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Loading home…Mendeleev’s law: properties of elements are a periodic function of their atomic masses.
He arranged elements in periods and groups and left gaps for undiscovered elements.
Limitation 1: position of isotopes was unclear because they have different masses but identical properties.
Limitation 2: anomalous pairs (e.g. Ar/K, Te/I) did not fit strictly by increasing atomic mass.
Some dissimilar elements appeared together and some similar ones were separated.
These problems were resolved by the modern law based on atomic number.
The modern periodic law states that the physical and chemical properties of elements are a periodic function of their atomic numbers.
Atomic number (Z) is the number of protons in the nucleus.
Moseley’s work on X-ray spectra established atomic number as the correct basis.
The modern long-form table arranges elements in order of increasing Z.
Periods and groups then reflect repeating outer electronic configurations.
This removes Mendeleev’s isotope and anomalous-pair difficulties.
Groups are vertical columns; elements in a group have the same valence-shell configuration pattern.
Periods are horizontal rows; period number equals the highest principal quantum number being filled.
There are 7 periods and 18 groups in the long-form table.
s-block: Groups IA–IIA (1–2); p-block: Groups IIIA–VIIIA (13–18).
d-block: transition elements (Groups 3–12); f-block: lanthanides and actinides.
Blocks classify elements by the subshell that receives the differentiating electron.
Across a period, nuclear charge increases while electrons enter the same shell.
Attraction to the nucleus increases, so atomic radius generally decreases.
Down a group, a new shell is added and shielding increases.
The outermost electrons are farther from the nucleus, so radius increases.
Cations are smaller than parent atoms; anions are larger.
Trends help compare reactivity and bond lengths.
Ionization energy (IE) is the energy needed to remove the most loosely held electron from a gaseous atom.
Across a period, IE generally increases as nuclear charge rises and radius falls.
Down a group, IE decreases because radius and shielding increase.
Exceptions occur at half-filled/full subshells (e.g. N vs O, Be vs B patterns).
Successive IEs jump sharply after valence electrons are removed.
High IE marks non-metals; low IE marks metals.
d-block elements have their differentiating electron in a d-subshell.
They lie between s- and p-blocks and show a transition in properties.
Many have partially filled d-orbitals in atoms or common ions.
They often show variable oxidation states, coloured ions, and catalytic activity.
They form complexes readily due to available d-orbitals.
Typical examples include Fe, Cu, Ni and Zn (Zn is sometimes discussed as a borderline case).
The modern periodic law states that properties of elements are a periodic function of atomic number.
Elements are arranged in order of increasing Z in the long-form periodic table.
A period is a horizontal row; there are seven periods.
A group is a vertical column of elements with similar valence configurations.
s-block elements have valence configuration ns¹ or ns² (Groups IA and IIA).
p-block elements fill np orbitals (Groups IIIA to VIIIA / 13–18).
d-block elements are transition metals with (n−1)d and ns electrons being filled.
f-block elements are inner transition metals (lanthanides and actinides) filling 4f/5f.
Block classification follows the subshell that receives the last (differentiating) electron.
Periodicity of properties arises because outer electronic structures repeat.
Metals dominate the left and centre; non-metals are at the upper right.
Noble gases in Group VIIIA have filled valence shells and low reactivity.
Hydrogen is unique and often placed separately or above Group IA.
Understanding blocks helps predict bonding type and common oxidation states.
Exam answers should state the law, define period/group, and give block examples.
Mention Moseley’s role in replacing atomic mass with atomic number.
Relate Group VA general configuration ns² np³ as a p-block example.
State that the third period contains eight elements (Na to Ar).
Conclude that the modern table organises chemical behaviour systematically.
Keep nuclear charge, shielding and distance as the three explanatory factors in every trend answer.
Atomic radius generally decreases across a period due to increasing effective nuclear charge.
Atomic radius increases down a group as shells and shielding increase.
First ionization energy generally increases across a period and decreases down a group.
Smaller atoms hold electrons more tightly, so IE is higher.
Electronegativity (EN) is the tendency of an atom in a bond to attract shared electrons.
EN generally increases across a period and decreases down a group.
Fluorine has the highest electronegativity among the elements.
Caesium (and francium) are among the most electropositive metals.
Electropositivity is the tendency to lose electrons; it opposes EN trends.
Noble gases are often excluded from simple EN scales because they form few compounds.
Metallic character increases down a group and decreases across a period.
Non-metallic character shows the opposite pattern.
These trends explain acid–base character of oxides and reactivity of metals/non-metals.
Always cite nuclear charge, shielding, and distance as the three key factors.
Give one concrete example for each trend (e.g. Li→F radius fall; F highest EN).
Note exceptions at half-filled/full subshells for IE.
Transition metals show smaller horizontal changes because d electrons shield imperfectly.
A clear exam answer defines each property then states both period and group trends.
Link EN to bond polarity and IE to ease of cation formation.
Summarise with a short “left/right, up/down” mental map.
Ionization energy depends mainly on nuclear charge, atomic radius, and shielding.
Higher nuclear charge increases IE; larger radius and more shielding decrease IE.
Subshell stability (full or half-filled) can raise IE relative to neighbours.
The first IE removes the outermost electron; the second IE removes the next, and so on.
Each successive IE is larger because the ion is more positive and holds electrons more tightly.
A very large jump occurs after all valence electrons have been removed.
Example pattern: for Mg (ns²), IE₁ and IE₂ are moderate; IE₃ is much larger.
That jump shows magnesium has two valence electrons.
For Na (ns¹), IE₂ is far larger than IE₁, confirming one valence electron.
Across a period, increasing nuclear charge raises IE despite the same shell.
Down a group, extra shells and shielding lower IE.
Gaseous atoms are the standard reference state for IE definitions.
Units are often kJ mol⁻¹ in school tables.
IE data help place elements as metals (low IE) or non-metals (high IE).
They also support electronic configuration deduced from the periodic table.
Exam answers should list factors, define successive IE, and interpret one jump.
Relate shielding to why IE falls down Group IA alkali metals.
Mention that removing an electron from a cation is harder than from a neutral atom.
Do not confuse IE with electron affinity or electronegativity.
Conclude that successive IE is a tool to find the number of valence electrons.
Write the electronic configuration and identify the valence shell.
The period number equals the highest principal quantum number n in the configuration.
The block is the subshell being filled: s, p, d or f.
For s- and p-block, the group follows the number of valence electrons.
Example: ns² np³ corresponds to Group VA (15), such as N, P, As.
Example: ns¹ is Group IA (alkali metals).
For d-block, group number relates to (n−1)d + ns electrons (with known exceptions).
Atomic number Z places the element uniquely in order.
Valence configuration predicts common oxidation states and bonding type.
Metals tend to lose electrons; non-metals tend to gain or share.
Noble-gas configurations (ns² np⁶) mark Group VIIIA.
Transition metals often show variable oxidation states due to close ns/(n−1)d energies.
Comparing configurations of neighbours explains periodic similarities.
Isotopes share Z and thus the same table position despite different mass.
Exam method: config → period → block → group → properties.
Give one full worked example (e.g. 1s² 2s² 2p⁶ 3s² 3p⁵ → Period 3, p-block, Group VIIA).
State that Moseley’s atomic number idea underpins this placement.
Mention that incomplete or excited configs are not used for standard placement.
Position prediction is a core FBISE skill linking structure to the table.
Always double-check electron count equals Z.
Metallic character is the tendency to lose electrons and form cations.
Non-metallic character is the tendency to gain or share electrons.
Across a period, metallic character decreases and non-metallic character increases.
Down a group, metallic character increases (especially in Groups IVA–VIA).
Alkali metals (Group IA) are the most strongly metallic family.
Halogens and oxygen-family non-metals are strongly non-metallic.
Caesium is among the most electropositive (metallic) elements.
Fluorine is the most electronegative and highly non-metallic.
Diagonal relationships (e.g. Li–Mg) show intermediate behaviour.
Metalloids along the staircase show mixed properties.
Oxide character follows the trend: metal oxides are basic; non-metal oxides are acidic.
Amphoteric oxides appear near the metalloid region.
Reactivity of metals often increases down Group IA/IIA as IE falls.
Reactivity of non-metals often decreases down a group as EN and oxidizing power fall.
These trends unify radius, IE, and EN ideas into chemical behaviour.
Use examples: Na more metallic than Mg; Cl more non-metallic than P in the same period.
State exceptions carefully for transition metals.
Exam answers should define both characters and give period + group trends.
Link electropositivity of Cs to easy electron loss.
End with a short summary map of the table’s metal/non-metal regions.