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Periodic Table and Periodic Trends. Test 2 Material Along with Atomic Structure: 16 out of 75 MC questions Free Response every year. Broad Periodic Table Classifications. Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)
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Periodic Table and Periodic Trends Test 2 Material Along with Atomic Structure: 16 out of 75 MC questions Free Response every year
Broad Periodic Table Classifications • Representative Elements(main group): filling s and p orbitals (Na, Al, Ne, O) • Transition Elements: filling dorbitals (Fe, Co, Ni) • Lanthanide and Actinide Series(inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)
Periodicity • Trends in atomic radii, ionization energies, and electron affinities are all tied to two basic structural features of the atoms: • Effective nuclear charge (Zeff) • Principal energy level of valence electrons
Periodic Trends • Explained in terms of: • Increasing/Decreasing nuclear charge (Z) • Increasing/Decreasing distance from nucleus (n) • Increasing/Decreasing shielding by core electrons
Ionization Energy The quantity of energy required to remove an electron from the gaseous atom or ion—form a positive ion.
Low ionization energy—electron is easy to remove • High ionization energy—electron is difficult to remove
Periodic Trends • First ionization energy: • increasesfrom left to right across a period; higher nuclear charge • decreases going down a group; electrons farther from nucleus
Values of first ionization energy for the elements in the first six periods Page 326
First Ionization Energy • Notice that the trend is not smooth. • First ionization energy decreases between s- and p- orbital filling and just after p-orbital is half full. • Extra stability at these points
Trends in ionization energies (kJ/mol) for the representative elements. p. 326
Beyond First Ionization • Each successive electron is harder to remove—large + charge in relation to # of electrons • Large increase in ionization energy occurs when going to a lower energy level (n changes)
Electron Affinity The energy change in energy associated with the addition of an electron to a gaseousatom. X(g) + e X(g)
Electron Affinity • More exothermic for atoms that easily accept electrons; energy is more negative • Generally becomes more negative from left to right across a period—several exceptions such as adding to a 2p orbital with one electron already there (N vs C)
Electron affinity values for atoms among the first 20 elements that form stable, isolated X- ions.
Trend in Electron Affinity Becomes less exothermic down a group
Periodic Trends • Atomic Radii: • decrease going from left to right across a period; more positive charge in nucleus • increasegoing down a group; more energy levels
The radius of an atom (r) = half the distance between the nuclei in a molecule consistingof identical atoms.
Trends in Transition Elements Sc Ni Ti V Cr Mn Fe Co • Although there is a slight contraction at the beginning of the series, the atoms are all approximately the same size.
Ions Formed • Varies by group • Very predictable for representative elements • +1 +2 +3 +/-4 –3 –2 –1 0 • Lose or gain electrons to reach a Noble gas configuration
Sizes of Ions • Size increases down a group; more energy levels • In isoelectric ions, size decreases as positivity increases; more + nucleus holds electrons closer. • Positive ions are much smaller than negative ions—nuclear attraction; less repulsion
Group 1A—Alkali Metals • 1 valence electron • Very reactive—very close to stable configuration • Lose electron to form +1 ion • Donate electrons—reducing agents • Reactivity increases down the group
Group 2A—Alkaline Earth Metals • 2 valence electrons • Very reactive, but less than alkali metals—not as close to stable config. • Lose 2 electron to form +2 ion • Donate electrons—reducing agents • Reactivity increases down the group
Group 3A • 3 valence electrons (ns2, np1) • Typically forms +3 ions • Increasingly metallic down the group • Forms bonds with more covalent character than groups 1 or 2
Transition Metals • Show more similarity than rep. ele. both within periods and groups • Electrons added in d & f orbitals are actually inner electrons & cannot bond as easily as the outer s & p electrons—chemical properties do not change as much.
Transition Metals—General Properties • Typical metals: shiny, good conductors, flexible (malleable & ductile) • Melting points, hardness, and reactivity vary, so transition metals have varied uses.
Uses forTransition Metals • Hg—low melting point + good conductivity—thermometers and thermostats • W—high melting point + ductility—filaments in lightbulbs
Uses forTransition Metals • Fe & Ti—strength + hardness—structural materials • Cu, Ag & Au—soft—jewelry and art • Cu—low resistance—wiring
Reactions with Transition Metals • Most react with O2 to form oxides • Cr, Ni, Co—oxides protect surface • Fe—oxide flakes off exposing more metal to oxidation • Au, Ag, Pt & Pd—Noble metals—do not readily form oxides
Ionic Compounds with Transition Metals • Often paramagnetic (unpaired e-s) • Usually more than one oxidation state • Cations are often complex ions in which ligands (Lewis base ions or molecules) surround metal ion
Ionic Compounds with Transition Metals Most are colored because ions absorb specific wave-lengths of light. Electrons in split d orbitals can rearrange.
Electron Configuration • Full (n)s orbital; filling (n-1)d orbital; empty (n)p orbital • Exception: Cu—4s1, 3d10
Breaking the Rules • Since the energy of a 3d orbital is less than that of a 4s orbital, ions formed from first row transition metals lose their 4s electrons first • Ex: Mn = [Ar] 4s2, 3d5 Mn2+ = [Ar] 3d5
Lanthanides & Actinides Technically part of the transition metals Filling f orbitals with electrons Rare metals with few uses All actinides are radioactive—nuclear fuels (U & Pu)
Bonding 8 out of 75 M/C Questions Free Response—Every year
In General: • All bonds occur because of electrostatic attractions. • Formation of molecules and the state of matter of a substance depends on the attractions between electron clouds of one atom and nucleus of another atom.
Bonding—General Rules • A metal and a nonmetal bond so that charges on the ions cancel. • When two nonmetals react to form a covalent bond, they share electrons in a way that gives both atoms a Noble gas configuration.
Bonding—General Rules • nonmetal + representative metal binary ionic compound: ions form to give the nonmetal the valence electron configuration of the next noble gas atom, and valence orbitals of the metal are emptied
Bonding—General Rules • Most bonds are combinations of ionic/covalent character • The more different the atoms bonding are, the more ionic character of the bond. • More similar—more covalent
Covalent Sharing Molecules Structural formula Ionic Transfer / charged Compounds Formula Unit Word Association
Coulomb’s Law • Describes energy of interaction between ions • E = 2.31 x 10-19 J*nm(Q1Q2 / r) • Q1 & Q2numerical ion charges • r distance between ion centers • Negative ans. means ion pair is more stable than individual ions.
Bond Length • Distance between bonding atoms at which energy is minimized • Atoms position themselves to minimize repulsions and maximize attraction & thus achieve lowest possible energy.
Predicting Formulas for Ionic Compounds • Metal—positive charge equal to # of valence electrons • Nonmetal—negative charge equal to # of electrons away from next Noble gas • Compound—charges must cancel
Predict Formulas for: • Potassium sulfide • Barium chloride • Aluminum oxide • Magnesium phosphide
Exceptions to the Rule: • Sn—forms both +2 and +4 ions • Pb—forms both +2 and +4 ions • Bi—forms +3 and +5 ions • Tl—forms +1 and +3 ions “no simple explanation for this behavior”