Classification of Elements and Periodicity in PropertiesClass 11 Chemistry Notes

WHY DO WE NEED TO CLASSIFY ELEMENTS?

Elements are the basic building blocks of all matter. As more elements were discovered over time (from 31 in 1800 to 114 at present), it became very difficult to study the chemistry of each element and its countless compounds individually.

To solve this problem, scientists looked for a way to organize their knowledge by classifying the elements. This systematic arrangement, known as the Periodic Table, helps us:

• Organize and understand the chemical facts about elements.
• See trends and relationships between different elements.
• Predict the properties of new or undiscovered elements.

GENESIS OF PERIODIC CLASSIFICATION

The development of the periodic table was a gradual process, built on the work of many scientists.

Dobereiner's Triads

In the early 1800s, German chemist Johann Dobereiner noticed that some elements could be grouped into sets of three, which he called Triads.

• Property of Triads: The elements in a triad had similar chemical properties. The atomic weight of the middle element was approximately the average of the atomic weights of the other two elements.
• Example: In the triad of Lithium (Li), Sodium (Na), and Potassium (K), the atomic weight of Na (23) is close to the average of Li (7) and K (39). $\frac{7 + 39}{2} = 23$
• Limitation: This rule only worked for a few known elements and was dismissed as a coincidence.

Newlands' Law of Octaves

In 1865, English chemist John Alexander Newlands arranged the elements in order of increasing atomic weight. He observed that every eighth element had properties similar to the first one.

• Musical Analogy: He compared this relationship to the octaves in music, where the eighth note resembles the first. This is known as the Law of Octaves.
• Example: Starting from Lithium (Li), the eighth element is Sodium (Na). Both have similar properties. The next eighth element is Potassium (K), which is also similar.
• Limitation: This law was only true for elements up to Calcium. It was not widely accepted at the time.

Mendeleev's Periodic Law

The modern periodic table owes its development to Russian chemist Dmitri Mendeleev and German chemist Lothar Meyer, who worked independently. Mendeleev is generally credited because he published his findings first and made bold predictions.

Mendeleev's Periodic Law states:

The properties of the elements are a periodic function of their atomic weights.

Mendeleev arranged elements in a table with horizontal rows and vertical columns (groups) based on increasing atomic weight. Elements with similar properties were placed in the same group.

Key Features of Mendeleev's Table:

1. Systematic Arrangement: He used a broader range of physical and chemical properties, especially the formulas of compounds formed with oxygen (oxides) and hydrogen (hydrides).
2. Correcting Atomic Weights: He prioritized grouping elements with similar properties over strictly following the order of atomic weights. For example, he placed Iodine (lower atomic weight) after Tellurium because Iodine's properties were similar to fluorine, chlorine, and bromine in Group VII.
3. Predicting New Elements: Mendeleev left gaps in his table for elements that he believed were yet to be discovered. He predicted the existence and properties of elements he called Eka-aluminium (later discovered as Gallium) and Eka-silicon (later discovered as Germanium). His predictions were remarkably accurate, which made his periodic table famous.

MODERN PERIODIC LAW AND THE PRESENT FORM OF THE PERIODIC TABLE

Mendeleev's table was based on atomic mass. However, in 1913, English physicist Henry Moseley discovered that the atomic number (Z), which represents the number of protons in the nucleus, is a more fundamental property of an element than its atomic mass.

This led to the Modern Periodic Law, which states:

The physical and chemical properties of the elements are periodic functions of their atomic numbers.

This law is the foundation of the modern periodic table. Since the atomic number determines the number of electrons, and the arrangement of electrons (electronic configuration) determines chemical properties, the periodic law is a direct result of the periodic variation in electronic configurations.

The Long Form of the Periodic Table

The most widely used version of the periodic table is the "long form".

• Periods: The horizontal rows are called periods. There are 7 periods in total. The period number corresponds to the highest principal quantum number (n) of the elements in that period.
  • 1st period: 2 elements
  • 2nd & 3rd periods: 8 elements each
  • 4th & 5th periods: 18 elements each
  • 6th period: 32 elements
  • 7th period: Incomplete, but would also have a theoretical maximum of 32 elements.
• Groups: The vertical columns are called groups or families. There are 18 groups. Elements in the same group have similar outer electronic configurations and therefore similar chemical properties.
• Lanthanoids and Actinoids: To keep the table's structure neat, the 14 elements of the 6th period (Lanthanoids) and 7th period (Actinoids) are placed in separate panels at the bottom.

NOMENCLATURE OF ELEMENTS WITH ATOMIC NUMBERS > 100

For newly discovered, highly unstable elements with atomic numbers greater than 100, the International Union of Pure and Applied Chemistry (IUPAC) established a systematic naming system to be used until an official name is approved.

The name is derived directly from the digits of the atomic number using the following numerical roots:

How to Name an Element:

1. Combine the roots for each digit in the atomic number.
2. Add the suffix "-ium" at the end.
3. The symbol consists of the first letter of each root.

Given

• Atomic number (Z) = 120

To Find

• IUPAC temporary name
• IUPAC temporary symbol

Solution

The digits are 1, 2, and 0.

• The root for 1 is "un".
• The root for 2 is "bi".
• The root for 0 is "nil".

Combining the roots and adding "-ium": un + bi + nil + ium = unbinilium. The symbol is formed from the first letters of the roots: u + b + n = Ubn.

Final Answer The name is unbinilium and the symbol is Ubn.

ELECTRONIC CONFIGURATIONS OF ELEMENTS AND THE PERIODIC TABLE

The arrangement of elements in the periodic table is a direct reflection of their electronic configurations. An element's location reveals the quantum numbers of the last electron added to the atom.

Electronic Configurations in Periods

The period number (n) tells you the principal energy level (shell) of the outermost electrons.

• Period 1 (n=1): Fills the 1s orbital. Contains 2 elements (Hydrogen, Helium).
• Period 2 (n=2): Fills the 2s and 2p orbitals. Contains 8 elements (Lithium to Neon).
• Period 3 (n=3): Fills the 3s and 3p orbitals. Contains 8 elements (Sodium to Argon).
• Period 4 (n=4): Fills the 4s, 3d, and 4p orbitals. The filling of the 3d orbitals creates the first series of transition elements. This period has 18 elements.
• Period 5 (n=5): Similar to period 4, it fills the 5s, 4d, and 5p orbitals. It also has 18 elements.
• Period 6 (n=6): Fills the 6s, 4f, 5d, and 6p orbitals. The filling of the 4f orbitals creates the lanthanoid series (inner-transition elements). This period has 32 elements.
• Period 7 (n=7): Fills the 7s, 5f, 6d, and 7p orbitals. The filling of the 5f orbitals creates the actinoid series.

Solution

When the principal quantum number is $n=5$, the available orbitals are 5s, 5p, and 5d. However, the order of filling based on energy is $5s < 4d < 5p$.

• The 5s subshell has 1 orbital, which can hold 2 electrons.
• The 4d subshell has 5 orbitals, which can hold 10 electrons.
• The 5p subshell has 3 orbitals, which can hold 6 electrons.

The total number of electrons that can be accommodated is $2 + 10 + 6 = 18$.

Final Answer Therefore, there are 18 elements in the 5th period.

Groupwise Electronic Configurations

Elements in the same group have similar valence shell electronic configurations (the same number of electrons in their outermost orbitals). This similarity is the reason they have similar chemical properties.

For example, all Group 1 elements (alkali metals) have a valence shell configuration of ns¹:

• Li: [He] 2s¹
• Na: [Ne] 3s¹
• K: [Ar] 4s¹

ELECTRONIC CONFIGURATIONS AND TYPES OF ELEMENTS: S-, P-, D-, F- BLOCKS

Based on which type of orbital the last electron enters, elements can be classified into four blocks.

The s-Block Elements

• Location: Groups 1 (alkali metals) and 2 (alkaline earth metals).
• Electronic Configuration: Outermost configuration is $ns^1$ (Group 1) or $ns^2$ (Group 2).
• Properties:
  • They are all highly reactive metals.
  • They have low ionization enthalpies, meaning they lose their outermost electron(s) easily.
  • They form positive ions: $1+$ for Group 1 and $2+$ for Group 2.
  • Their compounds are predominantly ionic (except for some compounds of Li and Be).

The p-Block Elements

• Location: Groups 13 to 18.
• Electronic Configuration: Outermost configuration varies from $ns^2np^1$ to $ns^2np^6$.
• Properties:
  • This block contains metals, non-metals, and metalloids.
  • The s-block and p-block elements together are called Representative Elements or Main Group Elements.
  • Group 18 (Noble Gases): Have a completely filled valence shell ($ns^2np^6$), making them very stable and chemically unreactive.
  • Group 17 (Halogens): Have highly negative electron gain enthalpies and readily gain one electron to form a stable noble gas configuration.
  • Group 16 (Chalcogens): Tend to gain two electrons.
  • Metallic character increases down a group, while non-metallic character increases from left to right across a period.

The d-Block Elements (Transition Elements)

• Location: Groups 3 to 12, in the center of the periodic table.
• Electronic Configuration: General outer configuration is $(n-1)d^{1-10}ns^{0-2}$. The last electron enters an inner d orbital.
• Properties:
  • They are all metals.
  • They form a bridge between the highly reactive s-block metals and the less reactive p-block elements.
  • They often form colored ions.
  • They exhibit variable valence (oxidation states).
  • Many are used as catalysts.

The f-Block Elements (Inner-Transition Elements)

• Location: The two rows at the bottom of the table.
• Series:
  • Lanthanoids: Following Lanthanum (La), filling the 4f orbitals.
  • Actinoids: Following Actinium (Ac), filling the 5f orbitals.
• Electronic Configuration: General outer configuration is $(n-2)f^{1-14}(n-1)d^{0-1}ns^2$.
• Properties:
  • They are all metals.
  • Properties within each series are very similar.
  • Actinoid elements are radioactive. Elements after uranium are called Transuranium Elements.

Metals, Non-metals and Metalloids

The periodic table also shows a broad classification based on properties:

• Metals: Comprise over 78% of known elements, located on the left side. They are typically solids (except mercury), have high melting points, are good conductors of heat and electricity, are malleable (can be flattened), and ductile (can be drawn into wires).
• Non-Metals: Located on the top right side. They are usually solids or gases with low melting points, are poor conductors, and are brittle (not malleable or ductile).
• Metalloids (or Semi-metals): Elements bordering the zig-zag line that separates metals and non-metals (e.g., Si, Ge, As). They have properties intermediate between metals and non-metals.

PERIODIC TRENDS IN PROPERTIES OF ELEMENTS

The properties of elements show predictable patterns or trends as you move across a period or down a group.

Trends in Physical Properties

Atomic Radius

The size of an atom is difficult to measure precisely because the electron cloud doesn't have a sharp boundary. We estimate it using:

• Covalent Radius: Half the distance between the nuclei of two identical atoms bonded together in a molecule (used for non-metals).
• Metallic Radius: Half the distance between two adjacent metal cores in a metallic crystal (used for metals).

Trends in Atomic Radius:

1. Across a Period (Left to Right): Atomic radius generally decreases.
   • Reason: Electrons are added to the same valence shell, but the number of protons in the nucleus (nuclear charge) increases. This stronger attraction pulls the electrons closer to the nucleus, making the atom smaller.
2. Down a Group (Top to Bottom): Atomic radius increases.
   • Reason: A new principal energy level (shell) is added for each element down the group. These outer electrons are farther from the nucleus. Also, the inner electrons "shield" the outer electrons from the full pull of the nucleus, further increasing the size.

Ionic Radius

The ionic radius is the radius of an atom's ion.

• Cations (Positive Ions): A cation is smaller than its parent atom.
  • Reason: It has lost one or more electrons, so there are fewer electrons being pulled by the same number of protons. The remaining electrons are pulled closer to the nucleus.
• Anions (Negative Ions): An anion is larger than its parent atom.
  • Reason: It has gained one or more electrons. The added electrons increase the repulsion among the electrons in the valence shell, causing the electron cloud to expand.

Isoelectronic Species are atoms and ions that have the same number of electrons (e.g., $O^{2-}, F^-, Na^+, Mg^{2+}$ all have 10 electrons).

• Trend: For isoelectronic species, the radius decreases as the nuclear charge (number of protons) increases. The species with the most protons will pull the electrons in most tightly and will be the smallest.
• Example Order (Largest to Smallest): $O^{2-} > F^- > Na^+ > Mg^{2+}$

Ionization Enthalpy ($\Delta_i H$)

Ionization Enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state. $X(g) \rightarrow X^+(g) + e^-$ It is always a positive value because energy is always needed to remove an electron.

• First Ionization Enthalpy ($\Delta_i H_1$): Energy to remove the first electron.
• Second Ionization Enthalpy ($\Delta_i H_2$): Energy to remove the second electron from a positive ion. It is always higher than the first because it's harder to remove an electron from a positively charged ion.

Trends in Ionization Enthalpy:

1. Across a Period (Left to Right): Ionization enthalpy generally increases.
   • Reason: As atomic size decreases and effective nuclear charge increases, the outermost electrons are held more tightly, requiring more energy to remove them.
2. Down a Group (Top to Bottom): Ionization enthalpy decreases.
   • Reason: The outermost electron is in a higher energy level, farther from the nucleus, and is shielded by inner electrons. It is held less tightly and is easier to remove.

Anomalies in the Trend:

• Be vs. B: Beryllium (Be) has a higher ionization enthalpy than Boron (B). This is because Be has a stable, filled 2s orbital ($1s^22s^2$), while the electron removed from B is a single, less stable 2p electron ($1s^22s^22p^1$).
• N vs. O: Nitrogen (N) has a higher ionization enthalpy than Oxygen (O). Nitrogen has a stable, half-filled 2p subshell ($2p^3$), where each electron is in a separate orbital. In Oxygen ($2p^4$), one 2p orbital contains a pair of electrons, and the repulsion between these two electrons makes it easier to remove one of them.

Electron Gain Enthalpy ($\Delta_{eg} H$)

Electron Gain Enthalpy is the enthalpy change that occurs when an electron is added to a neutral gaseous atom to form a negative ion. $X(g) + e^- \rightarrow X^-(g)$

• A negative $\Delta_{eg} H$ means energy is released (exothermic process), and the atom readily accepts an electron.
• A positive $\Delta_{eg} H$ means energy is absorbed (endothermic process), and the atom does not easily accept an electron.

Trends in Electron Gain Enthalpy:

1. Across a Period (Left to Right): Electron gain enthalpy generally becomes more negative.
   • Reason: As atomic size decreases and nuclear charge increases, the atom has a stronger attraction for an incoming electron. Halogens (Group 17) have the most negative values.
2. Down a Group (Top to Bottom): Electron gain enthalpy generally becomes less negative.
   • Reason: The atom is larger, and the added electron is farther from the nucleus, experiencing a weaker attraction.

Anomaly:

• F vs. Cl: Chlorine (Cl) has a more negative electron gain enthalpy than Fluorine (F), even though F is higher in the group. This is because the F atom is very small. An incoming electron entering the compact 2p subshell experiences significant repulsion from the electrons already present. In the larger Cl atom, the incoming electron enters the 3p subshell where there is more space, and repulsion is less.

Electronegativity

Electronegativity is a measure of the ability of an atom in a chemical compound to attract the shared pair of electrons towards itself. It is a relative value, not a measurable quantity. The most common scale is the Pauling scale, where Fluorine (the most electronegative element) is assigned a value of 4.0.

Trends in Electronegativity:

1. Across a Period (Left to Right): Electronegativity increases.
   • Reason: The increased nuclear charge and smaller atomic radius cause a stronger attraction for electrons.
2. Down a Group (Top to Bottom): Electronegativity decreases.
   • Reason: The larger atomic size means the nucleus is farther from the shared electrons, resulting in a weaker attraction.

Relationship to Metallic Character:

• Electronegativity is inversely related to metallic character. Elements with low electronegativity are metals (they tend to lose electrons).
• Electronegativity is directly related to non-metallic character. Elements with high electronegativity are non-metals (they tend to gain or share electrons).

Periodic Trends in Chemical Properties

Periodicity of Valence or Oxidation States

Valence is the combining capacity of an element. For representative elements, it is often equal to the number of valence electrons or eight minus the number of valence electrons.

Oxidation State is the charge an atom would have in a molecule if electrons were completely transferred based on electronegativity differences.

• In $OF_2$, Fluorine is more electronegative, so its oxidation state is -1. Oxygen must be +2 to balance the two fluorine atoms.
• In $Na_2O$, Oxygen is more electronegative, so its oxidation state is -2. Each Sodium atom must be +1.

The valence of elements shows a periodic trend. For example, the formula of hydrides across the third period changes predictably: $NaH, MgH_2, AlH_3, SiH_4, PH_3, H_2S, HCl$.

Anomalous Properties of Second Period Elements

The first element in each of the s- and p-block groups (Li, Be, B, C, N, O, F) shows properties that are different from the other members of its group. Reasons for Anomalous Behavior:

1. Small Size: They are much smaller than other elements in their groups.
2. High Electronegativity: They have the highest electronegativity in their groups.
3. Large Charge/Radius Ratio: This leads to more covalent character in their compounds.
4. Absence of d-orbitals: They only have s and p orbitals (4 valence orbitals total), so their maximum covalency is 4. Other members of the group have d-orbitals and can expand their valence shell (e.g., Al forms $[AlF_6]^{3-}$ but B only forms $[BF_4]^-$).

Diagonal Relationship: The first element of a group often shows similarities to the second element of the next group (e.g., Li is similar to Mg, and Be is similar to Al). This is called a diagonal relationship.

Periodic Trends and Chemical Reactivity

Chemical reactivity is also a periodic property.

• Across a Period: Reactivity is highest at the two extremes and lowest in the middle.
  • Extreme Left (Group 1): Highly reactive metals. They have very low ionization enthalpies and readily lose an electron to form a cation.
  • Extreme Right (Group 17): Highly reactive non-metals. They have very negative electron gain enthalpies and readily gain an electron to form an anion.
• Down a Group:
  • For Metals (e.g., Group 1): Reactivity increases. As size increases, the ionization enthalpy decreases, making it easier to lose the valence electron.
  • For Non-metals (e.g., Group 17): Reactivity decreases. As size increases, the ability to attract an electron (electronegativity) decreases.
• Nature of Oxides:
  • Across a Period: The oxides of elements change from basic on the left (e.g., $Na_2O$), to amphoteric in the middle (e.g., $Al_2O_3$), to acidic on the right (e.g., $Cl_2O_7$).
    • $Na_2O + H_2O \rightarrow 2NaOH$ (a strong base)
    • $Cl_2O_7 + H_2O \rightarrow 2HClO_4$ (a strong acid)