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Periodicity Unit 2: Ch. 6 Coral Gables Senior High Ms. Kiely Pre-IB Chemistry I Bell-Ringer #13 a) b) c) d) A B C D Bell-Ringer #13 Answer: b) B TEST Oct 10 and 11 Quiz Review Periodicity: Atomic Radii Atomic Radius is not as simple as defining the radius in geometry class! Remember, the electron cloud doesn’t have a definite edge. Energy levels and orbitals are not at fixed distances from the nucleus. Atomic Radius: half the distance between the nuclei of two bonded identical atoms. Explanations for trends: 1. Nuclear charge and the, therefore, effective nuclear charge experienced by valence electrons (explains why trend increases across period but decreases down group) 2. Shielding of the valence electrons (explains specifically why trend decreases down a group but increases across a period) Across a Period: Atomic Radii Trends -Atomic radii decreases left to right across a period because the attraction between the nucleus and the valence electrons increases as the nuclear charge increases, causing the atom to “shrink” inwards. Remember that within a period, all elements have the same amount of energy levels. Down a Group: Atomic Radii Trends -Atomic radii increases down a group as the number of occupied energy levels increases. -As you go down a group, each element has an additional energy level than the one above it. Indeed the atoms get bigger because of more protons and neutrons in the nucleus just like they do across a period, but the additional energy level explains why “shrinking” isn’t a factor here- the shielding effect is at play here. Nuclear Charge: the amount of positive charge in an atom’s nucleus; given by the number of protons in the nucleus; atomic number. It increases by one between successive elements in the periodic table, as a proton is added to the nucleus. Effective Nuclear Charge: how “effective” the nuclear charge is at attracting or pulling its own valence electrons. Shielding effect: inner electrons, electrons that are positioned between the nucleus of an atom and the valence electrons of that atom; inner electrons repel the valence electrons away from the nucleus. The shielding effect causes the nuclear charge of an atom to not be as effective as it could be in attracting its own electrons. These concepts explain why the radii of atoms of elements down a group increase while across a period the radii decrease. Does the effective nuclear charge increase across a period? Why does the shielding effect remain constant across a period? Why does the shielding effect increase down a group? Periodicity: Ionic Radii Ionic Radius: 1) Positive ions (cations) are smaller than their parent atoms. 2) Negative ions (anions) are larger than their parent atoms. Trends in radii size are the same as for regular atomic radii. Periodicity: Ionic Radii CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so the radius DECREASES. ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Answers