Transcript Chapter 6-Periodic Table & Periodic Law

Section 6.3 - Periodic Trends
• Objective:
- Compare period & group trends for shielding, atomic radius,
ionic radius, ionization energy, & electronegativity.
Shielding (or screening):
• The valence e- are blocked from the full positive charge
of the nucleus (effective nuclear charge) by the inner
(core) e-.
• As the average number of core e- increases, the
effective nuclear charge decreases.
– Concept of shielding will play a large role in a lot of the trends.
Shielding:
• Trend within the period (left to right): Generally decreases
• Why:
– Number of energy levels & core e- stays the same, but nucleus is
increasing.
– Increased attraction between nucleus and valence e-.
• Trend down a group: Generally increases
• Why:
– Number of energy levels & core e- increases.
– Valence e- farther from nucleus & more blocked by inner e-.
Atomic Radius:
• ½ the distance between adjacent nuclei of identical atoms.
• Trend within the period (left to right): Generally decreases
• Why:
– Number of energy levels & core e- stays the same, but nucleus is increasing.
– Increased attraction between nucleus and valence e-.
– This attraction pulls the e- closer to the nucleus and makes the atom smaller.
• Trend down a group: Increases
• Why:
– Number of energy levels increases & core eincreases.
– Each energy level is larger than the next.
– Valence e- farther from the nucleus and
more blocked by the inner e-.
Examples – Place each group of elements in order of increasing
atomic radius:
1. S, Al, Cl, Mg, Ar, Na
2. K, Li, Cs, Na, H
3. Ca, As, F, Rb, O, K, S, Ga
Examples – Place each group of elements in order of increasing atomic
radius:
1. S, Al, Cl, Mg, Ar, Na
Ar < Cl < S < Al < Mg < Na
2. K, Li, Cs, Na, H
H < Li < Na < K < Cs
3. Ca, F, As, Rb, O, K, S, Ga
F < O < S < As < Ga < Ca < K < Rb
• Ionic Radius: – Distance between the nucleus and the
outermost electron in ions (can’t be determined directly).
• Trend between atom & ion:
– Cations are smaller than original atom. (Losing e-, the atom has
unequal positive charge that attracts the valence e- closer to the nucleus.)
– Anions are larger than original atom and cations. (Adding
negative e-, adds to repulsion between valence e-, pushing them apart.)
• Trend within the period (left to right): Representative
Elements → Decreases.
– Cations: size decreases.
– Anions: the size drastically increases compared to the positive
ions, and then decreases across the period.
• Trend down a group: Increases for both cations & anions.
– Same reason as atomic radii trend.
Examples – Choose the larger species in each case:
1. Na or Na+
2. Br or Br3. N or N34. O- or O25. Mg2+ or Sr2+
6. Mg2+ or O27. Fe2+ or Fe3+
• Ionization Energy:
Energy required to remove an electron from a gaseous atom
(also called First Ionization Energy, I1)
Na (g) + 496 kJ  Na+ (g) + eThe second ionization energy, I2, is the energy required to remove the
next available electron:
Na+ (g) + 4562 kJ  Na2+ (g) + e• NOTICE:
– Ionization Energy increases for each electron removed from the same
element.
– The larger ionization energy, the more difficult it is to remove the electron.
Variations in Successive Ionization Energies
• There is a sharp increase in ionization energy when a
core electron is removed.
• Notice the large increase after the last valence electron is
removed. This chart can be used to determine the number of
valence electrons in an atom of an element.
• Trend within the period: Increases
• Why:
– Electrons are more difficult to remove from smaller atoms.
– Closer to the nucleus and increased nuclear charge.
• Trend down a group: Decreases
• Why:
– Electrons are easier to remove from large atoms.
– Farther away from the nucleus so less energy is needed to
remove them.
• Notice the trend in ionization energy is inversely
related to trends in atomic radii.
Examples – Put each set in order of increasing first ionization energy:
1. P, Cl, Al, Na, S, Mg
2. Ca, Be, Ba, Mg, Sr
3. Ca, F, As, Rb, O, K, S, Ga
Examples – Put each set in order of increasing first ionization energy:
1. P, Cl, Al, Na, S, Mg
2. Ca, Be, Ba, Mg, Sr
3. Ca, F, As, Rb, O, K, S, Ga
1. Na < Al < Mg < S < P < Cl
2. Ba < Sr < Ca < Mg < Be
3. Rb < K < Ca < Ga < As < S < O < F
ELECTRONEGATIVITY:
• Ability of an atom to attract electrons in a chemical bond to itself.
• Chemist Linus Pauling set electronegativities on a scale.
– 0.7 (Cs) to 4.0 (F)
– Used to help determine types of bonding (ionic or covalent) that are
occurring in a compound.
– Noble gases are not usually given electronegativity values.
• Trend within the period: Increases
• Why:
– Atoms become smaller, so shared electrons are closer to the nucleus.
• Trend down a group: Decreases
• Why:
– Atoms become larger, so shared electrons are farther from the nucleus.
Electronegativity
Examples – put each set in order by increasing electronegativity:
1. Na, Li, Rb, K, Fr
2. Cl, Ca, F, P, Mg, S, K
Examples – put each set in order by increasing electronegativity:
1. Na, Li, Rb, K, Fr
2. Cl, Ca, F, P, Mg, S, K
1. Fr < Rb < K < Na < Li
2. K < Ca < Mg < P < S < Cl < F
Review:
1. As you move across a period, left to right, describe
what generally happens (decreases, increases, or
remains the same) to:
a.
Number of valence electrons
b.
Ionization energy
c.
Atomic radius
2. Give a brief explanation for your answers to a-c.
3. Identify the element from the clues given:
a. This element has a smaller atomic radius than phosphorous, it
has a smaller ionization energy than fluorine, and is
chemically similar to iodine.
b. This element has the smallest ionization energy of any
element in Period 4.