Section 6.3 - Periodic TrendsSection 6.3 - Periodic Trends• Objective:

- Compare period & group trends for shielding, atomic radius, ionic radius, ionization energy, & electronegativity.

Shielding (or screening):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.

• 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-.

Shielding: Shielding:

Atomic Radius: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 e-

increases.

– 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 Br-

3. N or N3-

4. O- or O2-

5. Mg2+ or Sr2+

6. Mg2+ or O2-

7. 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) + e-

The 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: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: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.