Periodic Trends and Properties in the Elements

Periodic Properties

Introduction to Periodic Properties

When elements are arranged in order of increasing atomic mass (or atomic number), their chemical and physical properties recur periodically. This periodicity forms the basis of the modern periodic table. Dmitri Mendeleev used these recurring patterns to predict the properties of elements that had not yet been discovered.

• Periodic Law: Properties of elements are periodic functions of their atomic numbers.

• Prediction of Properties: Mendeleev's predictions for undiscovered elements were later confirmed by experimental data.

Element	Mendeleev's Predicted Properties	Actual Properties
Gallium (eka-aluminum)	Atomic mass: about 68 amu Density: 5.9 g/cm3 Formula of oxide: E2O3 Formula of chloride: ECl3	Atomic mass: 69.7 amu Density: 5.91 g/cm3 Formula of oxide: Ga2O3 Formula of chloride: GaCl3
Germanium (eka-silicon)	Atomic mass: about 72 amu Density: 5.5 g/cm3 Formula of oxide: EO2 Formula of chloride: ECl4	Atomic mass: 72.64 amu Density: 5.35 g/cm3 Formula of oxide: GeO2 Formula of chloride: GeCl4

Effective Nuclear Charge (Z*)

Definition and Calculation

Periodic properties are largely determined by the effective nuclear charge (Z*), which measures how strongly valence electrons feel the attraction of the nucleus. It is a balance between the attractive force of the nucleus and the repulsion by other electrons (shielding).

• Formula: $Z^* = Z - \text{shielding}$

• Shielding: Electrons in lower shells shield more effectively than those in lower subshells; electrons in the same subshell do not shield each other.

Example: For lithium (Li), the outermost 2s electron is shielded by the two 1s electrons. $Z^* \approx 3 - 2 = 1$.

Trends in Z*

• As you go across a period, Z increases and shielding increases only slightly, so $Z^*$ increases.

• As you go down a group, additional shells increase shielding significantly, so $Z^*$ increases less rapidly.

Atom	Z*
Li	+1.28
Be	---
B	+2.58
C	+3.22
N	+3.85
O	+4.49
F	+5.13

Stability: Full shells or subshells are particularly stable due to high effective nuclear charge.

Atomic and Ionic Radii

Atomic Radii

The atomic radius is defined as half the bond distance in a diatomic molecule. Atomic radii can be estimated using bond distances.

• Example: $r_{\text{Cl}} = \frac{198\,\text{pm}}{2} = 99\,\text{pm}$

• Trend Down a Group: Atomic radii increase due to the addition of electron shells.

• Trend Across a Period: Atomic radii decrease as $Z^*$ increases, pulling electrons closer to the nucleus.

Transition Metals: Atomic sizes stay about the same across the transition metals because electrons are added to inner shells, so both Z and shielding increase, keeping $Z^*$ roughly constant.

Ionic Radii

• Cations are smaller than their parent atoms because the loss of electrons increases the proton/electron attraction.

• Anions are larger than their parent atoms because the gain of electrons decreases the proton/electron attraction.

Example: Li (152 pm) loses an electron to become Li+ (78 pm); F (71 pm) gains an electron to become F- (133 pm).

Ionization Energy (IE)

Definition and Trends

Ionization energy is the energy required to remove an electron from an atom in the gas phase.

• Equation: $\text{Mg (g)} + 738\,\text{kJ} \rightarrow \text{Mg}^+ (g) + e^-$

• Successive ionization energies increase because the electron/proton ratio decreases, making it harder to remove additional electrons.

• Large jumps in IE occur when removing electrons from a full shell.

Element	IE1	IE2	IE3	IE4
Na	496	4560	---	---
Mg	738	1450	7730	---
Al	578	1816	2740	11600
Si	786	1577	3230	4360
P	1012	1900	2910	4960
S	1000	2250	3350	4550
Cl	1251	2290	3820	5150
Ar	1520	2660	3930	5770

• Trend Across a Period: IE increases as $Z^*$ increases.

• Trend Down a Group: IE decreases as atomic size increases.

• Exceptions: N vs. O, P vs. S, Be vs. B, Mg vs. Al due to subshell stability and electron configurations.

Electron Affinity (EA)

Definition and Trends

Electron affinity is the energy released when a neutral atom accepts an electron. It is usually negative, indicating energy is released.

• Equation: $\text{Cl} (g) + e^- (g) \rightarrow \text{Cl}^- (g)$, $\Delta E = -349\,\text{kJ/mol}$

• EA becomes more negative across a period (increased $Z^*$), but less negative down a group (increased shielding).

• Noble gases and Group II metals have EA ≈ 0 because added electrons would enter a new shell or subshell, experiencing complete shielding.

• Nitrogen also has EA ≈ 0 due to a half-filled subshell.

Electronegativity (χ)

Definition and Trends

Electronegativity (χ) is a measure of the ability of an atom in a bond to attract electrons to itself. Values range from 0.8 to 4.0 (Pauling scale).

• Trends are similar to those of ionization energy and electron affinity.

• Electronegativity increases across a period and decreases down a group.

Element	Electronegativity (χ)
F	4.0
O	3.5
N	3.0
Cl	3.0
H	2.1

Bond Polarity

Differences in electronegativity between atoms in a bond result in bond polarity. The greater the difference, the more polar the bond.

• Example: O-H bond ($\Delta \chi = 1.4$) is more polar than O-F bond ($\Delta \chi = 0.5$).

• The arrow in bond diagrams points toward the more electronegative (negative) end.

Summary Table: Periodic Trends

Property	Across a Period	Down a Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	More negative	Less negative
Electronegativity	Increases	Decreases

Additional info: These notes cover the core periodic trends and their explanations, suitable for General Chemistry students preparing for exams or seeking a concise review of periodic properties.