s-Block and p-Block Elements notes

MDCAT Chemistry

This chapter explains periodic trends, the arrangement of elements into s, p, d and f blocks, and the properties and reactions of Groups I, II and IV. It also covers important compounds and applications related to these groups, including oxides, hydroxides, carbon compounds, nitrogen dioxide, arsenic poisoning of catalysts, and common hydrated salts.

Periodic Properties and Their Trends

Periodic properties are properties that show regular variation with increasing atomic number. The main causes of these trends are effective nuclear charge, atomic size, shielding effect and the number of occupied shells.

Across a period, nuclear charge increases while electrons are added to the same shell. Down a group, new electron shells are added, so atomic size generally increases and the outer electron is farther from the nucleus.

Atomic radius generally decreases from left to right across a period because effective nuclear charge increases.
Atomic radius increases from top to bottom in a group because new electron shells are added.
The correct size comparison is Ba > Mg because barium has more occupied electron shells.
Ionization enthalpy is the energy required to remove an electron from an isolated gaseous atom in its ground state.
Ionization enthalpy generally increases from left to right across a period and decreases down a group.
Removing an electron from an orbital with a higher principal quantum number is generally easier than removing one from an orbital with a lower principal quantum number.
The statement that removal from an orbital with lower n is easier than from a higher n orbital is incorrect.
The first ionization enthalpy of Be is greater than that of Li because Be has a more stable filled 2s subshell. Thus, the statement IE1 of Li > IE1 of Be is incorrect.
Electron affinity generally becomes more negative across a period, although there are exceptions due to stable subshell arrangements.
Electronegativity generally increases across a period and decreases down a group. Fluorine has the highest electronegativity.

Blocks of the Periodic Table

The block of an element is determined by the subshell into which its last electron enters. The periodic table contains s-block, p-block, d-block and f-block elements.

s-Block elements are mainly Groups I and II. Their valence electrons enter an s subshell. p-Block elements are Groups III to VIII, or Groups 13 to 18, and their last electron enters a p subshell.

s-Block elements have general outer electronic configurations ns1 or ns2.
Group I elements are called alkali metals and have the general configuration ns1.
Group II elements are called alkaline earth metals and have the general configuration ns2.
p-Block elements have general outer configurations ns2np1 to ns2np6.
The last group of the p-block contains noble gases with a complete valence shell, except helium, which has 1s2.
d-Block elements are called transition elements and have electrons entering the (n-1)d subshell.
Transition metals show strong metallic bonding because both s and d electrons can take part in bonding.
f-Block elements include lanthanides and actinides. Their differentiating electron enters an f subshell.
Metals are generally present on the left side, nonmetals on the right side, and metalloids along the zigzag boundary.
Arsenic and antimony are metalloids because they show both metallic and nonmetallic properties.

Group I Elements and Their General Properties

Group I elements include lithium, sodium, potassium, rubidium, caesium and francium. They have one electron in the outermost shell and commonly form M+ ions.

Their single valence electron is removed easily. Therefore, alkali metals are highly reactive and are not found in the free state in nature. Reactivity generally increases down the group.

The general electronic configuration of Group I elements is [noble gas] ns1.
Group I elements have oxidation state +1 in their common compounds.
Melting points and boiling points generally decrease down Group I because metallic bonding becomes weaker as atomic size increases.
Alkali metals are soft and can be cut with a knife. Their softness generally increases down the group.
They are good reducing agents because they lose their outer electron readily.
Sodium does not normally show a +2 oxidation state because removal of its second electron requires a high second ionization enthalpy from a stable noble gas configuration.
Ionic mobility in aqueous solution depends on the hydrated size of the ion. Among alkali metal ions, Rb+ has maximum ionic mobility because its hydrated ion has relatively small effective size.
Lithium shows anomalous behaviour because of its very small size, high charge density, high ionization enthalpy and strong hydration energy.
Lithium has a diagonal relationship with magnesium and differs from the other Group I elements in several properties.

Reactions of Group I Elements

Alkali metals react readily with water, oxygen, hydrogen and halogens. The products depend on the metal and the reaction conditions. Their compounds are mainly ionic and soluble in water.

Lithium differs from sodium and potassium in some reactions. For example, lithium forms lithium nitride directly with nitrogen, while the other common alkali metals do not react with nitrogen under ordinary conditions.

With water: 2Na + 2H2O → 2NaOH + H2. The reaction produces heat and hydrogen gas.
Potassium reacts more vigorously with water than sodium, while lithium reacts comparatively slowly because of its high melting point and protective oxide layer.
With oxygen, lithium mainly forms Li2O, sodium commonly forms Na2O2, and potassium, rubidium and caesium can form superoxides such as KO2.
With hydrogen, alkali metals form ionic hydrides. For example, 2Na + H2 → 2NaH.
With chlorine, alkali metals form chlorides. For example, 2K + Cl2 → 2KCl.
Lithium reacts with nitrogen to form lithium nitride: 6Li + N2 → 2Li3N.
Potassium permanganate has the formula KMnO4.
Alkali metal carbonates are generally stable to heat, but lithium carbonate decomposes on heating: Li2CO3 → Li2O + CO2.
The solubility of alkali metal hydroxides generally increases down the group.

Group II Elements and Their General Properties

Group II elements include beryllium, magnesium, calcium, strontium, barium and radium. Their atoms have two valence electrons and commonly form M2+ ions.

Group II elements are less reactive than Group I elements because removal of two electrons requires more energy. Their compounds are often ionic, but the compounds of beryllium show greater covalent character.

The general electronic configuration of Group II elements is [noble gas] ns2.
Group II elements commonly have oxidation state +2.
Their melting points are generally higher than those of Group I elements because Group II atoms provide two binding electrons for metallic bonding.
Atomic size increases down Group II. Thus, Ba is larger than Mg.
Ionization enthalpy decreases down the group because the valence electrons are farther from the nucleus.
Electrical conductivity increases from Be to Ra because metallic character and the number of available shells increase down the group.
Beryllium compounds show covalent character due to the small size and high charge density of Be2+.
The basic strength of Group II hydroxides increases down the group because the metal hydroxides become more ionic.
The solubility of Group II hydroxides generally increases down the group.
The solubility of Group II sulfates generally decreases down the group. BaSO4 is very insoluble.

Reactions and Compounds of Group II Elements

Group II metals react with oxygen, water, acids and halogens. The ease of reaction generally increases down the group, although beryllium behaves differently because of its protective oxide layer and strong bonding.

Many Group II compounds occur as hydrated salts. Their crystal water can be removed by heating. The behaviour of hydroxides and carbonates on heating is also important.

Magnesium burns in oxygen to form magnesium oxide: 2Mg + O2 → 2MgO.
Calcium reacts with water to form calcium hydroxide and hydrogen: Ca + 2H2O → Ca(OH)2 + H2.
Beryllium does not react readily with water because of its protective oxide film.
Group II metals react with acids to form salts and hydrogen. For example, Mg + 2HCl → MgCl2 + H2.
Group II carbonates decompose on heating: MCO3 → MO + CO2.
Group II nitrates generally decompose on heating to give the metal oxide, nitrogen dioxide and oxygen.
Strontium hydroxide is dehydrated to strontium oxide on heating: Sr(OH)2 → SrO + H2O.
Epsom salt is magnesium sulfate heptahydrate, MgSO4.7H2O.
Blue vitriol or blue stone is copper sulfate pentahydrate, CuSO4.5H2O. Copper belongs to the d-block, but this hydrated salt is a commonly studied compound.
The flame colours of calcium, strontium and barium compounds are brick red, crimson red and apple green, respectively.

Group IV Elements and Their Properties

Group IV, also called Group 14, contains carbon, silicon, germanium, tin and lead. Their valence shell configuration is ns2np2. The group changes from nonmetallic character at the top to metallic character at the bottom.

Carbon and silicon commonly show a +4 oxidation state. Tin and lead can show both +2 and +4 states. The stability of the +2 state increases down the group because of the inert pair effect.

The general electronic configuration of Group IV elements is ns2np2.
Carbon is a nonmetal, silicon and germanium are metalloids, and tin and lead are metals.
Group IV elements show oxidation states +4 and +2, with the +2 state becoming more stable down the group.
Catenation is the ability of an element to form chains or rings with atoms of the same element. Carbon shows the strongest catenation.
Carbon forms strong multiple bonds because of its small atomic size and effective p orbital overlap.
Carbon monoxide, CO, has one lone pair on the oxygen atom in its commonly represented Lewis structure.
Lead suboxide has the formula Pb2O.
The inert pair effect is the reluctance of the ns2 electrons to participate in bonding, especially in heavier p-block elements.
The metallic character and atomic size generally increase down Group IV.
Covalent character is more prominent in the compounds of carbon and silicon, while ionic character increases in compounds of the heavier elements.

Important p-Block Compounds, Industrial Chemistry and Environmental Chemistry

Several p-block compounds have important laboratory, industrial and environmental roles. Their formulas, reactions and properties are connected with oxidation states and molecular structure.

Nitrogen oxides affect air quality. Other gases affect the atmosphere by absorbing infrared radiation. Industrial catalysts can be poisoned by impurities.

Nitrogen dioxide, NO2, absorbs visible light and gives smoggy air a brown tint.
NO2 is a greenhouse gas and an air pollutant. Nitrogen and oxygen are not greenhouse gases in their normal diatomic forms.
Concentrated nitric acid reacts with tin to form metastannic acid, commonly represented as hydrated tin(IV) oxide, H2SnO3, along with nitrogen oxide products depending on conditions.
In the Contact Process for sulfuric acid manufacture, V2O5 is used as the catalyst for the oxidation of SO2 to SO3.
Arsenic impurities act as a poison for the V2O5 catalyst in the Contact Process.
Argon is used in electric bulbs to prevent oxidation of the hot tungsten filament because it is an inert gas.
Cinnabar is mercury(II) sulfide, HgS, and is an important ore of mercury.
Ammonium chloride is added before ammonium hydroxide in Group III salt analysis. Its NH4+ ions decrease the ionization of NH4OH by the common ion effect.
The common ion effect reduces the concentration of OH− ions and helps precipitate Group III hydroxides selectively.
Carbon dioxide, CO2, is a greenhouse gas, while N2 and O2 are not greenhouse gases.

Key terms

Periodic property: A property of elements that changes regularly with increasing atomic number.
Effective nuclear charge: The net positive attraction experienced by an electron after shielding by other electrons is considered.
Atomic radius: A measure of the size of an atom, usually represented by half the distance between nuclei of bonded identical atoms.
Ionization enthalpy: The energy required to remove an electron from an isolated gaseous atom or ion.
Electron affinity: The energy change when an electron is added to an isolated gaseous atom.
Electronegativity: The tendency of an atom in a compound to attract the shared electron pair towards itself.
s-Block: The part of the periodic table in which the last electron enters an s subshell.
p-Block: The part of the periodic table in which the last electron enters a p subshell.
Alkali metal: A Group I element that forms a soluble hydroxide and commonly produces a +1 ion.
Alkaline earth metal: A Group II element that commonly forms a +2 ion and a basic oxide or hydroxide.
Metalloid: An element showing a mixture of metallic and nonmetallic properties, such as arsenic or antimony.
Diagonal relationship: A similarity between certain elements of neighbouring periods and groups, such as lithium and magnesium.
Hydration energy: The energy released when gaseous ions become surrounded by water molecules.
Catenation: The ability of an element to form covalent bonds with atoms of the same element.
Inert pair effect: The reluctance of the outer ns2 electron pair of heavier p-block elements to take part in bonding.
Common ion effect: The decrease in ionization of a weak electrolyte caused by adding a strong electrolyte containing a common ion.
Catalyst poison: An impurity that reduces or destroys the activity of a catalyst.
Greenhouse gas: A gas that absorbs outgoing infrared radiation and contributes to warming of the atmosphere.

Test yourself on s-Block and p-Block Elements

Free s-Block and p-Block Elements MCQs with an explanation on every answer. No account needed.

Practise s-Block and p-Block Elements MCQs

More for s-Block and p-Block Elements in the MDCAT pack

A one-page revision sheet for this chapter
5 s-Block and p-Block Elements mnemonics
Chapter-wise Ratta Cards and a Quiz Builder for your own tests

See the MDCAT packAlready have it? Open Revision

Chemistry shortcuts

Finding the limiting reactant and percentage composition

Convert every given mass or volume into moles first. The reactant that produces the least amount of the required product is the limiting reactant.

Write the balanced equation and calculate moles using n = mass/Mr.
Use the mole ratio to calculate the product. For percentage composition, use percentage = mass of element in one mole of compound divided by molar mass, multiplied by 100.
Example: Percentage of nitrogen in KNO3 = 14/101 × 100 = 13.86%.
Answer: 13.86% nitrogen.

Use gas volume at molar volume only when the gas conditions are stated or are standard conditions.

Using gas volume, pressure and temperature relations

At the same temperature and pressure, gas volume is directly proportional to the number of molecules. For changing conditions, use P1V1/T1 = P2V2/T2.

At constant temperature and pressure, divide or multiply the volume in the same ratio as the number of molecules.
Example: 10 mL H2 contains 2 × 10^3 molecules. Oxygen in 200 mL contains 20 × 2 × 10^3 = 4 × 10^4 molecules.
Answer: 4 × 10^4 molecules.
For a rigid container, increasing temperature increases molecular speed and mean free path if the gas remains in the same phase.

The direct volume to molecule ratio does not apply when temperature or pressure changes.

16 more Chemistry shortcuts are in the MDCAT pack. Already have it? See all shortcuts