FSc Part 1 Chemistry Chapter 1 — Periodic Table and Periodic Properties

The periodic table is called the “Symbol of Chemistry”. The modern table holds 118 elements, arranged in tabular form by atomic number. Many elements — gold, silver, iron, phosphorus, sulfur, zinc and arsenic — were known since prehistoric times, but the first real classification came in the 18th century.

• Antoine Lavoisier attempted to classify the known elements into metals and nonmetals.
• Döbereiner (1829) grouped elements into triads — groups of three with similar properties — noting that the atomic weight of the middle element was roughly the average of the other two. Example: lithium, sodium, potassium (⁷Li, ²³Na, ³⁹K).
• John Newlands (1864) was the first to observe periodicity: arranging the 62 known elements in increasing order of atomic masses, he found that every eighth element resembled the first in properties.
• Lothar Meyer plotted atomic weight against atomic volumes in the same year; the curves also showed periodicity.
• Dmitri Mendeleev (1869), the father of the Periodic Table, arranged 63 elements by increasing atomic mass, aligning similar properties into groups — and left gaps for undiscovered elements, predicting their masses and properties accurately. When those elements were found, his predictions proved correct.
• Moseley (1913) determined exact atomic numbers using X-ray emission and removed the flaws of Mendeleev's table by arranging elements by atomic number instead of atomic mass. The Periodic Law was modified to: “the properties of elements are periodic functions of their atomic numbers.”

The modern periodic table makes the properties of elements easier to understand. Its main features are:

• 118 elements are grouped in ascending order of atomic number.
• There are seven horizontal rows called periods and eighteen vertical columns called groups. (In older versions there were eight vertical groups, divided into A-Groups and B-Groups.)
• Elements in the same group show similar chemical properties because they have the same number of valence electrons. Their physical properties, however, change gradually from top to bottom in a group.
• Elements in a period show a gradual change in properties moving from left to right.
• Other ways of grouping — blocks, families and categories — give further understanding of the elements.

Remember the rule of thumb: down a group = similar chemistry, slowly changing physics; across a period = gradual change in everything.

Elements are broadly classified into three kinds:

• Metals tend to lose electrons to form positive ions. Examples: iron, copper, gold, silver.
• Nonmetals tend to gain electrons to form negative ions. Examples: chlorine, sulfur, phosphorus.
• Metalloids separate the metals from the nonmetals and show some properties of metals and some of nonmetals.

Most periodic tables show a “stair-step line”. The line begins at boron (B) and extends down to polonium (Po), including Si, Ge, As, Sb and Te. Elements to the left of the line are metals; elements just to the right of the line show properties of both metals and nonmetals and are called metalloids or semimetals; elements far to the right are nonmetals. The exception is hydrogen, the first element on the table.

Elements can be classified by the subshell containing their valence electrons:

• The valence electrons of the elements in the first two groups are in the “s” subshell — these are s-block elements.
• Transition elements belong to the d-block.
• The elements in the two series at the bottom of the table — the Lanthanides and Actinides — are f-block elements.
• The remaining elements, in groups 13 to 18 including the inert gases in the last group, belong to the p-block.

Knowing the block an element belongs to tells us about its characteristics, chemical reactivity, oxidation states, electronegativity, ionization energy and electron filling.

An element family is a set of elements sharing common properties. There are five famous families:

• i) Alkali metals — Group 1 (Li, Na, K, Rb, Cs, Fr): they produce alkalis when they react with water, have one valence electron, and are the most reactive metals.
• ii) Alkaline earth metals — Group 2 (Be, Mg, Ca, Sr, Ba, Ra): metals found mainly in the earth, forming alkalis; they have two valence-shell electrons and are less reactive than alkali metals.
• iii) Transition metals: the largest family, in the middle of the table — four series of d-block elements plus the lanthanides and actinides below. Examples: Cr, Fe, Co, Ni, Cu, Zn. They show variable oxidation states and mostly form coloured compounds.
• iv) Chalcogens — Group 16 (O, S, Se, Te, Po, Lv): named because most copper ores (Greek chalkos) are oxides or sulfides. Oxygen and sulfur are nonmetals, Se, Te and Po are metalloids, and livermorium is a metal.
• v) Halogens — Group 17 (F, Cl, Br, I, At, Ts): “salt-formers” — they react easily with alkali and alkaline earth metals to form salts. They are highly reactive nonmetals with high electron affinities and need only one electron to complete their outermost shell.
• vi) Noble gases — Group 18 (He, Ne, Ar, Kr, Xe, Rn, Og): unreactive elements at the extreme right. Due to their complete outermost shells, they are almost entirely unreactive under normal conditions and rarely form compounds.

An element's position tells us its electronic structure:

• Period number = shell number (n) — the number of electron shells around the nucleus. For example, an element X in the 3rd period has three electron shells, with its valence electrons in the 3rd shell. Which subshell holds them depends on the element's block: if X in the 3rd period is in the s-block, its valence electrons are in the 3s subshell.
• Group number = number of valence electrons. For example, an element X in the 3rd period and group 2 has two valence electrons in its outermost shell. So X (3rd period, group 2, s-block) has two valence electrons in the 3s subshell — that is magnesium (Mg).
• Another example: X belongs to group 13 and period 3 — it has 3 valence electrons and 3 shells. Its configuration is: 1s² 2s² 2p⁶ 3s² 3p¹ — aluminium (Al).

So the table works both ways: from position you can deduce configuration, and from configuration you can find the position.

The atomic radius measures the size of an atom: it is half the distance between two identical atoms bonded together. It can vary with the type of bond (covalent, metallic, van der Waals) and the state of the atom, and is typically measured in picometer (pm) or Angstrom (Å).

Factors affecting atomic radius: number of shells, effective nuclear charge, and the shielding effect of inner electrons.

Periodic trends:

• Across a period (left → right), atomic radius decreases because increasing nuclear charge pulls the electron cloud closer.
• Down a group (top → bottom), atomic radius increases because additional electron shells are added, so more shielding makes the atom larger despite the increase in nuclear charge (which is outweighed).

The ionic radius measures the size of an ion in a crystal lattice: the distance from the nucleus of an ion to its outermost electron shell, measured in pm or Å.

• When an atom loses electrons to become a positive ion (cation), it generally becomes smaller than the neutral atom — the loss of an electronic shell and stronger pull of the nucleus on the remaining electrons.
• When an atom gains electrons to become a negative ion (anion), it generally becomes larger than the neutral atom — added electrons increase electron repulsion, so the electron cloud expands.

Trends: Moving across a period left to right, the ionic radius of cations decreases (increasing nuclear charge pulls electrons closer) and anions also decrease across a period for the same reason. Moving down a group, both cations and anions increase in size because the shell number (n) increases, so the distance between nucleus and outermost electrons grows, outweighing the increased nuclear charge. Book values include: Li⁺ 60 pm (atom 152), Na⁺ 95 (186), K⁺ 133 (231), Mg²⁺ 65 (160), Ca²⁺ 99 (197), Al³⁺ 50 (143), F⁻ 136 (64), Cl⁻ 181 (99), O²⁻ 140 (66), S²⁻ 184 (104).

First ionization energy (ΔH_i1) is “the energy needed to remove one electron from each atom in one mole of atoms of the element in the gaseous state to form one mole of gaseous 1+ ions.”

Na_(g) → Na⁺_(g) + e⁻   ΔH_i1 = 494 kJ mol⁻¹
Ca_(g) → Ca⁺_(g) + e⁻   ΔH_i1 = 590 kJ mol⁻¹
Ca⁺_(g) → Ca²⁺_(g) + e⁻   ΔH_i2 = 1150 kJ mol⁻¹
Ca²⁺_(g) → Ca³⁺_(g) + e⁻   ΔH_i3 = 4940 kJ mol⁻¹

An element can have as many ionization energies as its atomic number. Factors affecting ionization energy:

• i) Nuclear charge: greater effective nuclear charge → stronger attraction → harder to remove an electron → higher ionization energy.
• ii) Size of atom or ion: in bigger atoms the attraction between nucleus and outermost electrons is weaker → ionization energy decreases as size increases.
• iii) Electronic arrangement: half-filled and completely-filled orbitals are stable, so ionization energy is higher when the electron is removed from a fully-filled or half-filled subshell. (a) Noble gases have the highest ionization energies in their periods due to highly stable fully-filled shells (ns² np⁶). (b) Oxygen has lower ionization energy than nitrogen due to the half-filled subshell of nitrogen: ₇N = 1s² 2s² 2p_x¹ 2p_y¹ 2p_z¹ (more stable), ΔH_i1 = 1403 kJ mol⁻¹; ₈O = 1s² 2s² 2p_x² 2p_y¹ 2p_z¹ (less stable), ΔH_i1 = 1365 kJ mol⁻¹.
• iv) Shielding effect: greater shielding → easier to remove valence electrons → lower ionization energy.
• v) Spin-pair repulsion: when electrons are spin-paired in the same orbital, repulsion makes removal slightly easier. Manganese (Mn) has two spin-paired electrons in its 4s orbital, so its ionization energy is relatively lower; chromium (Cr) has one unpaired electron in 4s, so removing it needs more energy (no spin-pairing repulsion). ₂₄Cr = [Ar] 3d⁵4s¹ (more stable); ₂₅Mn = [Ar] 3d⁵4s² (less stable).

Periodic trends: Going down a group, ionization energy decreases (atomic size and shielding increase). In Group 1: Li > Na > K > Rb > Cs. Moving across a period left to right, ionization energy increases (shells unchanged, effective nuclear charge increases). Noble gases have the highest values in a period and alkali metals the lowest.

First electron affinity (ΔH^o_ea1) is the enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous uni-negative ions under standard conditions.

Cl_(g) + e⁻ → Cl⁻_(g)   ΔH^o_ea1 = −348.8 kJ mol⁻¹. Energy is released, so the first electron affinity carries a negative sign.

Second electron affinity (ΔH^o_ea2) is the energy required to add electrons to 1 mole of uni-negative gaseous ions to form 1 mole of gaseous −2 ions. For oxygen, the first addition releases energy (ΔH^o_ea1 = −142 kJ mol⁻¹), but adding a second electron to the negative O⁻ ion needs 844 kJ mol⁻¹ because the negatively charged ion repels the incoming electron: O⁻_(g) + e⁻ → O²⁻_(g), ΔH^o_ea2 = +844 kJ mol⁻¹.

Factors affecting electron affinity: (i) Size of atom — smaller atom → stronger nuclear attraction for the incoming electron → greater electron affinity. (ii) Nuclear charge — greater nuclear charge → stronger attraction → greater electron affinity. (iii) Electronic configuration — electron affinity is low when the electron is added to a half-filled subshell rather than a partially filled one. The values for nitrogen and phosphorus (group 15) are very low because of their half-filled “np” subshells in the valence shell (N = 2s² 2p³, P = 3s² 3p³); these very stable half-filled p-subshells have very little tendency to accept an extra electron.

Periodic trends: As atomic size increases down a group, the incoming electron feels less attraction, so electron affinity generally decreases down the group — observed in halogens (At < I < Br < Cl < F). Electron affinities become more negative as we move left to right in a period, firstly due to increased nuclear charge and secondly due to decreasing atomic radius. Book Table 1.1 (kJ mol⁻¹): Group 17 — F −328.0, Cl −349.0, Br −324.0, I −295.0, At −270.1; Group 1 — Li −60.0, Na −53.0, K −48.0, Rb −47.0, Cs −46.0.

Electronegativity is the power of an atom to attract the shared pair of electrons toward itself in a molecule. Linus Pauling developed a scale of dimensionless electronegativity values. On this scale, alkali metals have minimum electronegativity (0.8) and halogens have maximum values (F = 4.0).

Factors affecting electronegativity: (i) Atomic size — a larger atomic size gives a lower electronegativity, because electrons far from the nucleus feel a weaker force of attraction. (ii) Effective nuclear charge — a higher effective nuclear charge gives a greater electronegativity, because increased nuclear charge attracts bonded electrons more strongly. This is why electronegativity increases across a period from left to right: Li in period 2 is 1.0 and F is 4.0.

Periodic trends: Moving left to right along a period, electronegativity increases (increasing nuclear charge, decreasing size). In groups it decreases from top to bottom (increasing size and shielding). For example, in the halogen group, electronegativity decreases from fluorine (4.0) to iodine (2.5).

Normally, metals on the left side of the table have lower electronegativity values than nonmetals. Hence, metals are electropositive and nonmetals are electronegative, relatively.

The metallic character of elements is typically their tendency to lose electrons. Elements on the left side of the table have a greater tendency to lose their outermost electrons to achieve noble gas configuration. Elements on the right side tend to gain electrons. So the metallic character of an element depends largely on its valence shell electronic configuration.

Consequently, metallic character decreases across a period from left to right: increasing nuclear charge pulls the electron cloud closer to the nucleus, making it more difficult to lose electrons; also, across the period the nuclear charge increases while atomic size decreases, strengthening attraction to valence electrons. Contrarily, metallic character increases down a group because atomic size and shielding increase, reducing nuclear attraction on valence electrons. The increase in metallic character (ease of losing electrons) makes the element more reactive — hence cesium is far more reactive and electropositive than sodium or lithium.

1.8.1 With Water: Sodium is more reactive than magnesium towards water. Na reacts vigorously with water to form sodium hydroxide and hydrogen, while Mg reacts more slowly to form magnesium hydroxide and hydrogen. However, magnesium reacts with steam more vigorously to make magnesium oxide and hydrogen gas.

2Na_(s) + 2H₂O_(l) → 2NaOH_(aq) + H_2(g)
Mg_(s) + 2H₂O_(l) → Mg(OH)_2(aq) + H_2(g)
Mg_(s) + 2H₂O_(g) → MgO_(s) + 2H_2(g)

1.8.2 With Oxygen: Sodium burns in oxygen with a golden yellow flame to produce a white solid mixture of sodium oxide and sodium peroxide. Sodium is kept under kerosene oil to prevent its reaction with air. It reacts vigorously with oxygen in open air to form peroxide: 2Na_(s) + O_2(g) → Na₂O_2(s). Under special conditions like limited O₂ or high temperature, sodium oxide is formed: 4Na_(s) + O_2(g) → 2Na₂O_(s). Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide: 2Mg_(s) + O_2(g) → 2MgO_(s).

1.8.3 With Chlorine: Chlorine reacts with both metals to give soluble salts. It reacts exothermically with sodium, with a golden yellow flame, to give white solid sodium chloride; magnesium also reacts with chlorine to give white solid magnesium chloride.

2Na_(s) + Cl_2(g) → 2NaCl_(s)
Mg_(s) + Cl_2(g) → MgCl_2(s)

1.9 Trends in bonding: Oxides of groups 1, 2 and 3 (e.g. Na₂O) have more ionic character — they exist as giant ionic lattices with strong electrostatic forces between oppositely charged ions. Oxides of groups 4, 5, 6 and 7 (e.g. SO₂) are more covalent — they exist as covalent molecules with weak intermolecular forces. This transition is a result of increasing electronegativity and decreasing ionic character. Similarly, chlorides of groups 1, 2 and 3 (e.g. NaCl) are predominantly ionic, while chlorides of groups 4, 5, 6 and 7 (e.g. PCl₃) are covalent; covalent character increases due to decrease in electronegativity difference between the halogen and the other atom.

1.9.1 Classification of oxides: Oxides are binary compounds formed by the reaction of oxygen with other elements, classified as neutral, amphoteric, basic or acidic by their characteristics.

• i) Basic oxides: an oxide that when combined with water gives off an alkali. Metals react with oxygen to give basic oxides, usually ionic. Group 1 and 2 form basic oxides when they react with oxygen. Examples: Na₂O, CaO, BaO. Group 2 hydroxides' solubility increases down the group, so alkalinity also increases down the group.
  Na₂O_(s) + H₂O_(l) → 2NaOH_(aq)
  CaO_(s) + H₂O_(l) → Ca(OH)_2(aq)
• ii) Acidic oxides: an oxide that when combined with water gives an acid. Nonmetals react with oxygen to form acidic oxides held together by covalent bonds. Silicon dioxide is an acidic oxide as it can react with bases. Examples in period 3: P₂O₃, P₂O₅, SO₃, SO₂.
  P₂O_5(s) + H₂O_(l) → H₃PO_4(aq)
  SO_2(g) + H₂O_(l) → H₂SO_3(aq)
  Reactions with bases: SiO_2(s) + 2NaOH_(aq) → Na₂SiO_3(aq) + H₂O_(l); P₂O_5(s) + 6NaOH_(aq) → 2Na₃PO_4(aq) + 3H₂O_(l); SO_2(g) + 2NaOH_(aq) → Na₂SO_3(aq) + H₂O_(l)
• iii) Amphoteric oxides: oxides that can react with both acids and bases — they can behave as either an acid or a base depending on conditions. Aluminium oxide (Al₂O₃) is insoluble in water but reacts with hydrochloric acid to form aluminium chloride and water, and with sodium hydroxide to form sodium aluminate and water.
  Al₂O_3(s) + 6HCl_(aq) → 2AlCl_3(s) + 3H₂O_(l)
  Al₂O_3(s) + 2NaOH_(aq) → 2NaAlO_2(s) + H₂O_(l)

1.9.2 Classification of chlorides: Chlorine forms compounds with other elements known as chlorides. These show characteristic behaviour when added to water, resulting in solutions that can be acidic or neutral.

• i) Neutral chlorides: salts that, when dissolved in water, produce a neutral solution with a pH close to 7. At the start of period 3, sodium and magnesium chlorides do not react with water. The solutions contain positive metal ions and negative chloride ions surrounded by water molecules — now known as hydrated ions; this process is hydration.
  NaCl_(s) → Na⁺_(aq) + Cl⁻_(aq)   (pH = 7)
  MgCl_2(s) → Mg²⁺_(aq) + 2Cl⁻_(aq)   (pH = 6.5)
  Group 1 and group 2 chlorides are also neutral, with few exceptions.
• ii) Acidic chlorides: If we move in period 3, from aluminium to sulfur, all chlorides react with water to make an acidic solution with pH less than 7 — this process is called hydrolysis. When AlCl₃ is added to water, aluminium and chloride ions are in solution. Al³⁺ ion is hydrated and causes a water molecule to lose an H⁺ ion, making the solution acidic.
  AlCl_3(s) + 3H₂O_(l) → Al(OH)_3(s) + 3HCl_(aq)
  Other examples: SiCl_4(l) + 2H₂O_(l) → SiO_2(s) + 4HCl_(aq); PCl_3(l) + 3H₂O_(l) → H₃PO_3(s) + 3HCl_(aq)

1.10 Variation in oxidation number in oxides and chlorides: The oxidation number of an atom is the formal charge on that atom in a molecule or ion (also called oxidation state). In ionic compounds, an atom's oxidation number is the charge it appears to have. The oxidation number of an element of the 3rd period in its oxide or chloride corresponds to the number of electrons used for bonding and is always positive, because oxygen and chlorine are more electronegative than any of these elements. It matches the group number, reflecting the total number of valence electrons.

Book Table 1.2 — oxides: Na in Na₂O +1, Mg in MgO +2, Al in Al₂O₃ +3, Si in SiO₂ +4, P in P₄O₁₀/P₂O₅ +5/+3, S in SO₃ +6, S in SO₂ +4; chlorides: Na in NaCl +1, Mg in MgCl₂ +2, Al in AlCl₃ +3, Si in SiCl₄ +4, P in PCl₅ +5, P in PCl₃ +3, S in SCl₂ +2. In oxides the oxidation number increases from +1 in Na to +6 in S; in chlorides it increases from +1 in Na to +5 in P. Phosphorus and sulfur exhibit several oxidation numbers because they can expand their octet by exciting electrons into empty 3d orbitals — for instance, in SO₂ sulfur has +4 because only four electrons are used for bonding, while in SO₃ sulfur has +6 because all six electrons are used for bonding.