Comprehensive Study Notes on the Periodic Table and Periodicity

Historical Background of the Periodic Table

The periodic table of elements, often referred to as the "Symbol of Chemistry," represents one of the most critical turning points in scientific history. In its modern form, 118 elements are arranged in a tabular format based on their atomic numbers. While elements such as Gold, Silver, Iron, Phosphorus, Sulfur, Zinc, and Arsenic have been known since pre-historic times, the first formal classification attempt occurred in the 18th century when Antoine Lavoisier categorized known elements simply as metals and nonmetals.

In 1829, Johann Wolfgang Döbereiner introduced the concept of triads, which were groups of three elements with similar chemical properties. He observed that the atomic weight of the middle element was approximately the mathematical average of the weights of the other two. A primary example provided is the triad of lithium, sodium, and potassium, represented as $^7\text{Li}$, $^{23}\text{Na}$, and $^{39}\text{K}$. By 1864, English chemist John Newlands observed periodicity among the 62 then-known elements by arranging them in increasing order of atomic mass. His "Law of Octaves" suggested that every eighth element shared properties with the first.

Simultaneously in 1864, Lother Meyer developed curves by plotting atomic weight against atomic volumes, which also demonstrated periodicity. However, Dmitri Mendeleev is recognized as the father of the Periodic Table. In 1869, he arranged 63 elements by increasing atomic mass. Mendeleev’s success lay in his decision to leave gaps for undiscovered elements, accurately predicting their atomic masses and properties. In 1913, Henry Moseley used X-ray emission to determine exact atomic numbers, resolving Mendeleev's discrepancies. This led to the Modern Periodic Law: the properties of elements are periodic functions of their atomic numbers.

Main Features of the Modern Periodic Table

The modern table organizes 118 elements to facilitate a better understanding of their properties. Elements are arranged in ascending order of atomic number. The table consists of seven horizontal rows known as periods and eighteen vertical columns known as groups. In older versions, these groups were categorised into eight A-Groups and ten B-Groups.

Elements within the same group share similar chemical properties due to having an identical number of valence electrons, though they exhibit a gradual change in physical properties from top to bottom. Conversely, elements in a period show a gradual change in chemical and physical properties when moving from left to right. Elements can be further categorized into blocks, families, and categories to enhance comprehension of their behaviors.

Metals, Non-metals, and Metalloids

Elements are broadly divided into three categories based on their electronic tendencies. Metals, such as iron, copper, gold, and silver, tend to lose electrons to form positive ions (cations). Non-metals, such as chlorine, sulfur, and phosphorus, tend to gain electrons to form negative ions (anions). Hydrogen is unique as the first element and is categorized as a non-metal despite its position.

Metalloids (or semimetals) separate the metals and non-metals on the table, exhibiting properties of both. A "stair-step line" typically identifies these, beginning at boron ($\text{B}$) and extending down to polonium ($\text{Po}$), including silicon ($\text{Si}$), germanium ($\text{Ge}$), arsenic ($\text{As}$), antimony ($\text{Sb}$), and tellurium ($\text{Te}$). Elements located to the left of this line are metals, those to the far right are non-metals, and those bordering the line are metalloids.

Blocks in the Periodic Table

Elements are classified into blocks based on the subshell that contains their valence electrons:

1. s-block: Elements in the first two groups (Groups 1 and 2), where valence electrons occupy the "s" subshell.
2. p-block: Elements in Groups 13 to 18 (including noble gases), where valence electrons occupy the "p" subshell.
3. d-block: Transition elements located in the middle of the table, where the "d" subshell is being filled.
4. f-block: The two series at the bottom (Lanthanides and Actinides), characterized by the filling of the "f" subshell.

Identifying an element's block provides critical information regarding its chemical reactivity, oxidation states, electronegativity, and ionization energy.

Element Families

There are five famous families of elements with shared characteristics:

1. Alkali Metals (Group 1): Includes $\text{Li}$, $\text{Na}$, $\text{K}$, $\text{Rb}$, $\text{Cs}$, and $\text{Fr}$. They have one valence electron and are the most reactive metals, forming alkalis when reacting with water.
2. Alkaline Earth Metals (Group 2): Includes $\text{Be}$, $\text{Mg}$, $\text{Ca}$, $\text{Sr}$, $\text{Ba}$, and $\text{Ra}$. They have two valence electrons, are found primarily in the earth, and are less reactive than Group 1.
3. Transition Metals: The largest family, occupying the d-block and f-block. Examples include $\text{Cr}$, $\text{Fe}$, $\text{Co}$, $\text{Ni}$, $\text{Cu}$, and $\text{Zn}$. They are known for variable oxidation states and forming colored compounds.
4. Chalcogens (Group 16): Includes $\text{O}$, $\text{S}$, $\text{Se}$, $\text{Te}$, $\text{Po}$, and $\text{Lv}$. The name comes from the Greek "chalkos" (copper) because copper ores are often oxides or sulfides. In this group, oxygen and sulfur are non-metals, while $\text{Se}$, $\text{Te}$, and $\text{Po}$ are metalloids and livermorium is a metal.
5. Halogens (Group 17): Includes $\text{F}$, $\text{Cl}$, $\text{Br}$, $\text{I}$, $\text{At}$, and $\text{Ts}$. Meaning "salt-former," these are highly reactive non-metals with high electron affinities that easily accept one electron.
6. Noble Gases (Group 18): Includes $\text{He}$, $\text{Ne}$, $\text{Ar}$, $\text{Kr}$, $\text{Xe}$, $\text{Rn}$, and $\text{Og}$. They are almost entirely unreactive due to complete outermost shells. While once called "inert," the term was changed to "noble" after compounds like xenon hexafluoroplatinate ($\text{XePtF}_6$) were discovered.

Periodic Arrangement and Electronic Configuration

The position of an element provides its electronic structure. The period number corresponds to the shell number ($n$). If an element is in Period 3, it has three electron shells. The group number indicates the number of valence electrons. For example, magnesium ($\text{Mg}$) is in Period 3, Group 2, and the s-block. This means its valence electrons are in the $3s$ subshell, specifically identified as $3s^2$.

For an element "X" in Group 13 and Period 3, it has 3 valence electrons across 3 shells. The configuration is constructed as follows:

• 1st shell: $1s^2$
• 2nd shell: $2s^2 2p^6$
• 3rd shell (valence): $3s^2 3p^1$ Total configuration: $1s^2 2s^2 2p^6 3s^2 3p^1$.

Periodicity of Atomic and Ionic Radii

Atomic radius is half the distance between two identical bonded atoms, typically measured in picometers ($pm$) or Angstrom (\text{Å}). Factors affecting it include the number of shells, effective nuclear charge, and the shielding effect of inner electrons. Across a period, atomic radius decreases because increasing nuclear charge pulls the electron cloud closer. Down a group, atomic radius increases as additional electron shells are added, which outweighs the increase in nuclear charge.

Ionic radius measures the size of an ion in a crystal lattice. Cations (positive ions) are smaller than their parent atoms because the loss of an electron shell and increased nuclear pull on remaining electrons causes the cloud to contract. Anions (negative ions) are larger than their neutral atoms because the addition of electrons increases electronic repulsion, causing the cloud to expand. Across a period, ionic radii for both cations and anions generally decrease. Down a group, both increase due to the addition of more electron shells.

Ionization Energy: Factors and Trends

Ionization Energy ($\Delta H_{i1}$) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form 1+ ions. Successive ionization energies exist ($\Delta H_{i2}, \Delta H_{i3}$), as seen in calcium:

• $\text{Ca}_{(g)} \rightarrow \text{Ca}^+_{(g)} + e^- \quad \Delta H_{i1} = 590\,kJ\,mol^{-1}$
• $\text{Ca}^+_{(g)} \rightarrow \text{Ca}^{2+}_{(g)} + e^- \quad \Delta H_{i2} = 1150\,kJ\,mol^{-1}$
• $\text{Ca}^{2+}_{(g)} \rightarrow \text{Ca}^{3+}_{(g)} + e^- \quad \Delta H_{i3} = 4940\,kJ\,mol^{-1}$

Factors affecting Ionization Energy:

1. Nuclear charge: IE increases as effective nuclear charge increases.
2. Atomic size: IE decreases as atomic size increases because the nucleus has a weaker hold on distant electrons.
3. Electronic arrangement: Stable half-filled or completely-filled orbitals require more energy to disrupt. For example, Nitrogen ($1s^2 2s^2 2p^3$) has a higher IE than Oxygen ($1s^2 2s^2 2p^4$) because nitrogen has a more stable half-filled subshell. Noble gases have the highest IE in their periods due to completed shells.
4. Shielding effect: More inner electrons result in lower IE due to reduced nuclear attraction on valence electrons.
5. Spin-pair repulsion: Electrons in the same orbital repel each other. Manganese ($25\text{Mn}: [Ar] 3d^5 4s^2$) has a lower IE for removing one paired $4s$ electron compared to Chromium ($24\text{Cr}: [Ar] 3d^5 4s^1$), where the single electron lacks this repulsion.

Periodic trends show that IE increases across a period (due to nuclear charge) and decreases down a group (due to size and shielding).

Electron Affinity and Enthalpy Changes

First electron affinity ($\Delta H_{ea1}$) is the enthalpy change when adding 1 mole of electrons to 1 mole of gaseous atoms to form uni-negative ions. For chlorine: $\text{Cl}_{(g)} + e^- \rightarrow \text{Cl}^-_{(g)} \quad \Delta H_{ea1} = -348.8\,kJ\,mol^{-1}$ (energy released, negative sign).

Second electron affinity ($\Delta H_{ea2}$) is the energy required to add an electron to a negative ion. For oxygen: $\text{O}_{(g)} + e^- \rightarrow \text{O}^-_{(g)} \quad \Delta H_{ea1} = -142\,kJ\,mol^{-1}$$\text{O}^-_{(g)} + e^- \rightarrow \text{O}^{2-}_{(g)} \quad \Delta H_{ea2} = +844\,kJ\,mol^{-1}$ The positive sign indicates energy is absorbed because the existing negative charge repels the incoming electron.

Factors affecting Electron Affinity include the size of the atom (smaller atoms attract electrons more strongly), nuclear charge (greater charge increases attraction), and configuration (half-filled subshells in $\text{N}$ and $\text{P}$ lead to very low electron affinity). Generally, EA becomes more negative across a period and less negative down a group, although fluorine has a lower EA than chlorine despite its smaller size.

Electronegativity and the Pauling Scale

Electronegativity is the power of an atom in a molecule to attract a shared pair of electrons toward itself. Linus Pauling, the only person to receive two unshared Nobel Prizes (Chemistry 1954 and Peace 1962), developed the standard scale. Alkali metals have minimum values (around 0.8), while fluorine has the maximum (4.0). Electronegativity increases across a period due to increasing nuclear charge and decreasing size, and decreases down a group due to increased shielding and size. Consequently, metals are generally electropositive, while non-metals are electronegative.

Variation in Metallic Character

Metallic character refers to the tendency of an atom to lose electrons. Across a period, metallic character decreases as nuclear charge increases and shells are pulled tighter, making electron loss difficult. Down a group, metallic character increases because the shielding effect and increased size reduce nuclear attraction on valence electrons. Therefore, Cesium is much more reactive and electropositive than Lithium or Sodium.

Reactions of Sodium and Magnesium

With Water: Sodium reacts vigorously with liquid water: $2\text{Na}_{(s)} + 2\text{H}_2\text{O}_{(l)} \rightarrow 2\text{NaOH}_{(aq)} + \text{H}_{2(g)}$ Magnesium reacts slowly with liquid water to form magnesium hydroxide: $\text{Mg}_{(s)} + 2\text{H}_2\text{O}_{(l)} \rightarrow \text{Mg(OH)}_{2(aq)} + \text{H}_{2(g)}$ However, magnesium reacts vigorously with steam to form magnesium oxide: $\text{Mg}_{(s)} + \text{H}_2\text{O}_{(g)} \rightarrow \text{MgO}_{(s)} + \text{H}_{2(g)}$

With Oxygen: Sodium burns with a golden yellow flame. It is stored under kerosene to prevent reaction with air. In open air, it forms peroxide ($\text{Na}_2\text{O}_2$): $2\text{Na}_{(s)} + \text{O}_{2(g)} \rightarrow \text{Na}_2\text{O}_{2(s)}$ Under limited oxygen, it forms the oxide: $4\text{Na}_{(s)} + \text{O}_{2(g)} \rightarrow 2\text{Na}_2\text{O}_{(s)}$ Magnesium burns with an intense white flame (useful in flares/fireworks): $2\text{Mg}_{(s)} + \text{O}_{2(g)} \rightarrow 2\text{MgO}_{(s)}$

With Chlorine: Both metals form soluble salts. Sodium produces a yellow flame and white solid $\text{NaCl}$. Magnesium produces white solid $\text{MgCl}_2$: $2\text{Na}_{(s)} + \text{Cl}_{2(g)} \rightarrow 2\text{NaCl}_{(s)}$$\text{Mg} + \text{Cl}_{2(g)} \rightarrow \text{MgCl}_{2(s)}$

Bonding Trends in Period 3 Oxides and Chlorides

Oxides and chlorides of Group 1, 2, and 3 (e.g., $\text{Na}_2\text{O}$, $\text{NaCl}$) are predominantly ionic with giant ionic lattice structures and high electrostatic forces. Moving to Groups 4 through 7 (e.g., $\text{SO}_2$, $\text{PCl}_5$), the bonding character shifts to covalent with weak intermolecular forces. This transition is driven by the decreasing difference in electronegativity between the atoms.

Classification and Reactions of Oxides

1. Basic Oxides: Metals from Groups 1 and 2 form these. They react with water to form alkalis. Solubility of Group 2 hydroxides increases down the group.

   • $\text{Na}_2\text{O}_{(s)} + \text{H}_2\text{O}_{(l)} \rightarrow 2\text{NaOH}_{(aq)}$
   • $\text{CaO} + \text{H}_2\text{O} \rightarrow \text{Ca(OH)}_{2(aq)}$

2. Acidic Oxides: Non-metals form these via covalent bonds. They react with water to form acids and react with bases to form salts and water.

   • $\text{P}_2\text{O}_{3(s)} + 3\text{H}_2\text{O}_{(l)} \rightarrow 2\text{H}_3\text{PO}_{3(aq)}$
   • $\text{SO}_{2(g)} + \text{H}_2\text{O}_{(l)} \rightarrow \text{H}_2\text{SO}_{3(aq)}$
   • $\text{SiO}_{2(s)} + 2\text{NaOH}_{(aq)} \rightarrow \text{Na}_2\text{SiO}_{3(aq)} + \text{H}_2\text{O}_{(l)}$

3. Amphoteric Oxides: These can react with both acids and bases. Aluminum oxide ($\text{Al}_2\text{O}_3$) is a key example:

   • $\text{Al}_2\text{O}_{3(s)} + 6\text{HCl}_{(aq)} \rightarrow 2\text{AlCl}_{3(s)} + 3\text{H}_2\text{O}$
   • $\text{Al}_2\text{O}_{3(s)} + 2\text{NaOH}_{(aq)} \rightarrow 2\text{NaAlO}_{2(s)} + \text{H}_2\text{O}_{(l)}$

Classification and Reactions of Chlorides

1. Neutral Chlorides: Sodium and magnesium chlorides do not hydrolyze significantly; they undergo hydration to form neutral solutions with pH values near 7 ($pH=7$ for $\text{NaCl}$, $pH=6.5$ for $\text{MgCl}_2$):

   • $\text{NaCl}_{(s)} \rightarrow \text{Na}^+_{(aq)} + \text{Cl}^-_{(aq)}$

2. Acidic Chlorides: From aluminum to sulfur, chlorides react with water through hydrolysis, producing acidic solutions.

   • $\text{AlCl}_{3(s)} + 3\text{H}_2\text{O}_{(l)} \rightarrow \text{Al(OH)}_{3(s)} + 3\text{HCl}_{(aq)}$
   • $\text{SiCl}_{4(l)} + 2\text{H}_2\text{O}_{(l)} \rightarrow \text{SiO}_{2(s)} + 4\text{HCl}_{(aq)}$
   • $\text{PCl}_{3(g)} + 3\text{H}_2\text{O}_{(l)} \rightarrow \text{H}_3\text{PO}_{3(s)} + 3\text{HCl}_{(aq)}$

Oxidation Numbers in Period 3 Elements

The oxidation number (or state) is the formal charge on an atom in a molecule. In Period 3, the highest oxidation state in oxides and chlorides generally corresponds to the number of valence electrons (e.g., $\text{Na}$ is +1, $\text{Mg}$ is +2, $\text{Al}$ is +3, etc.).

Questions & Discussion

Quick Check 1.1:

• Q: Why are the elements in Groups 1 and 2 known as s-block elements?
• A: Their valence electrons are located in the "s" subshell.
• Q: Name the elements in the chalcogen family and give two characteristics.
• A: $\text{O, S, Se, Te, Po, Lv}$. Characteristics include forming copper ores (oxides/sulfides) and having six valence electrons.

Quick Check 1.3:

• Q: Which factors affect atomic and ionic radii?
• A: Number of shells, effective nuclear charge, and the shielding effect.
• Q: Predict the relative sizes of a lithium atom and its ion, $\text{Li}^+$.
• A: The lithium atom is larger than the ion because the ion has lost its outer electron shell.

Quick Check 1.8:

• Q: Why is $\text{AlCl}_3$ an acidic halide, but $\text{NaCl}$ is not?
• A: $\text{AlCl}_3$ undergoes hydrolysis in water, producing $H^+$ ions and $\text{HCl}$, while $\text{NaCl}$ only undergoes hydration, creating a neutral solution.
• Q: Predict if $\text{PCl}_3$ and $\text{NCl}_3$ are acidic or basic.
• A: They are acidic because they are non-metal chlorides that hydrolyze in water to form acids.