Periodic Trends: Periodic PropertiesPeriodic Trends: Periodic Properties

Certain physical and chemical properties recur at regular intervals, and/or vary in regular fashion, when the elements are arranged according to increasing atomic number.

Melting point, boiling point, hardness, density, physical state, and chemical reactivity are periodic properties.

We will examine several periodic properties that are readily explained using electron configurations.

Periodic Trends

• Atomic Size – – Group – increase in size down the group

• (addition of energy levels)

– Period – decrease in size across a period• (increase in the # protons in the nucleus

pulling on the same energy level)

Periodic Properties: Atomic RadiusPeriodic Properties: Atomic Radius

Half the distance between the nuclei of two atoms is the atomic radius.

Covalent radius: half the distance between the nuclei of two identical atoms joined in a molecule.

Metallic radius: half the distance between the nuclei of adjacent atoms in a solid metal.

Periodic Properties: Atomic RadiusPeriodic Properties: Atomic Radius

• Atomic radius increases from top to bottom within a group.

• The value of n increases, moving down the periodic table.

• The value of n relates to the distance of an electron from the nucleus.

Periodic Properties: Atomic RadiusPeriodic Properties: Atomic Radius

• Atomic radius decreases from left to right within a period.

• Why? The effective nuclear charge increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.

Mg has a greater effective nuclear charge than Na, and is smaller than Na.

Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron.Zeff = Zactual - Electron shielding

Multielectron and ShieldingMultielectron and Shielding

Atomic Radii of the ElementsAtomic Radii of the Elements

ExampleExample

With reference only to a periodic table, arrange each set of elements in order of increasing atomic radius:

(a)Na, Al, P

(b) Te, S, Se

(c) Ge, Sb, Sn

Ionic RadiiIonic Radii

The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion.

Ionic Radii of CationsIonic Radii of Cations

• Cations are smaller than the atoms from which they are formed; the value of n usually decreases. Also, there is less electron–electron repulsion.

Ionic Radii of AnionsIonic Radii of Anions

• Anions are larger than the atoms from which they are formed.

• Effective nuclear charge is unchanged, but additional electron(s) increase electron–electron repulsion.

• Isoelectronic species have the same electron configuration; size decreases with effective nuclear charge.

Atomic Radii and the Periodic TableAtomic Radii and the Periodic Table

ExampleExample

Use the standard periodic table to arrange the following species in the expected order of increasing radius:

Sr2+, Ru3+, Rb+, I–, Br–

Isoelectronic atoms= Sr+2, Rb+1, Br-1Atom Protons Electrons Configuration Rank

Sr+2 38 36 [Kr]5s2 [Kr]

1

Ru+3 44 41 [Kr]5s24d6 [Kr]4d5

4

Rb+1 37 36 [Kr]5s1 [Kr]

2

I-1 53 54 [Kr]5s24d105p5 [Xe]

5

Br-1 35 36 [Ar]4s23d104p5 [Kr]

3

Periodic Trend• Ionization Energy –

– Group – decreases down the group• (addition of energy levels makes it harder

for the nucleus to hold on and easier to remove electrons)

– Period – increases across the period• (addition of protons pulling on the

electrons makes it harder to remove the electrons)

Ionization EnergyIonization Energy

• Ionization energy (I) is the energy required to remove an electron from a ground-state gaseous atom.

• I is usually expressed in kJ per mole of atoms.

M(g) M+(g) + e– ΔH = I1

M+(g) M2+(g) + e– ΔH = I2

M2+(g) M3+(g) + e– ΔH = I3

Ionization Energy TrendsIonization Energy Trends

• I1 < I2 < I3– Removing an electron from a positive ion is more difficult than

removing it from a neutral atom.

• A large jump in I occurs after valence electrons are completely removed (why?).

• I1 decreases from top to bottom on the periodic table.– n increases; valence electron is farther from nucleus.

• I1 generally increases from left to right, with exceptions.– Greater effective nuclear charge from left to right holds electrons

more tightly.

Selected Ionization EnergiesSelected Ionization Energies

Compare I2 to I1 for a 2A element, then for the corresponding 1A element.

Why is I2 for each 1A element so much greater than I1?

Why don’t we see the same trend for each 2A element? I2 > I1 … but only about twice as great …

ExampleExample

Why is lithium’s I2 14 times as much as the I1 while beryllium is only twice as much?

Li [He]2s1Li+1 [He] or 1s2Li+2 1s1

Be [He]2s2Be+1 [He]2s1Be+2 [He]

Selected Ionization EnergiesSelected Ionization Energies

General trend in I1: An increase from left to right, but …

…I1 drops, moving from 2A to 3A.

The electron being removed is now a p electron (higher energy, easier to remove than an s).

I1 drops again between 5A and 6A.

Repulsion of the paired electron in 6A makes that electron easier to remove.

ExampleExample

Why is boron's I1 less than that of beryllium?

B

Be

ExampleExample

Why is oxygen’s I1 less than that of nitrogen?

O

N

First Ionization EnergiesFirst Ionization Energies

Second Ionization Energies• The amount of energy required to remove the

second electron once the first is gone.– Always greater than THEIR first IE due to an INCREASED

effective nuclear charge (a greater proton to electron ratio), making the atom smaller.

– However, 2nd Ionization energies can be compared BETWEEN different atoms by evaluating the +1 ion of THAT atom.

• If the +1 ion is a noble gas the 2nd IE will be LARGE. If the +1 ion is NOT a noble gas the 2nd IE will be SMALLER.

Second Ionization Energies

Example: 1st IE 2nd IE

Na

[Ne]3s2

Small (wants to lose 1e-)

Big (Doesn’t want to lose 2e-)

Mg

[Ne]3s2

Bigger than Na (doesn’t want to lose any)

Smaller than Na (once 1e- is gone, it is easy to lose 2e-)

Periodic Trends: Periodic PropertiesPeriodic Trends: Periodic Properties

Certain __________________________recur at regular intervals, and/or vary in regular fashion, when the elements are arranged according to ______________________.

Melting point, boiling point, hardness, density, physical state, and chemical reactivity are periodic properties.

We will examine several periodic properties that are readily explained using electron configurations.

Periodic Trends

• Atomic Size – – Group –

– Period –

Periodic Properties: Atomic RadiusPeriodic Properties: Atomic Radius

Half the distance between the nuclei of two atoms is the _____________________.

_____________________: half the distance between the nuclei of two identical atoms joined in a molecule.

_____________________: half the distance between the nuclei of adjacent atoms in a solid metal.

Periodic Properties: Atomic RadiusPeriodic Properties: Atomic Radius

• Atomic radius __________ from top to bottom within a group.

• The value of n increases, moving down the periodic table.

• The value of n relates to the __________ of an electron from the nucleus.

Periodic Properties: Atomic RadiusPeriodic Properties: Atomic Radius

• Atomic radius __________ from left to right within a period.

• Why? The ________________________increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.

Mg has a greater effective nuclear charge than Na, and is smaller than Na.

Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron.Zeff = Zactual - Electron shielding

Multielectron and ShieldingMultielectron and Shielding

Atomic Radii of the ElementsAtomic Radii of the Elements

ExampleExample

With reference only to a periodic table, arrange each set of elements in order of increasing atomic radius:

(a)Na, Al, P

(b) Te, S, Se

(c) Ge, Sb, Sn

Ionic RadiiIonic Radii

The ___________ of each ion is the portion of the distance between the nuclei occupied by that ion.

Ionic Radii of CationsIonic Radii of Cations

• __________ are _________ than the atoms from which they are formed; the value of n usually decreases. Also, there is less electron–electron repulsion.

Ionic Radii of AnionsIonic Radii of Anions

• ________ are _______than the atoms from which they are formed.

• Effective nuclear charge is unchanged, but additional electron(s) increase electron–electron repulsion.

• _____________ species have the same electron configuration; size decreases with effective nuclear charge.

Atomic Radii and the Periodic TableAtomic Radii and the Periodic Table

ExampleExample

Use the standard periodic table to arrange the following species in the expected order of increasing radius:

Sr2+, Ru3+, Rb+, I–, Br–

Periodic Trend• Ionization Energy –

– Group –

– Period –

Ionization EnergyIonization Energy

• ____________________(I) is the energy required to remove an electron from a ground-state gaseous atom.

• I is usually expressed in kJ per mole of atoms.

M(g) M+(g) + e– ΔH = I1

M+(g) M2+(g) + e– ΔH = I2

M2+(g) M3+(g) + e– ΔH = I3

Ionization Energy TrendsIonization Energy Trends

• I1 < I2 < I3– Removing an electron from a __________ is more difficult than

removing it from a ____________.

• A large jump in I occurs after valence electrons are completely removed (why?).

• I1 ____________ from top to bottom on the periodic table.– n increases; valence electron is farther from nucleus.

• I1 generally _________ from left to right, with exceptions.– Greater effective nuclear charge from left to right holds electrons

more tightly.

Selected Ionization EnergiesSelected Ionization Energies

Compare I2 to I1 for a 2A element, then for the corresponding 1A element.

Why is I2 for each 1A element so much greater than I1?

Why don’t we see the same trend for each 2A element? I2 > I1 … but only about twice as great …

ExampleExample

Why is lithium’s I2 14 times as much as the I1 while beryllium is only twice as much?

Selected Ionization EnergiesSelected Ionization Energies

General trend in I1: An increase from left to right, but …

…I1 drops, moving from 2A to 3A.

The electron being removed is now a p electron (higher energy, easier to remove than an s).

I1 drops again between 5A and 6A.

Repulsion of the paired electron in 6A makes that electron easier to remove.

ExampleExample

Why is boron's I1 less than that of beryllium?

ExampleExample

Why is oxygen’s I1 less than that of nitrogen?

First Ionization EnergiesFirst Ionization Energies

Second Ionization Energies• The amount of energy required to remove the

second electron once the first is gone.– Always greater than _____ first IE due to an

____________ effective nuclear charge (a greater proton to electron ratio), making the atom smaller.

– However, 2nd Ionization energies can be compared ________ different atoms by evaluating the +1 ion of ____ atom.

• If the +1 ion is a noble gas the 2nd IE will be _____. If the +1 ion is NOT a noble gas the 2nd IE will be _________.

Second Ionization Energies

Example: 1st IE 2nd IE

Na Small (wants to lose 1e-)

Big (Doesn’t want to lose 2e-)

Mg Bigger than Na (doesn’t want to lose any)

Smaller than Na (once 1e- is gone, it is easy to lose 2e-)