1 / 62

Chapter 2

Chapter 2. Electrons and the Periodic Table CHE 101 Sleevi. Modern Atomic Theory. Early 1900’s – Birth of Quantum Mechanics 1913 – Bohr proposed fixed orbits for electrons based on energy Studied hydrogen emission spectrum Worked for single electron Model did not resolve all observations

alicel
Download Presentation

Chapter 2

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 2 Electrons and the Periodic Table CHE 101 Sleevi

  2. Modern Atomic Theory • Early 1900’s – Birth of Quantum Mechanics • 1913 – Bohr proposed fixed orbits for electrons based on energy • Studied hydrogen emission spectrum • Worked for single electron • Model did not resolve all observations • 1926 – Schrödinger proposed quantum mechanical model of the atom • First mathematical model 2

  3. Modern Atomic Theory

  4. Modern Atomic Theory Emission spectra due to change in energy levels of electrons: Ground State – electrons in atoms occupy lowest possible energy levels Excited State – electrons occupy higher energy levels by absorbing energy When e- falls to lower energy state, light is emitted

  5. Modern Atomic Theory Neon Helium Sunlight Hydrogen Continuous spectrum Discrete spectrum Figure 2.4

  6. Quantum Mechanical Model of the Atom Continuous Quantized 6

  7. Quantum Mechanical Model of the Atom • Principal energy levels (shells) • correspond to the discrete energy levels in the atom that can be occupied by electrons • designated by quantum number n (n = 1, 2, 3, 4, …) • contain sublevels • quantum number n designates number of sublevels at that energy level • Sublevels (subshells) • correspond to orbital type within a principal energy level (s, p, d, f) 7

  8. Orbitals • Solution to the Schrödinger equation • Shape describes the probability of where electron will be found in space, at that energy (90%) • Overlay the nucleus • Increase in size as the principal energy level increases • Contain no more than 2 electrons 8

  9. Orbitals Fig 2.8. (a) A burning sparkler produces a roughly spherical cloud of sparks. (b) Most (90%) of the sparks from this sparkler are located within a sphere with a 5 cm radius. (c) The sphere that contains 90% of the sparks represented by a solid surface.

  10. Electronic StructureOrbital Shapes • Thes orbital has a spherical shape. • Thep orbital has a dumbbell shape. 10

  11. Electronic StructureOrbital Shapes • Thed orbitals have multi-lobed shape. 11

  12. Energy Levels in an Atom 12

  13. Orbitals and Electrons by Sublevel 13

  14. Max Number of Electrons by Principal Energy Level 14

  15. Electronic StructureShells Shells with larger numbers (n) are farther from the nucleus, have a larger volume, and can therefore hold more electrons. Number of Electrons in a Shell Shell (n) 32 4 increasing number of electrons 3 18 increasing energy 2 8 1 2 15

  16. Electron Configurations • Arrangement of electrons around the nucleus • Based on set of rules that comes from Quantum Mechanical model of the atom 16

  17. Electron Configurations • Rules determine how orbitals are filled, based on number of electrons in the atom • Aufbau Principle • Pauli Exclusion Principle • Hund’s Rule 17

  18. Aufbau Principle • Electrons enter orbitals of lowest energy first • Filling order: 1s 2s 2p 3s 3p 4s 3d 4p 5s ….. • Number in front = quantum number (PEL) • Letter designates sublevel (orbital type) • Maximum number of electrons depends upon the orbital type 18

  19. Aufbau Diagram 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 19

  20. Aufbau Diagram 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 20

  21. Pauli Exclusion Principle • An orbital can contain 0, 1, or 2 electrons • Electrons have a property called spin • When there are two electrons in one orbital they have opposite spin 21

  22. Hund’s Rule • When electrons occupy orbitals of equal energy, one electron enters each orbital with parallel spin a. b. c. 22

  23. Following the rules… • Elements through Vanadium follow these rules • Most elements after that follow the rules • There are exceptions 23

  24. Writing Configurations • Electron Configuration Diagram (aka orbital diagram) • Use horizontal lines for orbitals, • Arrows for electrons, • Direction of arrows for spin • Label rows with PEL and orbital designation (1s, 2s, 2p, etc.) • Follow Aufbau Rule, Pauli Exclusion Principle, Hund’s rule 24

  25. Increasing Energy of Orbitals Energies of subshells overlap starting with 4s and 3d

  26. Examples Lecture Problem #1 Draw electron configuration diagrams for the following elements: • Atomic #9 • Magnesium • Element in Group 6A, Period 3 • Atomic Number 23 • Rubidium 26

  27. Electron Configuration Shorthand • Write PEL and orbital type as number, letter • Use superscript for number of electrons in the orbital 1s22s22p3 • Element can be identified by adding up the superscripts (# electrons) 27

  28. Electron Configurations 28 Noble Gas Notation = condensed configuration with Noble Gas Core

  29. Mapping Aufbau Diagram to the Periodic Table Electrons enter s orbitals (1A, 2A) Electrons enter p orbitals (3A – 8A) Electrons enter d orbitals (Transition metals) Electrons enter f orbitals (Inner transition metals) Principal energy level for s and p given by the row number 29

  30. Electron Configurationsand the Periodic Table The Blocks of Elements in the Periodic Table

  31. Shorthand and Noble GasElectron Configurations Lecture Problem #2 Write the shorthand notation and noble gas notation for the electron configurations for the elements in #1.

  32. Valence Electrons • Outermost electrons • Highest occupied principal energy level (largest quantum number) • Number of valence electrons = group number for the A groups (representative elements) • Determine the chemical properties of the elements • Determine number and type of bonds 32

  33. More Practice! Lecture Problems #3 & #4 3. Circle the portion of the shorthand configuration in #2 that represents the valence electrons. 4. Draw the valence shell configuration diagram for the element in each of the representative groups in Period 4. (symbol, shorthand configuration) 33

  34. Valence ElectronsElectron-Dot Symbols • Dots representing valence electrons are placed on the four sides of an element symbol. • Each dot represents one valence electron. • For 1–4 valence electrons, single dots are used. With > 4 valence electrons, the dots arepaired. H Element: O Cl C # of Valence electrons: 1 4 6 7 Electron-dot symbol: H C O Cl 34

  35. Practice! Lecture Problem #5: Draw Lewis dot symbols for each of the following: S P Sr C Br Kr Na Al H O B N

  36. Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties 36

  37. Periodic Trends • Properties have trends with respect to location of elements on periodic table • Atomic Size • Ionization Energy • Metallic character • Electron affinity 37

  38. Periodic TrendsMetallic Character • The metallic character of atoms increasesdown a column, as the valence e− are farther from the nucleus. Increases Decreases • The metallic character of atoms decreases across a row. 38

  39. Periodic TrendsAtomic Size • The size of atoms increasesdown a column, as the valence e− are farther from the nucleus. Increases Decreases • The size of atoms decreases across a row, as the number of protons in the nucleus increases. • The increasing # of protons pulls the e− closer to the nucleus, making the atoms smaller. 39

  40. Periodic TrendsIonization Energy The ionization energy is the energy needed to remove an electron from a neutral atom. Na + energy Na+ + e– Ionization energies decrease down a column as the valence e− get farther away from the positively charged nucleus. Decreases Increases Ionization energies increase across a row as the number of protons in the nucleus increases. 40

  41. Periodic TrendsElectron Affinity Electron affinity is the attraction of an atom to gain one or more electrons F + e–  F– No significant trend in electron affinity down a group Increases Electron affinity increases across a row. 41

  42. Examples! Lecture Problem #6: Arrange the atoms in order of increasing size • B, C, Ne • F, S, Al • Kr, Ne, Xe • Ca, Mg, Be

  43. Examples! Lecture Problem #7: Arrange the atoms in order of increasing ionization energy • P, Si, S • Ne, Kr, Ar • C, F, Be • Ca, Al, N

  44. Formation of Ions • Ion: charged particle formed from a neutral atom by adding or removing electrons. The number of protons remains the same. • Cation: positive ion, formed from metal atoms by removing one or more electrons • Anion: negative ion, formed from nonmetal atoms by adding one or more electrons • Ion formula: element symbol with the charge of the ion as a superscript (on the right)

  45. How do we determine the charge of the ion? Consider the valence electrons • Octet rule • ions are formed by the gain or loss of electrons to attain a noble gas configuration (8 e- in the valence shell)

  46. Formation of IonsCations and Anions Cations are positively charged ions. A cationhas fewer electrons (e−) than protons.

  47. Formation of IonsCations and Anions Anions are negatively charged ions. An anion has more e− than protons.

  48. A few points to remember… • Electron transfer only occurs when there is an atom to receive the electrons • Number of electrons gained or lost usually does not exceed three • Group 4A nonmetals do not usually form ions

  49. IonsRelating Group Number to Ionic Charge forGroups 1A–3A the cation charge = the group number group 1A: M 1 valence e− M+ + e− M 2 valence e− M2+ + 2e− group 2A: M 3 valence e− M3+ + 3e− group 3A:

  50. Valence Electrons and Ion Formation • Metals form cations

More Related