Study sheet: Periodic Table and Periodic Trends

📋 Course Outline

1. Historical Development of the Periodic Table
2. Modern Table Structure and Classification
3. Blocks and Element Families
4. Periodic Arrangement and Electron Configuration
5. Atomic and Ionic Radii
6. Ionization Energy
7. Electron Affinity and Electronegativity
8. Period Three Reactions and Bonding
9. Oxides, Chlorides, and Oxidation States

📅 Key Dates

1. 1829Döbereiner grouped elements into triads of three elements with similar properties, in which the atomic weight of the middle element was roughly the average of the other two.
2. 1864John Newlands arranged 62 known elements in increasing atomic mass and observed that every eighth element resembled the first in properties.
3. 1869Dmitri Mendeleev arranged 63 elements by increasing atomic mass, left gaps for undiscovered elements, and accurately predicted their masses and properties.
4. 1913Moseley used X-ray emission to determine exact atomic numbers and rearranged the elements by atomic number, leading to the modern Periodic Law.

📖 1. Historical Development of the Periodic Table

📝 Essential Points

• The modern periodic table contains 118 elements arranged in increasing atomic number and is described as the “Symbol of Chemistry.”

💡 Memory Hook

Lavoisier → Döbereiner → Newlands → Mendeleev → Moseley

📖 2. Modern Table Structure and Classification

★ Must-know

• The modern periodic table has seven horizontal rows called periods and eighteen vertical columns called groups.

📌 Elements in the same group have similar chemical properties because they have the same number of valence electrons, but their physical properties change gradually from top to bottom.

📌 Metals tend to lose electrons and form positive ions, nonmetals tend to gain electrons and form negative ions, and metalloids display properties of both metals and nonmetals.

Further detail

📌 Elements within a period show a gradual change in properties from left to right.

💡 Memory Hook

Groups share chemical properties; periods show gradual variation

📖 3. Blocks and Element Families

★ Must-know

📌 The s-block contains elements whose valence electrons occupy s subshells, the d-block contains transition elements, the p-block contains groups 13–18, and the f-block contains lanthanides and actinides.

• Alkali metals are Group 1 elements Li, Na, K, Rb, Cs, and Fr; they have one valence electron, form alkalis with water, and are the most reactive metals.

• The alkaline earth metals are:

  • Be
  • Mg
  • Ca
  • Sr
  • Ba
  • Ra

• Halogens are Group 17 elements F, Cl, Br, I, At, and Ts; they are highly reactive nonmetals that readily accept one electron and form salts.

Further detail

• The noble gases are:
  • He
  • Ne
  • Ar
  • Kr
  • Xe
  • Rn
  • Og

💡 Memory Hook

s-d-p-f blocks

📖 4. Periodic Arrangement and Electron Configuration

★ Must-know

📌 The period number indicates the number of electron shells, while the group number indicates the number of valence electrons for the relevant main-group elements.

• An element in Group 13 and Period 3 has three valence electrons in the third shell and the configuration 1s² 2s² 2p⁶ 3s² 3p¹.

Further detail

• An element in Period 3 and Group 2 is magnesium because it has three shells, two valence electrons, and a 3s² valence configuration.

💡 Memory Hook

Period number gives shells; group number gives valence electrons → configuration

📖 5. Atomic and Ionic Radii

🔑 Key Concepts & Definitions

• Atomic radius : half the distance between two identical bonded atoms and is commonly measured in picometres or angstroms

📝 Essential Points

📌 Atomic radius generally decreases from left to right across a period because increasing nuclear charge pulls the electron cloud closer.

📌 Atomic radius increases down a group because additional shells increase shielding and enlarge the atom despite the higher nuclear charge.

📌 A cation is generally smaller than its neutral atom because electron loss removes a shell and strengthens nuclear attraction on the remaining electrons, whereas an anion is generally larger because added electrons increase repulsion.

💡 Memory Hook

Across a period: smaller; down a group: larger

📖 6. Ionization Energy

🔑 Key Concepts & Definitions

• Ionization energy : decreases down a group because increasing atomic size and shielding make valence electrons easier to remove, and it increases from left to right across a period because effective nuclear charge increases

📝 Essential Points

• In Group 1, first ionization energy decreases in the order Li > Na > K > Rb > Cs.

📌 Spin-pair repulsion slightly lowers ionization energy because paired electrons repel one another, making one paired electron easier to remove.

• Chromium has the configuration [Ar] 3d⁵4s¹, whereas manganese has [Ar] 3d⁵4s², so removing a paired 4s electron from manganese is relatively easier.

💡 Memory Hook

More shielding and larger size → easier electron removal → lower ionization energy

📖 7. Electron Affinity and Electronegativity

🔑 Key Concepts & Definitions

• Electron affinity : the enthalpy change when one mole of gaseous atoms gains one mole of electrons to form one mole of gaseous uninegative ions
• Electronegativity : Linus Pauling — Electronegativity is the power of an atom to attract a shared pair of electrons toward itself in a molecule, and the Pauling scale gives fluorine a maximum value of 4.0 and alkali metals a minimum value of 0.8.

★ Must-know

📐 Formula — For oxygen, the first electron affinity is $\Delta H^\circ_{ea^1} = -142\ \mathrm{kJ\ mol^{-1}}$, whereas the second is $\Delta H^\circ_{ea^2} = +844\ \mathrm{kJ\ mol^{-1}}$ because the negative ion repels the incoming electron.

📌 Electron affinity generally becomes more negative from left to right across a period and decreases down a group as atomic size and distance from the nucleus increase.

Further detail

📐 Formula — For chlorine, the first electron affinity is represented by $\mathrm{Cl(g) + e^- \rightarrow Cl^-(g)}$ with $\Delta H^\circ_{ea^1} = -348.8\ \mathrm{kJ\ mol^{-1}}$.

💡 Memory Hook

Electron affinity concerns added electrons; electronegativity attracts shared electrons

📖 8. Period Three Reactions and Bonding

★ Must-know

📐 Formula — Sodium reacts with water according to $\mathrm{2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)}$.

📌 Oxides and chlorides of Groups 1–3 are predominantly ionic, whereas those of Groups 4–7 are more covalent because electronegativity increases across the period.

Further detail

📐 Formula — Sodium burns in oxygen to form sodium peroxide according to $\mathrm{2Na(s) + O_2(g) \rightarrow 2Na_2O_2(s)}$, while limited oxygen or high temperature can produce sodium oxide.

📐 Formula — Sodium and magnesium react with chlorine to form soluble salts: $\mathrm{2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)}$ and $\mathrm{Mg(s) + Cl_2(g) \rightarrow MgCl_2(s)}$.

💡 Memory Hook

Water → oxygen → chlorine

📖 9. Oxides, Chlorides, and Oxidation States

🔑 Key Concepts & Definitions

• Basic oxide : reacts with water to produce an alkali and is usually ionic; Group 1 and Group 2 metals form basic oxides such as Na₂O, CaO, and BaO
• Acidic oxide : An acidic oxide reacts with water to produce an acid and is generally a covalent nonmetal oxide, such as SO₂, SO₃, P₂O₃, or P₂O₅.
• Amphoteric oxide : An amphoteric oxide reacts with both acids and bases; aluminium oxide reacts with hydrochloric acid and sodium hydroxide.

★ Must-know

📌 Group 1 and Group 2 chlorides are generally neutral in water, whereas chlorides from aluminium to sulfur in Period 3 hydrolyze to produce acidic solutions.

📌 The oxidation number of a Period 3 element in an oxide or chloride is positive because oxygen and chlorine are more electronegative than the Period 3 elements.

• In Period 3 oxides, oxidation number increases from +1 in sodium to +6 in sulfur, while in chlorides it increases from +1 in sodium to +5 in phosphorus.

Further detail

• Sulfur has oxidation number +4 in SO₂ and +6 in SO₃ because four or six electrons, respectively, are used for bonding.

💡 Memory Hook

Basic oxides react with water to form alkalis; acidic oxides form acids

📊 Synthesis Tables

Periodic Trends

Property	Across a period	Down a group
Atomic radius	Decreases	Increases
Ionization energy	Increases	Decreases
Electron affinity	Generally becomes more negative	Generally decreases
Electronegativity	Increases	Decreases

Oxide and Chloride Character

Category	Composition or position	Behavior
Basic oxide	Usually ionic metal oxide	Forms an alkali with water
Acidic oxide	Usually covalent nonmetal oxide	Forms an acid with water
Amphoteric oxide	Example: Al₂O₃	Reacts with both acids and bases
Neutral chloride	Groups 1–2 generally	Produces a nearly neutral solution
Acidic chloride	Aluminium to sulfur in Period 3	Hydrolyzes to produce an acidic solution