. Sulfur ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ 1s 2s 2p 3s 3p
- 33. 2. Electron-Configuration Notation: a. Uses no lines or arrows. b. The number of electrons is shown as superscripts. Ex. Sulfur Orbital ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ Notation 1s 2s 2p 3s 3p Electron Configuration 1s2 2s2 2p6 3s2 3p4 Notation
- 34. 3. Electron-Dot Notation: a. Electron-dot notation shows only the electrons in the highest energy level or valence shell. b. Valence electrons – only those electrons located in the outermost energy level. c. Valence electrons are represented by dots placed in pairs around the element symbol. d. Most atoms can have 8 electrons in their valence shell before the begin filling the next level. Exceptions: Hydrogen and Helium – the first energy level only has an s orbital so it fills up with only 2 electrons!
- 35. Example: Sulfur Orbital ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ Notation 1s 2s 2p 3s 3p Electron Configuration 1s2 2s2 2p6 3s2 3p4 Notation 6 valence Electron Dot Notation electrons in S the 3rd energy level *Note: Place single electrons before you begin pairing them.
- 36. More Examples: Hydrogen: ↑ 1s1 H 1 e- 1s Lithium: ↑↓ ↑ 1s2 2s1 Li 3 e- 1s 2s Sodium: ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ 11 e- 1s 2s 2p 3s 1s2 2s2 2p6 3s1 Na
- 37. The End