Periodic Table and Periodicity

This lesson covers all essential concepts to help you understand how elements are organized and the patterns they follow.

1. Historical Development of the Periodic Table

The Periodic Table is one of the most important organizational tools in chemistry. Let's explore its development:

1.1 Early Classification Attempts

Johann Döbereiner (1829): Proposed the Law of Triads - groups of three elements with similar properties had the middle element with properties that were approximately the average of the other two.
John Newlands (1864): Proposed the Law of Octaves - elements arranged by increasing atomic mass showed a repetition of properties at every eighth element.

1.2 Mendeleev's Periodic Table

Dmitri Mendeleev (1869): Created the first widely recognized periodic table by arranging elements by increasing atomic mass and similarity of chemical properties.
Key contributions:
- Left gaps for undiscovered elements and predicted their properties
- Sometimes placed elements out of atomic mass order to match chemical properties

1.3 Modern Periodic Table

Henry Moseley (1913): Discovered that elements should be arranged by atomic number (number of protons), not atomic mass.
This resolved irregularities in Mendeleev's table and established the Modern Periodic Table we use today.

2. Structure of the Modern Periodic Table

2.1 Organization Principles

Elements are arranged in order of increasing atomic number (number of protons).
Elements with similar properties are placed in the same column (group).
New rows (periods) begin when electron shells start filling.

2.2 Groups and Periods

Groups: Vertical columns labeled 1-18
- Elements in the same group have similar chemical properties
- This is due to the same number of electrons in their outer shell
Periods: Horizontal rows labeled 1-7
- Elements in the same period have the same number of electron shells
- Properties change progressively across a period

2.3 Blocks in the Periodic Table

s-block: Groups 1-2 (electron filling s-orbitals)
p-block: Groups 13-18 (electron filling p-orbitals)
d-block: Groups 3-12 (electron filling d-orbitals)
f-block: Lanthanides and Actinides (electron filling f-orbitals)

2.4 Classification of Elements

Metals: Left side of the periodic table
- Alkali metals (Group 1)
- Alkaline earth metals (Group 2)
- Transition metals (Groups 3-12)
- Post-transition metals (parts of Groups 13-16)
- Lanthanides and Actinides (f-block)
Non-metals: Right side of the periodic table
- Halogens (Group 17)
- Noble gases (Group 18)
- Other non-metals (parts of Groups 14-16)
Metalloids: Elements along the staircase line between metals and non-metals
- Display properties of both metals and non-metals
- Examples: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium

3. Periodic Trends

Elements display regular patterns in their properties as you move across periods or down groups in the Periodic Table. These are called periodic trends.

3.1 Atomic Radius

Across a period (left to right): Atomic radius generally decreases
- Reason: Increasing nuclear charge pulls electrons closer to the nucleus
Down a group: Atomic radius increases
- Reason: Additional electron shells increase the distance from nucleus to outer electrons

3.2 Ionization Energy

Definition: Energy required to remove an electron from a gaseous atom or ion.
Across a period (left to right): Ionization energy generally increases
- Reason: Higher nuclear charge and smaller atomic radius make it harder to remove electrons
Down a group: Ionization energy decreases
- Reason: Outer electrons are farther from the nucleus and experience more shielding

3.3 Electron Affinity

Definition: Energy change when a gaseous atom gains an electron.
Across a period (left to right): Electron affinity generally becomes more negative (more favorable)
- Reason: Atoms closer to completing their octet are more willing to accept electrons
Down a group: Electron affinity generally becomes less negative (less favorable)
- Reason: Larger atoms have less attraction for additional electrons

3.4 Electronegativity

Definition: Ability of an atom to attract shared electrons in a covalent bond.
Across a period (left to right): Electronegativity increases
- Reason: Stronger pull from nucleus due to higher nuclear charge and smaller atomic radius
Down a group: Electronegativity decreases
- Reason: Outer electrons are farther from the nucleus

3.5 Metallic Character

Definition: Tendency of an element to lose electrons and form positive ions.
Across a period (left to right): Metallic character decreases
- Reason: Increasing tendency to gain rather than lose electrons
Down a group: Metallic character increases
- Reason: Outer electrons are less tightly held and more easily lost

Note: There can be some irregularities in these trends due to electron configuration effects, particularly when moving from the s-block to p-block elements.

4. Specific Groups in the Periodic Table

4.1 Group 1: Alkali Metals

Elements: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), Francium (Fr)
Properties:
- Soft, silvery, highly reactive metals
- Low melting and boiling points compared to other metals
- Low density (Li, Na, and K float on water)
- Readily lose one electron to form +1 ions
- React vigorously with water to produce hydrogen gas and metal hydroxides
- Stored under oil to prevent reaction with air/moisture
Trends down the group:
- Increasing atomic radius
- Decreasing ionization energy
- Increasing reactivity

4.2 Group 2: Alkaline Earth Metals

Elements: Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), Radium (Ra)
Properties:
- Harder, denser, and stronger than Group 1 metals
- Less reactive than Group 1 but still reactive
- Higher melting and boiling points than Group 1
- Readily lose two electrons to form +2 ions
- React with water (less vigorously than Group 1) to form hydroxides and hydrogen gas
- Form basic oxides that react with water to form alkaline solutions
Trends down the group:
- Increasing atomic radius
- Decreasing ionization energy
- Increasing reactivity

4.3 Group 17: Halogens

Elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At), Tennessine (Ts)
Properties:
- Non-metals with diatomic molecules (e.g., F₂, Cl₂)
- Highly reactive, especially with metals to form salts
- Low melting and boiling points
- Strong oxidizing agents (gain electrons easily to form -1 ions)
Trends down the group:
- Increasing atomic radius
- Decreasing electronegativity
- Decreasing reactivity

4.4 Group 18: Noble Gases

Elements: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn), Oganesson (Og)
Properties:
- Colorless, odorless, monatomic gases
- Very low chemical reactivity (full outer electron shell)
- Low boiling points
- Used in lighting (e.g., neon signs) and as inert environments (e.g., argon in welding)
Trends down the group:
- Increasing atomic radius
- Increasing boiling point
- Slight increase in reactivity (only for heavier noble gases like xenon)

Self-Assessment Questions

1. What is metallic character?
It is the tendency of an element to lose electrons and form positive ions.
2. How does metallic character change across a period?
It decreases from left to right across a period.
3. How does metallic character change down a group?
It increases down a group as atomic size increases.
4. Why do alkali metals have high metallic character?
They have low ionization energy and readily lose their outer electron.
5. What is a common property of alkali metals?
They are soft, silvery metals that react vigorously with water.
6. How do alkaline earth metals differ from alkali metals?
They are harder, less reactive, and have higher melting points.
7. What happens to the reactivity of alkali metals down the group?
Reactivity increases down the group.
8. What ion do Group 2 elements typically form?
They form +2 ions by losing two outer electrons.
9. What is a key chemical property of halogens?
They are highly reactive nonmetals and strong oxidizing agents.
10. Why are noble gases unreactive?
They have full outer electron shells, making them stable and inert.
11. What trend is observed in ionization energy across a period?
Ionization energy increases across a period.
12. Which group contains elements that form salts with metals?
Group 17 (Halogens) form salts with metals.