Periodic Properties of the Elements

1

• Periodic Properties of the Elements
• For representative elements the similarities in
  chemical and physical properties can often be
  explained by the similarities in electron
  configurations.
• Elements in the same group have the same valence
  shell electron configuration except for the
  n-value
• Be 1s22s2 Mg 1s22s22p63s2 Ca
  1s22s22p63s23p64s2
• Some elements in the same group may have somewhat
  different properties because of the differences
  in their overall electron configuration
• N2 is a gas at room temperature whereas P exists
  in a couple of solid forms
• Origin of the Periodic Table starts with Dmitri
  Mendeleev and Lothar Meyer
• Both arranged the elements in order of atomic
  weight and placed elements with similar
  properties in columns.
• At the time only 63 or so elements were known and
  there were obvious missing elements in their
  periodic tables
• Mendeleev was able to predict fairly accurately
  the chemical and physical properties of some of
  these missing elements such as Ga and Ge
• Henry Moseley observed that the n of x-rays
  emitted by elements bombarded with electrons
  increased with atomic weight.
• He labeled the emissions using an integer he
  called the atomic number.

2

• Periodic Properties of the Elements
• The use of the atomic number improved the
  organization
• For some horizontal neighboring pairs, the
  element with the higher atomic number has a
  lower atomic weight.
• Compare Ar to K Co to Ni Te to I Th to Pa U
  to Np
• Atomic Size and Electron Shells
• As z increases, the maximum probability of
  finding the electrons in a shell with a
  particular n-value occurs at smaller distance
  from the nucleus.
• See Figure 7.3 in Brown, LeMay and Bursten
• This trend is associated with the increased
  Zeffective as n increases.
• Atoms do not have a fixed size because the
  probability of finding the electron in the
  valence shell decreases with distance from the
  nucleus but never gets to zero.
• One measure of the size of an atom is its
  covalent radius measured from interatomic
  distances of atoms in molecules.
• The C-C distance in diamond and other compounds
  is 1.54 Å, rC0.77 Å
• The Cl-Cl distance in Cl2 is 1.99 Å, rCl0.99 Å
• In CCl4, rCCl1.77 Å which is close to 0.77 Å
  0.99 Å1.76 Å

3

• Periodic Properties of the Elements
• Other Kinds of radii
• Ionic radii of monatomic ions
• Cations are smaller than neutral atoms
• Anions are larger than neutral atoms
• Metal atom radii measured from atom separations
  in metal crystals
• Periodic trends in atomic size
• In a vertical group, the lower the element the
  larger the atom
• In a horizontal period, moving to the right the
  smaller the atom
• These results are consistent with
• an increase in the radius of the valence shell
  with n
• Zeffective is fairly constant as we move
  down a group
• an increase in Zeffective with atomic number for
  valence shell electrons having the same n, i.
  e., for atoms in the same period
• As we move right in a period, the nuclear
  charge increases but the number of inner
  shell electrons does not change

4

• Periodic Properties of the Elements
• Ionization Energy the energy required to remove
  an electron from a gaseous atom or ion
• X(g) X(g) e- DE is the 1st ionization
  energy for X
• X-(g) X2(g) e- DE is the
  2nd ionization energy for X
• all ionization energies are endothermic the
  larger the ionization energy the harder it is to
  remove an electron from an atom or ion
• Table 7.2 shows the ionization energies for
  period 3 elements, Na to Ar
• Successive ionizations of the same element are
  more endothermic
• The nuclear charge is constant but interelectron
  repulsions decrease as electrons are removed
  from a valence shell
• There is a big endothermic jump in the ionization
  energy when the next after the last valence
  electron is removed
• The core electrons have experienced a much larger
  Zeffective than valence electrons
• These results support the idea that the valence
  electrons are important in forming molecular
  compounds because they can be shared with other
  atoms or transferred in chemical reactions

5

• Periodic Properties of the Elements
• Periodic trends in ionization energies
• Figure 7.6, page 231, plots the 1st ionization
  energies vs. atomic number for the first 54
  elements, H - Xe
• With some notable exceptions, the 1st ionization
  energy becomes more endothermic as we move to
  the right in any period. The alkali metals have
  the lowest values and the noble gases the
  highest.
• Within a group the 1st ionization energy
  decreases as we go down.
• Representative elements show large variation in
  the 1st ionization energy whereas the transition
  elements and f-block elements show little
  variation.
• These trends are explained in terms of Zeffective
  and electron nuclear distance
• If Zeffective increases or the electron-nucleus
  distance decreases, the ionization energy
  increases
• Moving to the right increases Zeffective and
  decreases electron-nucleus distance
  ionization energy increases
• Moving down a group changes Zeffective
  little but electron-nucleus distance increases
  ionization energy decreases

6

• Periodic Properties of the Elements
• Periodic trends in ionization energies
• Irregularities in going from group 2A to 3A
• For Be, 1s22s2, a 2s electron is lost and for B,
  1s22s22p1, a 2p electron is lost
• Zeffective(Be) is about 2 for the 2s electrons
  because the 2s electrons are not very effective
  in screening the nucleus from themselves.
• Zeffective(B) is a bit less than 2 for the 2p
  electron because the 2s electrons screen the
  nuclear charge from the 2p electrons
• For N, 1s22s22p3, an electron from a half-filled
  p orbital is lost, and for
• O, 1s22s22p4, an electron from a filled p
  orbital is lost.
• There is electron-electron repulsion between the
  electrons in the same p orbital which reduces
  the ionization energy for O.
• The Electron Affinity the energy change
  accompanying the gain of an electron by a
  gaseous atom
• X(g) 1e- X-(g) DE is the electron
  affinity for element X
• Most electron affinities are exothermic some are
  endothermic such as Be, Mg, N and the noble
  gases.

7

• Periodic Properties of the Elements
• Periodic trends in electron affinities
• Generally going to the right in a row or from the
  bottom to the top, the electron affinities
  become more exothermic.
• Zeffective for the added electron increases
  left-to-right or bottom-up
• The electron-nucleus distance decreases
  left-to-right and bottom-up
• These factors increase the attraction of a
  neutral atom for an electron
• Irregularities in these trends
• group 2A elements have less exothermic electron
  affinities than their previous 1A neighbor
• The added electron for 2A a element enters a
  higher energy p orbital and there is a lower
  electron-nucleus attraction than when a 1A
  element adds an electron to a half-filled s
  orbital
• Be(1s22s1) e Be-(1s22s2) Li(1s22s1)
  e Li-(1s22s2)
• group 5A elements have less exothermic electron
  affinities than their previous 4A neighbor
• The added electron for the 5A element enters a
  half-filled p orbital with greater
  electron-electron repulsion than when a 4A
  element adds an electron to an unfilled p
  orbital
• N(1s22s22p3) e N-(1s22s22p4)
  C(1s22s22p2) e C-(1s22s22p3)

8

• Periodic Properties of the Elements
• Irregularities in these trends
• Notice the 3A through 7A elements in the 2nd
  period show less exothermic electron affinities
  than their neighbors immediately below.
• This is ascribed to the extra small size of
  period 2 elements and increased
  electron-electron repulsions in the ion compared
  to period 3 elements
• Properties of Metals

Metallic luster
Good conductors of heat and electricity
As Solids malleable and ductile
Metal oxides are ionic solids that are
basic CaO(s) H2O(l ) Ca(OH)2(aq) CaO(s)
2HCl(aq) CaCl2(aq) H2O(l )
Only 2 are colored
Form cations in water
Low ionization energies
Many are easily oxidized by H and O2
Wide range of melting temperatures
9
Periodic Properties of the Elements Properties of
Non-metals
Do not have metallic luster
Tend to form molecular compounds with other
non-metals
Various colors some colorless
Solids may be brittle some hard, some soft
Tend to form anions or oxyanions in water
Poor conductors of heat and electricity
Most non-metal oxides form acids in
water SO3(g) H2O(l ) H2SO4(aq)
Melting temperatures tend to be lower than metals
Metals react with nonmetals to give ionic
compounds in which the nonmetal is anionic.
Properties of Metalloids the properties are
intermediate
Some - Si - are semiconductors
Some look metallic - Si- but are brittle
10

• Periodic Properties of the Elements
• Alkali metals
• Table 7.4, p 240, in Brown, LeMay and Bursten
  gives some properties of the alkali metals
  showing their low densities, low melting
  temperatures and low ionization energies
• They are the most active metals and are always
  found combined in nature
• They are isolated by electrolysis 2NaCl(l )
  2Na(s) Cl2(g)
• They react with most non-metals
• 2M(s) H2(g) 2MH(s)
• 2M(s) S(s) M2S(s)
• 2M(s) Cl2 (g) 2MCl(s)
• They react with H2O(l ) with sufficient release
  of heat to cause ignition of the H2 produced
• 2M (s) 2H2O(l ) H2(g) 2MOH(aq)
• Li reacts with O2(g) to form lithium oxide
  4Li(s) O2 (g) 2Li2O(s)
• Na reacts with O2(g) to form sodium peroxide
  2Na(s) O2 (g) Na2O2(s)
• K, Rb, Cs react with O2(g) to give metal
  superoxide M(s) O2(g) MO2(s)

Electricity
11

• Periodic Properties of the Elements
• Alkaline Earth Metals
• Table 7.5, page 242, in Brown, LeMay and Bursten
  gives some properties of the alkaline earth
  metals showing density, ionization energies
  and melting temperatures.
• Densities are low but higher than alkali metals
• Ionization energies are low, but higher than that
  of alkali metals
• Melting temperatures are much higher than alkali
  metals
• Their chemical activity increases with atomic
  mass
• Be does not react with water or steam even at red
  heat
• Mg does not react with water but will react with
  steam
• Mg(s) H2O(g) MgO(s) H2(g)
• Ca, Sr, Ba react with water at room temperature
• Ca(s) 2H2O(l ) Ca(OH)2(aq) H2(g)

12

• Periodic Properties of the Elements
• Group trends for selected non-metals
• Hydrogen often put in group 1A and 7A. It is not
  a metal.
• Forms compounds that form H ion in water
• Forms hydrides with active metals where H- is the
  form of hydrogen
• 2Na H2 2NaH
• Chalcogens Group 6A
• Some properties given in Table 7.6, p 244, Brown,
  LeMay and Bursten
• O, S, Se are nonmetals, Po is a radioactive metal
• O2(g) and O3(g) are two allotropes of oxygen
• Allotropes are different forms of an element in
  the same state
• O3(g) is a pungent smelling, poisonous gas, that
  can be formed in an electric discharge 3O2(g)
  O2(g) DHo 284.6 kJ
• O2 is a very good oxidizing agent(! )because it
  attracts electrons strongly from other metal
  and non-metal elements.
• Compounds of O22- and O2- slowly decompose
• 2H2O2(aq) 2H2O(l ) O2(g) DHo -196.1
  kJ

electricity
13

• Periodic Properties of the Elements
• Chalcogens Group 6A
• Sulfur exists as S8 molecules in its most stable
  form.
• S readily forms S2- ions because it attracts
  electrons from other elements
• S also reacts with other elements including
  oxygen S(s) O2(g) SO2(g)
• Halogens - salt formers Group 7A
• Some properties given in Table 7.7, p 246, Brown,
  LeMay and Bursten
• All halogens are non-metal, diatomic molecules
• The halogens have very exothermic electron
  affinities and very endothermic ionization
  energies, so much of the chemistry involves X-
  ions.
• F2 is the best oxidizing agents of all the
  elements and reacts vigorously with almost every
  other substance
• 2H2O(l ) 2F2(g) 4HF(aq) O2(g) DHo
  -758.7 kJ
• Cl2 can be produced from electrolysis of molten
  NaCl or brine
• 2NaCl(aq) 2H2O(l ) 2NaOH(aq) Cl2(g)
  H2(g)

14

• Periodic Properties of the Elements
• Halogens
• Cl2 is added to water (where its converted to
  HOCl) as a disinfectant
• Cl2(g) H2O(l ) HOCl(aq) HCl(aq)
• Noble gases group 8A
• Some properties given in Table 7.8, p 247, Brown,
  LeMay and Bursten
• Exist as monatomic gases
• Have very exothermic ionization energies and
  endothermic electron affinities
• These elements are not chemically very active.
• Only a few compounds with F or F and O are known
  for Xe and KrF2