1 / 20

PERIODIC TRENDS

PERIODIC TRENDS. CHAPTER 5. ATOMIC RADII. Defined as: ½ the distance between the nuclei of two identical atoms joined in a molecule Approximates the size of the electron cloud. ATOMIC RADIUS. Trends: Across a period: DECREASES

haley
Download Presentation

PERIODIC TRENDS

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. PERIODIC TRENDS CHAPTER 5

  2. ATOMIC RADII • Defined as: • ½ the distance between the nuclei of two identical atoms joined in a molecule • Approximates the size of the electron cloud

  3. ATOMIC RADIUS • Trends: • Across a period: DECREASES • why? Electrons are added to the same energy level and the increased nuclear charge causes the electrons to be bound more effectively, reducing the size of the electron cloud. • Down a group: INCREASES • Why? Electrons are being added to higher energy levels. Electron cloud size increases as n increases.

  4. ATOMIC RADIUS

  5. IONIZATION ENERGY • Defined as: • The energy needed to remove an electron from an atom to form a +1 ion. • Measured in kJ or kJ/mol Ex. Na + IE  Na+1 + 1 e- • Ion – an atom with a + or - charge

  6. IONIZATION ENERGY • Trends: • Across a period: INCREASES • Electrons feel more nuclear pull when closer to the nucleus, so they are harder to remove. • Metals lose electrons easily/ non-metals hold electrons more.

  7. IONIZATION ENERGY • Down a group: DECREASES • Why? • e- are farther from the nucleus in higher energy levels, so they do not feel the nuclear charge so they are more easily removed. • Inner electrons shield the nuclear charge.

  8. IONIZATION ENERGY

  9. 2nd IONIZATION ENERGY • The 2nd ionization energy is always higher than the first ionization energy • Why? • Na+1 + 2nd IE  Na+2 + 1e-1 • More protons are pulling on less electrons, electrons under control are harder to remove.

  10. CHECK POINT • Put the following in order of increasing atomic radius: • Na, Cl, Cs, I • Put the following in order of increasing ionization energy: • Be, He,Ca, Br • Which requires more energy? The first ionization of Fe or the second ionization of Fe? Why?

  11. ELECTRON AFFINITY • Defined as: • The energy released or required when a neutral atom gains an electron. • Ex. Cl + 1e-  Cl-1 + energy • Ex. Na + 1e- + energy  Na-1

  12. ELECTRON AFFINITY TREND • Trends: • Across a period: INCREASES, becomes more favorable • Why? Smaller atoms with higher nuclear charge will attract e- more effectively, making it easier to add anoter electron. • Exception: noble gases EA = 0

  13. ELECTRON AFFINITY TREND • Down a group: Decreases, becomes less favorable • Why? In larger atoms, the nuclear charge is shielded, so it is more difficult to attract the additional electron. Energy is required to add the electron to the atom. • Metals always have lower EA than non-metals. • Order: metals (+), noble gases (0), non-metals (-) • Ex. Place the following in order of increasing favorability for EA: Si, Cl, Ar, Na

  14. ELECTRON AFFINITY

  15. ELECTRONEGATIVITY • Defined as: • A measure of the ability of a bonded atom to attract a pair of electrons. • Pauling Electronegativity scale ranges from 0-4 • Highest electronegativity: Fluorine

  16. ELECTRONEGATIVITY TREND • Trends: • Across a period: Gradually increases • Smaller atoms with increased nuclear charge are better e pair attractors. • Non metals always have a higher e-neg than metals. • Noble gases do not have an e-neg value.

  17. ELECTRONEGATIVITY TREND • Down a group: Decreases • Larger atoms cannot attract e pairs as effectively due to the nuclear charge being shielded by inner electrons. • Ex. Place the following in order of increasing electronegativity: Na,Cl, Ar, S

  18. ELECTRONEGATIVITY

  19. IONIC RADII • Cations – (+ ions) lost electrons, so a higher p+/e- ratio causes the electrons to be pulled closer, reducing the size of the electron cloud. • Anions- (- ions) gained electrons, so a lower p+/e- ratio means the electrons are under less control and the e- cloud expands. • Trend: • Going down a group: increases as the atomic radius increases

  20. ISOELECTRONIC SPECIES • Elements and ions with the same number of electrons. • Ex. Kr, Br-1, Sr+2, Rb+1 • Ex. Find ions that would be isoelectronicAr and place them in order of increasing ionic/atomic radius

More Related