Chemistry Unit 4
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31 terms
Terms | Definitions |
|---|---|
 Define Shielding | A barrier made of inner-shell electrons which serves to decrease the pull of the nucleus on the outer electrons |
| Trend for shielding: Across | Constant |
| Trend for shielding: Down | Increase |
| Define atomic radius | The size of an atom. Half the distance between two nuclei in a diatomic molecule. |
| Trend for atomic radius: Across | decrease |
| Trend for atomic radius: down | increase |
| Explain the trends for atomic radius at the atomic level. | Across a period there are more protons pulling on the same number of energy levels. Down a group there are more energy levels and more shielding so the outermost electrons aren't held as tightly. |
| Define electronegativity | Tendency of an atom to attract BONDED electrons closer to itself. |
| Trend for electronegativity: Across | Increase |
| Trend for electronegativity: Down | decrease |
| Explain the trends for electronegativity at the atomic level. | Across a period, more protons means more attraction. Down a group more energy levels and more shielding means less attraction. |
| Define ionization energy | Amount of energy require to remove an electron |
| Trend for ionization energy: Across | increase |
| Trend for ionization energy: Down | decrease |
| Explain the trends for ionization energy at the atomic level. | Across a period more protons means more attraction making it more difficult to remove an electron. Down a group more energy levels and more shielding means the outer electrons aren't held as tightly, therefore it takes less energy to remove. |
| Explain the fall-back in ionization energy between groups 2A and 3A. | It is some added stability associated with the full s sublevel in group 2A therefore it is actually easier to remove the lone electron from the p sublevel in group 3A than it is to remove one of the electrons from s in 2A. |
| Explain the fall-back in ionization energy between group 5A and 6A. | In 5A, each p electron gets its own room/orbital. In 6A the pairing of electrons in the p sublevel begins because of electron repulsion, it is easier to remove a paired electron from 6A. |
| For multiple ionization energies, how do you predict where the largest JUMP will occur? | Usually after the noble gas configuration. The electron that comes from the next lower energy level. |
| Define electron affinity | The amount of energy absorbed or given off when an atom accepts an electron. |
| Trend for electron affinity: Across | Increase |
| Trend for electron affinity: Down | decrease |
| Explain the trend for electron affinity at the atomic level. | Across a period there are more protons and more attraction, therefore more energy is given off as the atom accepts an electron. Down a group there are more energy levels and shielding. There is less pull by the nucleus and the values get less negative. |
| Define reactivity | How easily a metal atom loses its electrons or how easily a non-metal atom gains electrons. |
| Trend for reactivity (metals): Across | decrease |
| Trend for reactivity(metals): Down | increase |
| Explain the trends for reactivity at the atomic level. | Since metals lose electrons in chemical reactions, those atoms that can easily lose electrons (low ionization energy) are most reactive. Since non-metals gain electrons in chemical reactions, those atoms that can easily gain electrons (high electronegativity/electron affinity) are most reactive. |
| Trend for reactivity (nonmetal): Across | Increase |
| Trend for reactivity (nonmetal): Down | decrease |
| How does ionic radius of a metal cation compare to its neutral atom? | smaller, goes down an energy |
| How does ionic radius of a nonmetal anion compare to its neutral atom? | Larger, more electron repulsion |
| State the trend for melting point of metals and nonmetals. | Melting points for metals generally decrease down a group. Melting points for nonmetals generally increase down a group. |
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