Chapter 7 Periodic Properties of the Elements CHEMISTRY The Central Science 9th Edition Chapter 7 Periodic Properties of the Elements Prentice Hall © 2003 Chapter 7

7.1: Development of the Periodic Table Dimitri Mendeleev and Lothar Meyer arranged the elements in order of increasing atomic weight Certain elements were missing from this scheme: Mendeleev noted: As properly belonged underneath P and not Si A missing element underneath Si He predicted a number of properties for this element In 1886 Ge was discovered Properties of Ge match Mendeleev’s predictions Modern periodic table: arrange elements in order of increasing atomic number Prentice Hall © 2003 Chapter 7

Effective Nuclear Charge Effective nuclear charge is the charge experienced by an electron on a many-electron atom The effective nuclear charge is not the same as the charge on the nucleus Due to the effect of the inner electrons Prentice Hall © 2003 Chapter 7

7.2: Effective Nuclear Charge Electrons are attracted to the nucleus, but repelled by the electrons that screen them from the nuclear charge The nuclear charge experienced by an electron depends on its distance from the nucleus the number of core electrons As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases OR…as the distance from the nucleus increases, S increases and Zeff decreases Prentice Hall © 2003 Chapter 7

The ns orbitals all have the same shape, but have different sizes and different numbers of nodes (places where the probability function = 0) Zeff = Z – S or Protons – Electrons The radial electron density is the probability of finding an electron at a given distance Prentice Hall © 2003 Chapter 7

The period trend is of greater significance than the group trend Zeff increases across a period Added electrons don’t increase shielding Zeff increases down a group Added electrons are in increasingly higher energy levels but they are less able to shield the outer electrons from the nucleus The period trend is of greater significance than the group trend Prentice Hall © 2003 Chapter 7

7.3: Sizes of Atoms and Ions Consider a simple diatomic molecule: The distance between the two nuclei is called the bond distance Half of the bond distance is called the covalent radius of the atom Prentice Hall © 2003 Chapter 7

Periodic Trends in Atomic Radii The properties of elements vary periodically Atomic size varies consistently through the periodic table Atomic radius increases down a group (more noticeable) Atomic radius decreases across a period There are two factors at work: Principal quantum number Electrons fill higher energy levels Effective nuclear charge Electrons fill the same energy level, no increase in shielding Prentice Hall © 2003 Chapter 7

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Trends in the Sizes of Ions Ion size is the distance between ions in an ionic compound Cations vacate the most spatially extended orbital and are smaller than the parent atom Anions add electrons to the most spatially extended orbital and are larger than the parent atom Remember: electrons are gained or lost to achieve a noble gas configuration Prentice Hall © 2003 Chapter 7

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O2- > F- > Na+ > Mg2+ > Al3+ For ions of the same charge, ion size increases down a group Electrons in higher energy levels All the members of an isoelectronic series have the same number of electrons (O2-, F-, Na+, Mg2+, Al3+) As nuclear charge increases in an isoelectronic series, the ions become smaller: O2- > F- > Na+ > Mg2+ > Al3+ Added electrons are in the same energy level, so no increase in shielding Removed electrons result in an increasingly positive nucleus which pulls electrons toward it Prentice Hall © 2003 Chapter 7

7.4: Ionization Energy The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom: Na(g)  Na+(g) + e- The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion: Na+(g)  Na2+(g) + e- The larger ionization energy, the more difficult it is to remove the electron Prentice Hall © 2003 Chapter 7

Prentice Hall © 2003 Chapter 7

Variations in Successive Ionization Energies There is a sharp increase in ionization energy when a core electron is removed: Text, 268 Prentice Hall © 2003 Chapter 7

Think: does it want to lose an electron??? Periodic Trends in Ionization Energies Think: does it want to lose an electron??? Ionization energy decreases down a group The outermost electron is more readily removed as electrons fill higher energy levels The s electrons are more effective at shielding than p electrons Irregularity in group 2: sharp increase in I1 Prentice Hall © 2003 Chapter 7

Prentice Hall © 2003 Chapter 7 Text, P. 270

Ionization energy generally increases across a period Across a period, Zeff increases It becomes more difficult to remove an electron Irregularity in group 15: sharp increase in I1 (half-filled sublevel) Transition metals and f-block elements show slow variations in I1 due to shielding Prentice Hall © 2003 Chapter 7

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Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5) Electron Configuration of Ions Cations: electrons removed from orbital with highest principle quantum number, n, first: Li (1s2 2s1)  Li+ (1s2) Fe ([Ar]3d6 4s2)  Fe3+ ([Ar]3d5) Anions: electrons added to the orbital with highest n: F (1s2 2s2 2p5)  F- (1s2 2s2 2p6) Prentice Hall © 2003 Chapter 7

Think: does it want to gain an electron??? 7.5: Electron Affinities Think: does it want to gain an electron??? Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion: Cl(g) + e-  Cl-(g) Electron affinity can either be exothermic (above) or endothermic: Ar(g) + e-  Ar-(g) Prentice Hall © 2003 Chapter 7

EA becomes more negative across a period A greater attraction between the atom and the added electron Groups 2, 15 and 18 EA has no real group trend Text, P. 272

7.6: Metals, Nonmetals, and Metalloids Text, P. 274

Metallic character refers to the properties of metals shiny or lustrous malleable and ductile oxides form basic ionic solids form cations in aqueous solution Metallic character increases down a group Metallic character decreases across a period Metals have low ionization energies Most neutral metals are oxidized rather than reduced form cations Prentice Hall © 2003 Chapter 7

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Metal oxide + water  metal hydroxide Most metal oxides are basic: Metal oxide + water  metal hydroxide Nonmetals Nonmetals are more diverse in their behavior than metals Properties are opposite those of metals Compounds composed only of nonmetals are molecular When nonmetals react with metals, nonmetals tend to gain electrons: metal + nonmetal  salt Prentice Hall © 2003 Chapter 7

Most nonmetal oxides are acidic: nonmetal oxide + water  acid Nonmetal oxide + base  salt + water The greater the nonmetallic character of the central atom, the stronger the acidic character of the oxide Metalloids Metalloids have properties that are intermediate between metals and nonmetals Si has a metallic luster but it is brittle Semiconductors Prentice Hall © 2003 Chapter 7

7.7: Group Trends for the Active Metals Group 1A: The Alkali Metals Alkali metals are all soft Chemistry dominated by the loss of their single s electron: M  M+ + e- Reactivity increases down the group Alkali metals react with water to form MOH and hydrogen gas: 2M(s) + 2H2O(l)  2MOH(aq) + H2(g) Prentice Hall © 2003 Chapter 7

Alkali metals produce different oxides when reacting with O2: 4Li(s) + O2(g)  2Li2O(s) (oxide) 2Na(s) + O2(g)  Na2O2(s) (peroxide) K(s) + O2(g)  KO2(s) (superoxide) (Rb and Cs, too) Alkali metals: Flame tests The s electron is excited by the flame and emits energy when it returns to the ground state Prentice Hall © 2003 Chapter 7

Na line (589 nm): 3p  3s transition Li line: 2p  2s transition K line: 4p  4s transition Prentice Hall © 2003 Chapter 7

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M(s) + 2H2O(l)  M(OH)2(aq) + H2(g) Group 2A: The Alkaline Earth Metals Alkaline earth metals are harder and more dense than the alkali metals The chemistry is dominated by the loss of two s electrons: M  M2+ + 2e- Mg(s) + Cl2(g)  MgCl2(s) Increasing activity down the group: Be does not react with water Mg will only react with steam Ca, Sr and Ba: M(s) + 2H2O(l)  M(OH)2(aq) + H2(g) Prentice Hall © 2003 Chapter 7

Group 2A: The Alkaline Earth Metals Text, P. 281 Prentice Hall © 2003 Chapter 7

7.8: Group Trends for Selected Nonmetals Hydrogen Colorless diatomic gas, H2 High IE because there are no inner electrons It can either gain another electron to form the hydride ion, H-, or lose its electron to become H+: 2Na(s) + H2(g)  2NaH(s) 2H2(g) + O2(g)  2H2O(g) H+ is a proton The aqueous chemistry of hydrogen is dominated by H+(aq) Prentice Hall © 2003 Chapter 7

Group 16: The Oxygen Group Metallic character increases down a group (O2 is a gas, Te is a metalloid, Po is a metal) Two important forms of oxygen: O2 and ozone, O3 Ozone can be prepared from oxygen: 3O2(g)  2O3(g) H = +284.6 kJ. Ozone is pungent and toxic Prentice Hall © 2003 Chapter 7

Group 16: The Oxygen Group Text, P. 283 Prentice Hall © 2003 Chapter 7

Oxygen (O2) is a potent oxidizing agent O2- ion has a noble gas configuration There are two oxidation states for oxygen: 2- and 1- Sulfur is another important member of this group Most common form of sulfur is yellow S8 Sulfur tends to form S2- in compounds (sulfides) Allotropes: different forms of the same element in the same state Prentice Hall © 2003 Chapter 7

Fluorine is one of the most reactive substances known Group 17: The Halogens The chemistry of the halogens is dominated by gaining an electron to form an anion: X2 + 2e-  2X- Fluorine is one of the most reactive substances known All halogens consists of diatomic molecules, X2 Prentice Hall © 2003 Chapter 7

Group 17: The Halogens Text, P. 285 Prentice Hall © 2003 Chapter 7

2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g) Chlorine is the most industrially useful halogen Produced by the electrolysis of brine (NaCl): 2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g) Hydrogen compounds of the halogens are all strong acids with the exception of HF Prentice Hall © 2003 Chapter 7

These are all nonmetals and monatomic Group 18: The Noble Gases These are all nonmetals and monatomic Unreactive because they have completely filled s and p sub-shells In 1962 the first compound of the noble gases was prepared: XeF2, XeF4, and XeF6 Xe has a low enough IE to react with substances that readily remove electrons To date the only other noble gas compounds known are KrF2 and HArF “Argon hydrofluoride” Prentice Hall © 2003 Chapter 7

Group 18: The Noble Gases Text, P. 286 Prentice Hall © 2003 Chapter 7