Download Chapter 7. Periodic Properties of the Elements

C H A P T E R 7 . P E R I O D I C P RO P E RT I E S O F T H E E L E M E N T S
Student LEARNING OUTCOMES
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Be able to know how position on periodic table & electron structure is used to highlight periodic properties
Know that the periodic table is an organizational tool; will help recall chemical facts
Know why, within a period, atomic radii decrease with increasing atomic number
Be able to understand slight irregularities in periodic trends for elements in each row after each ns subshell
becomes filled, and after np and (n – 1)d subshells become half filled
• Be aware of problems with the signs of electron affinities—in particular, why group 1A metals have
negative (exothermic) electronegativities
• Be able to understand behavior of elements in aqueous phase with periodic properties determined in gas
phase (ionization energy, electron affinity) or in solid phase (ionic radius)
• Know what isoelectronic species are with those with the same number of valence electrons
CHAPTER 7 OUTLINE
7.1 Development of the Periodic Table
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Periodic table is most significant tool chemists use for organizing and recalling chemical facts
Elements in same column contain same number of outer-shell electrons or valence electrons
Arrange elements to reflect trends in chemical and physical properties
Periodic table arises from the periodic patterns in electronic configurations of elements
• elements in same column contain same number of valence electrons
• trends within row or column form patterns help make predictions of chemical properties and reactivity
• Mendeleev and Meyer arranged elements in order of increasing atomic weight
• certain elements were missing from this scheme
• example, 1871 Mendeleev noted As belonged underneath P and not Si, this left missing element
underneath Si, could then predicted a number of properties for this element
• in 1886 Ge was discovered; the properties of Ge match Mendeleev’s predictions well
• In the modern periodic table, elements arranged in order of increasing atomic number
7.2 Effective Nuclear Charge
• Effective nuclear charge (Zeff) is charge experienced by electron on a many-electron atom
• Effective nuclear charge not same as charge on nucleus because of effect of inner electrons
• electron attracted to nucleus, but repelled by electrons that shield or screen it from full nuclear charge
• Nuclear charge experienced by electron depends on distance from nucleus and number electrons in
spherical volume out to electron in question
• As average number of screening electrons (S) increases, effective nuclear charge (Zeff) decreases Zeff = Z – S
• As distance from nucleus increases, S increases and Zeff decreases
• S is screening constant which represents portion of nuclear charge that is screened from valence electron by
other electrons in the atom
• value of S usually close to number of core electrons in atom
7.3 Sizes of Atoms and Ions
• Radius is determined by closest distances separating the nuclei during such collisions; the nonbonding radius
• Nonbonding atomic radii also called van der Waals radii
• used in space-filling models to represent the sizes of different elements
• Diatomic molecule
• distance between two nuclei called bonding atomic radius, shorter than nonbonding radius
• if 2 atoms are the same, then half the bond distance is called covalent radius of atom
Periodic Trends in Atomic Radii
• Atomic size varies consistently through periodic table
• down a group atoms become larger
• across a period atoms become smaller
• 2 factors at work: principal quantum number, n, and effective nuclear charge, Zeff.
• as principal quantum number increases (down group), distance of outermost electron from nucleus
becomes larger, atomic radius increases
• as move across periodic table, number of core electrons remains constant; nuclear charge increases, is an
increased attraction between nucleus and outermost electrons, causes atomic radius to decrease
Periodic Trends in Ionic Radii
• Ionic size is important: predicting lattice energy and determining way ions pack in a solid
• As atomic size is periodic, ionic size is also periodic
• In general:
• cations smaller than parent atoms, electrons removed from extended orbital, effective nuclear charge has
increased, the cation is smaller than parent atom
• anions are larger than their parent atoms, electrons have been added to extended orbital, means total
electron–electron repulsion increased, anions are larger than parent atoms
• Ions with the same charge, ionic size increases down a group
• All members of isoelectronic series have same number of electrons
• as nuclear charge increases in isoelectronic series ions become smaller: O2– > F– > Na+ > Mg2+ > Al3+
7.4 Ionization Energy
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Ionization energy of atom or ion is minimum energy required to remove electron from ground state
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1st ionization energy, I1, amount energy required to remove electron from gaseous atom: Na(g)  Na (g) + e
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2+
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2nd ionization energy, I2, is energy required to remove second electron: Na (g)  Na (g) + e
Larger ionization energy, more difficult to remove electron
A sharp increase in ionization energy when a core electron is removed
Variations in Successive Ionization Energies
• Ionization energies for element increase magnitude as successive electrons are removed
• as successive electrons removed, more energy required to pull electron away from increasingly more
positive ion
• A sharp increase in ionization energy occurs when inner-shell electron is removed
Periodic Trends in First Ionization Energies
• Ionization energy generally increases across a period
• across aperiod, Zeff increases, more difficult to remove electron
• 2 exceptions are removing first p electron and removing fourth p electron
2 0
• S electrons more effective at shielding than p electrons, forming s p configuration more favorable
• Second electron placed in p orbital, electron–electron repulsion increases, when electron removed, the s2p3
configuration more stable than starting s2p4 configuration, therefore a decrease in ionization energy
• Ionization energy decreases down a group
• means the outermost electron more readily removed a down a group
• as atom gets bigger, becomes easier to remove electron form most extended orbital
• example: noble gases the ionization energies follow the order:
He > Ne > Ar > Kr > Xe
Electron Configurations of Ions
• Derived from electron configurations of elements with required number of electrons added or removed
1
+
2 5
2 6
• Li: [He]2s becomes Li : [He]
F: [He]2s 2p
becomes F-: [He]2s 2p = [Ar]
• Transition metals tend to lose valence shell electrons first and then d electrons as required to reach the
desired charge on the ion, thus electrons removed from 4s before the 3d, etc
• Writing electron configurations of transition metal cations, the order of removal of electrons is not exactly
opposite to order which subshells were occupied when electron configuration of parent atom written
7.5 Electron Affinities
• Electron affinity is energy change when gaseous atom gains electron to form gaseous ion
• Electron affinity and ionization energy measure energy changes of opposite processes
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• electron affinity: Cl(g) + e  Cl (g)
∆E = –349 kJ/mol
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• ionization energy: Cl(g)  Cl (g ) + e
∆E = 1251 kJ/mol
• Electron affinity can either be exothermic or endothermic:
Ar(g) + e–  Ar– (g)
∆E > 0
• look at electron configurations to determine electron affinity is positive or negative
• extra electron in Ar needs placed in 4s orbital which is higher in energy than the 3p orbital
• added electron in Cl placed in 3p orbital to form stable 3p6 electron configuration
• Electron affinities not change greatly as move down in group
7.6 Metals, Nonmetals and Metalloids
• Metallic character refers to extent which element exhibits physical and chemical properties of metals
• metallic character increases down a group
• metallic character decreases from left to right across a period
Metal
• Metals are shiny and lustrous, malleable and ductile
• Metals are solids at room temperature (exception: mercury is liquid; gallium and cesium melt just above
room temperature) and have very high melting temperatures
• Metals tend to have low ionization energies and form cations easily
• Metals tend to be oxidized when they react
• Compounds of metals with nonmetals tend to be ionic substances
• Metal oxides form basic ionic solids
• most metal oxides are basic: Metal oxide + water  metal hydroxide Na2O(s) + H2O(l)  2NaOH(aq)
• metal oxides react with acidsform salts and water: Metal oxide + acid salt + water
NiO(s) + 2HNO3(aq)  Ni(NO3)2(aq) + H2O(l)
Nonmetals
• Nonmetals are more diverse in their behavior than metals
• Nonmetals are nonlustrous, poor conductors of heat & electricity, lower melting points than metals
• 7 nonmetallic elements exist as diatomic molecules under ordinary conditions:
• H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)
• Nonmetals react with metals, nonmetals tend to gain electrons: Metal + nonmetal  salt
2Al(s) + 3Br2(l)  2AlBr3 (s)
• Compounds composed entirely of nonmetals are molecular substances
• Most nonmetal oxides are acidic:
Nonmetal oxide + water  acid
CO2(g) + H2O(l)  H2CO3(aq)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
• Nonmetal oxides react with bases to form salts and water: Nonmetal oxide + base  salt + water
CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)
Metalloids
• Metalloids have properties are intermediate between those of metals and nonmetals
• example, Si has metallic luster but is brittle
• Metalloids found in semiconductor industry
7.7 Group Trends for the Active Metals
• Alkali metals (group 1A) and the alkaline earth metals (group 2A) called the active metals
Group 1A: The Alkali Metals
• Alkali metals are in Group 1A
• are all soft, chemistry by loss of single s electron: M  M+ + e–
• reactivity increases move down the group
• react with hydrogen to form hydrides
–
• hydrogen present as H , called hydride ion 2M(s) + H2(g)  2MH(s)
• react with water form MOH and hydrogen gas: 2M(s) + 2H2O(l)  2MOH(aq) + H2(g)
• produce different oxides reacting with O2:
4Li(s) + O2(g)  2Li2O(s) (oxide)
2Na(s) + O2 (g)  Na2O2(s) ( peroxide)
K(s) + O2 (g)  KO2(s) (superoxide)
• Emit characteristic colors when placed in flame
• s electron excited by flame and emits energy when returns to ground state
• Na line occurs at 589 nm (yellow), characteristic of 3p  3s transition
• Li line is crimson red
K line is lilac
Group 2A: The Alkaline Earth Metals
• Alkaline earth metals are harder and more dense than the alkali metals
2+
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• Chemistry dominated by loss of two s electrons: M  M + 2e
• Mg(s) + Cl2(g)  MgCl2(s)
2Mg(s) + O2(g)  2MgO(s)
• Reactivity increases down group
• Be does not reactwater
• Mg only react with steam
• Ca and elements below react with water at room temperature: Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
7.8 Group Trends for Selected Nonmetals
Hydrogen
• Hydrogen unique element
• often occurs as colorless diatomic gas, H2
• Reactions of hydrogen and nonmetals very exothermic: 2H2(g) + 2O2(g) 2H2O(l)
• Either gain another electron to form hydride ion, H–, or lose electron to become H+:
• 2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(l)
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• H is called a proton
o
∆H = –571.7 kJ
Group 6A: The Oxygen Group
• Move down the group the metallic character increases.
• O2 is a gas, Te is a metalloid, Po is a metal
• Two of the important forms of oxygen are O2 and ozone, O3
• O2(g) and O3(g) are allotropes
• allotropes are different forms of same element in same state as above
• ozone can be prepared from oxygen:
3O2(g)  2O3(g)
∆Ho = +284.6 kJ
• ozone is pungent and toxic
• oxygen a potent oxidizing agent since O2– ion has noble gas configuration
• two oxidation states for oxygen: –2 (e.g., H2O) and –1 (e.g., H2O2)
• Sulfur is another important member of this group
• most common form of sulfur is yellow S8
• sulfur tends to form S2– in compounds (sulfides)
Group 7A: The Halogens
• Group 7A elements known as halogens ("salt formers")
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• Chemistry of halogens dominated by gaining electron to form anion: X2 + 2e  2X
• Fluorine is one of most reactive substances known: 2F2(g) + 2H2O(l)  4HF(aq) + O2(g) ∆H = –758.9 kJ
• All halogens consist of diatomic molecules, X2
• Chlorine is most industrially useful halogen
• reaction between chorine and water produces hypochlorous acid (HOCl), used to disinfect swimming
pool water: Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq)
• Halogens react with hydrogen to form gaseous hydrogen halide compounds: H2(g) + X2 2HX(g)
• Hydrogen compounds of halogens are all strong acids with the exception of HF
Group 8A: The Noble Gases
• Group 8A elements known as noble gases
• All nonmetals and monoatomic
• Unreactive because have completely filled s and p subshells
• 1962 first compounds of noble gases were prepared: XeF2, XeF4, and XeF6