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Honors unit 3-4 periodic trends. Chap 8. Sizes of Atoms. Radius of a free atom / ion = distance from nucleus to outermost electron shell. The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. Effective Nuclear Charge ( Z eff ).

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sizes of atoms
Sizes of Atoms
  • Radius of a free atom / ion = distance from nucleus to outermost electron shell.
  • The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.
effective nuclear charge z eff
Effective Nuclear Charge (Zeff)
  • In a many-electron atom, electrons are both attracted to the nucleus (p#) and repelled by other electrons (core e#).
    • Shielding effect: Electrons within level and from previous levels block the effects of the (+) nucleus
  • Zeff is the actual attraction experienced by the valence electron
atomic size or ionic size
Atomic Size or Ionic Size

Influenced by three factors:

1. Energy Level (n number)

  • Higher energy level is further away.

2. Effective nuclear charge

Zeff= Z - S

  • Z is the atomic number and S is the number of core electrons.
  • More Zeffcharge pulls electrons in closer.

3. Outermost electron number

  • More outermost e- repel each other.
atomic radius trend
Atomic Radius Trend
  • Radius decreases across a period
    • Increased effective nuclear charge due to decreased shielding
  • Radius increases down a group
    • Each row on the periodic table adds a “shell” or energy level to the atom
  • Radius of transition metals
    • increase in size down the Group
    • atomic radii of transition metals roughly the same size across the same period
example 8 5 choose the larger atom in each pair

N or F, N is further left

  • N or F
  • C or Ge
  • N or Al
  • Al or Ge? opposing trends
  • N or F
  • C or Ge, Ge is further down
  • N or F
  • C or Ge
  • N or Al, Al is further down & left
Example 8.5 – Choose the Larger Atom in Each Pair
sizes of ions
Sizes of Ions
  • Ionic size depends upon:
    • Energy level (n#)
    • Effective nuclear charge Zeff.
    • Outermost e# : repulsion to each other.
sizes of ions1
Sizes of Ions
  • Cations are smaller than their parent atoms.
    • The outermost electron is removed and shell# are less.
sizes of ions2
Sizes of Ions
  • Anions are larger than their parent atoms.
    • Electrons are added and repulsions are increased.
sizes of ions3
Sizes of Ions
  • Ions increase in size as you go down a column.
    • Due to increasing value of n.
sizes of ions4
Sizes of Ions
  • In an isoelectronic series, ions have the same number of electrons.
  • Isoelectronic ionic size decreases with an increasing Zeff.
    • Who are isoelectronic? Li+, O2-, Mg2+
trends in ionic radius
Trends in Ionic Radius
  • Ions in same group have same charge
  • Ion size increases down the group
    • more valence shells (n), larger
  • Cations smaller than neutral atom; Anions bigger than neutral atom
  • Cations smaller than anions in the same period
    • except Rb+1 & Cs+1 bigger or same size as F-1 and O-2
    • The radii of isoelectronic ions decrease as the Z increases.
homework
Homework
  • Atomic radius: Page 359 55, 56, 59, 60, 61
  • Ion radius: Page 359 63 a and b, 64 a, b and c, 68, 69, 70
ionization energy ie
Ionization Energy (IE)
  • Definition: the energy required to remove an electron from an atom, starting from the valence e-.
  • first ionization energy = energy to remove electron from neutral atom
  • Bigger atomic size means less attraction to the e-, therefore less IE needed to remove the e-
  • Tends to increase across a period
    • As radius decreases across a period, the electron you are removing is closer to the nucleus and harder to remove
  • Tends to decrease down a group
    • Outer electrons are farther from the nucleus and easier to remove
example 8 8 choose the atom in each pair with the higher first ionization energy

Al or S, Al is further left

  • Al or S
  • As or Sb, Sb is further down
  • Al or S
  • As or Sb
  • N or Si, Si is further down & left
  • Al or S
  • As or Sb
  • N or Si
  • O or Cl? opposing trends
Example 8.8 – Choose the Atom in Each Pair with the Higher First Ionization Energy

Tro, Chemistry: A Molecular Approach

trends in electron affinity electronegativity
Trends in Electron Affinity/Electronegativity
  • Definition: A measure of the ability of an atom in a chemical compound to attract electrons
    • Electronegativity tends to increase across a period
      • As radius decreases, electrons get closer to the bonding atom’s nucleus
    • Electronegativity tends to decrease down a group
      • As radius increases, electrons are farther from the bonding atom’s nucleus (more shells)
homework1
Homework
  • Page 359 72, 73, 74, 77
metallic character
Metallic Character
  • Metals
    • conduct heat and electricity
    • most oxides basic and ionic
    • form cations in solution
    • lose electrons in reactions – oxidized - + charged
  • Nonmetals
    • electrical and thermal insulators
    • most oxides are acidic and molecular
    • form anions and polyatomic anions
    • gain electrons in reactions – reduced - - charged
  • metallic character increases left
  • metallic character increase down
example 8 9 choose the more metallic element in each pair

Sn or Te

  • P or Sb
  • Ge or In, In is further down & left
  • Sn or Te
  • P or Sb
  • Ge or In
  • S or Br? opposing trends
  • Sn or Te, Sn is further left
  • Sn or Te
  • P or Sb, Sb is further down
Example 8.9 – Choose the More Metallic Element in Each Pair

Tro, Chemistry: A Molecular Approach

trends in the alkali metals
Trends in the Alkali Metals
  • S block, having only 1 valence e-, 1+ charge when forming ion.
  • atomic radius increases down the column  electron affinity/ electron negetivity decreases down the column ionization energy decreases down the column
  • very low ionization energies
    • good reducing agents, easy to oxidize
    • very reactive, not found uncombined in nature
    • react with nonmetals to form salts
  • Replace H in water and form strong base AOH
trends in the halogens
Trends in the Halogens
  • P block, having only 7 valence e-, 1- charge when forming monatomic ion.
  • atomic radius increases down the column  electron affinity decreases down the column ionization energy decreases down the column
  • Very high electron affinities
    • good oxidizing agents, easy to reduce
    • very reactive, not found uncombined in nature
  • reactivity increases down the column
  • react with hydrogen to form HX acids
  • melting point and boiling point increases down the column
trends in the noble gases
Trends in the Noble Gases
  • atomic radius increases down the column
  • Full-shell electron configuration
  • Very high ionization energy. IE decreases down the column
  • very unreactive
    • only found in monatomic form in nature
    • Also called “inert” gas
  • Very low boiling points
    • all gases at room temperature
homework2
Homework
  • Page 360 79, 81, 85, 87