chapter 7 periodic properties n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Chapter 7 – Periodic Properties PowerPoint Presentation
Download Presentation
Chapter 7 – Periodic Properties

play fullscreen
1 / 39

Chapter 7 – Periodic Properties

105 Views Download Presentation
Download Presentation

Chapter 7 – Periodic Properties

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Chapter 7 – Periodic Properties

  2. The Periodic Law • When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties. • Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. • They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.

  3. ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

  4. Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity

  5. Effective Nuclear Charge • What keeps electrons from simply flying off into space? • Effective nuclear charge is the pull that an electron “feels” from the nucleus. • The closer an electron is to the nucleus, the more pull it feels. • As effective nuclear charge increases, the electron cloud is pulled in tighter.

  6. Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by the outermost electrons. • E(2s) < E(2p) • Z* increases across a period owing to incomplete shielding by inner electrons. • Estimate Z* = [ Z - (no. inner electrons) ] • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 • Be Z* = 4 - 2 = 2 • B Z* = 5 - 2 = 3 and so on!

  7. Effective Nuclear Charge, Z* • Atom Z* Experienced by Electrons in Valence Orbitals • Li +1.28 • Be ------- • B +2.58 • C +3.22 • N +3.85 • O +4.49 • F +5.13 Increase in Z* across a period

  8. Orbital Energies Orbital energies “drop” as Z* increases ChemNow Screens 8.9 - 8.13, Simulations

  9. Shielding • As more Energy levels are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. • The effective nuclear charge (enc) on those outer electrons is less, and so the outer electrons are less tightly held. • Shielding increases as you go down the column and stays the same across the period.

  10. 99 pm 198 pm Atomic Size • The radius of an atom is defined by the edge of its last energy level. • However, this boundary is fuzzy • An atom’s radius is the measured distance between the nuclei of 2 identical atoms chemically bonded together - divided by 2. • It is measured in picometers (pm) or angstroms (Å)

  11. Atomic Size • Size goes UP on going down a group. See Figure 7.8. • Because electrons are added further from the nucleus, there is less attraction. • Size goes DOWN on going across a period.

  12. Increase in Z* Atomic Size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge. Large Small

  13. Atomic Radii See Active Figure 7.8

  14. Trends in Atomic SizeSee Active Figure 7.8

  15. Sizes of Transition ElementsSee Figure 7.9

  16. Sizes of Transition ElementsSee Figure 7.9 • 3d subshell is inside the 4s subshell. • 4s electrons feel a more or less constant Z*. • Sizes stay about the same and chemistries are similar!

  17. Ion Sizes Does the size go up or down when losing an electron to form a cation?

  18. + + Li , 78 pm 2e and 3 p Ion Sizes Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • The electron/proton attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p

  19. Ion Sizes Does the size go up or down when gaining an electron to form an anion?

  20. - - F, 71 pm F , 133 pm 9e and 9p 10 e and 9 p Ion Sizes Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The electron/proton attraction has gone DOWN and so size INCREASES. • Trends in ion sizes are the same as atom sizes.

  21. Trends in Ion Sizes See Active Figure 7.12

  22. Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg2+ ions and not Mg3+? Why do nonmetals take on electrons?

  23. Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. PLAY MOVIE Mg (g) + 738 kJ f Mg+ (g) + e-

  24. Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ f Mg+ (g) + e- PLAY MOVIE Mg+ (g) + 1451 kJ f Mg2+ (g) + e- Mg+ has 12 protons and only 11 electrons. Therefore, IE for Mg+ > Mg.

  25. Ionization Energy Mg (g) + 735 kJ f Mg+ (g) + e- Mg+ (g) + 1451 kJ f Mg2+ (g) + e- PLAY MOVIE Mg2+ (g) + 7733 kJ f Mg3+ (g) + e- Energy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no.

  26. Trends in Ionization Energy See Active Figure 7.10

  27. Trends in Ionization Energy

  28. Trends in Ionization Energy • IE increases across a period because Z* increases. • Metals lose electrons more easily than nonmetals. • Metals are good reducing agents. • Nonmetals lose electrons with difficulty.

  29. Trends in Ionization Energy • IE decreases down a group • Because size increases. • Reducing ability generally increases down the periodic table. • See reactions of Li, Na, K

  30. Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy involved when an atom gains an electron to form an anion. A(g) + e- f A-(g) E.A. = ∆U If the element becomes more stable after gaining electron, energy is released (negative).

  31.     -  [He] O ion + electron       [He] O atom Electron Affinity of Oxygen ∆U is EXOthermic because O has an affinity for an e-. EA = - 141 kJ

  32.     [He] N atom + electron   N- ion    [He] Electron Affinity of Nitrogen Difficult due to electron-electron repulsions. EA = 0 kJ

  33. Trends in Electron Affinity See Active Figure 7.11

  34. Trends in Electron Affinity • Affinity for electron increases across a period (EA becomes more negative). • Affinity decreases down a group (EA becomes less negative or positive). Atom EA F -328 kJ Cl -349 kJ Br -325 kJ I -295 kJ Note effect of atom size on F vs. Cl

  35. Electronegativity • Electronegativity is a tendency of an element to attract electrons towards itself in a compound. • Electronegativity is a measure of an atom’s attraction for another atom’s electrons. • It is an arbitrary scale that ranges from 0 to 4. • F is the most electronegative element with an electronegativity of zero. • Generally, metals are electron givers and have low electronegativities. • Nonmetals are are electron takers and have high electronegativities. • The trends in electronegativity are same as electron affinity.

  36. Electronegativity Table

  37. Metallic Character • This is simple a relative measure of how easily atoms lose or give up electrons. • Your help sheet should look like this:

  38. Periodic Trend in the Reactivity of Alkali Metals with Water PLAY MOVIE Lithium PLAY MOVIE PLAY MOVIE Sodium Potassium

  39. Shielding remains the same