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Properties of Elements

Discover the significance of atomic radius, ionization energy, and electronegativity in determining an element's properties. Explore trends across periods and groups, as well as the impact of gaining or losing electrons on ionic radius. Learn about the characteristics and behavior of various groups of elements.

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Properties of Elements

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  1. Properties of Elements

  2. Atomic Radius A measure of the size on an atom. What determines an atom’s size? Remember, the nucleus is very very small and compact. It is the electrons that determine how big the atom is.

  3. Atomic Radius It is hard to measure where the moving e- are at any moment, so they can not be easily used to measure size. DEFINITION: Half the distance between the nuclei of two adjacent atoms in a crystal ½ distance

  4. Ionization Energy DEFINITION: The amount of energy it takes to remove the outermost e- from a neutral atom in the gas phase X + Ionization energy X+ + e- Neutral atom Cation electron

  5. Electronegativity DEFINITION: A measure of the attraction an atom has for electronswhen it is bonded to another atom. Scale is from 0.7 (low, Cs) to 4.0 (high, F)

  6. Trends in Atomic Radius Across a period: radiusdecreases because there are more protons in each successive atom’s nucleus, pulling harder on the e- and making the atom smaller Down a group: radiusincreases because the atoms have more energy levels farther from the nucleus, making the atom bigger

  7. Trends in Ionization Energy Across a period: I.E.increases because there are more protons in each successive atom’s nucleus, pulling harder on the e- and making it harder to remove the e- Down a group: I.E. decreases because the atoms have more energy levels farther from the nucleus, so the outer e- are less attracted to the nucleus and are therefore easier to remove. Also, inner e- “shield” the outer e- from the pull of the nucleus.

  8. Trends in Electronegativity Across a period: Electronegativityincreases because there are more protons in each successive atom’s nucleus, pulling harder on the e- Down a group: Electronegativitydecreases because the atoms have more energy levels farther from the nucleus, so the nucleus has less positive pull on the e-. Also, inner e- “shield” the outer e- from the pull of the nucleus

  9. Ionic Radius If an atom GAINS e-, it gets bigger in size So….negative ions (anions) are bigger than their neutral atom

  10. Ionic Radius If an atom LOSES e-, it gets smaller in size So… positive ions (cations) are smaller than their neutral atom

  11. Group 1 alkali metals • Electron configuration ends with S1 • Lose this outermost e- easily (lowIonization energy and electronegativity) forming +1 cations • VERY reactive! Francium is MOST reactive • Not found uncombined in nature • Form stable compounds with non metals like NaCl

  12. Group 2 alkaline earth metals • Electron configuration ends with S2 • Lose these 2 outermost e- easily (lowIonization energy and electronegativity) but not as easily as Group 1 metals losing only 1 e- • Form +2 cations • Reactive! (but not as much as Group 1) • Not found uncombined in nature • Form stable compounds with non metals like MgCl2

  13. Groups 3-12 transition metals • Highest energy level ends with S2 but d-orbitals are being filled • Tend to lose the S2e- easily, forming +2 cations, but many can also form +1 or +3 (multiple oxidation states) • Less reactive than Groups 1 or 2 • Form colorful ions and compounds

  14. Groups 13 • Electron configuration ends with S2P1 • Lose the three S2P1 e-, forming +3 cations • Both metalloids and metals in this group

  15. Groups 14 • Electron configuration ends with S2P2 • Don’t tend to form ions • Non metals, metalloids and metals in this group

  16. Groups 15 • Electron configuration ends with S2P3 • Tend to gain 3 e-, forming -3 anions • Non metals, metalloids and metals in this group

  17. Groups 16 • Electron configuration ends with S2P4 • Tend to gain 3 e-, forming -2 anions • Non metals, and metalloids in this group • Reactive! Tend to form stable compounds with metals like MgO

  18. Groups 17 Halogens • Electron configuration ends with S2P5 • Tend to gain one e-, forming -1 anions. Very highelectronegativity and ionization energy. (F is highest electronegativity with 4.0) • Non metals only in this group • Only group to have all three phases of matter at room temperature (s, l, g) • VERY Reactive! Not found uncombined in nature. Tend to form stable compounds with metals like NaCl. Most reactive is F.

  19. Electron configuration ends with S2P6 Energy level is full Do not lose or gain e-. Do not form ions. UNreactive! Not found combined with other elements in nature. Do not form compounds. Groups 18 Noble gases

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