Chemistry of the Nonmetals
Approximately 18 elements are classified as nonmetals, they lie above
and to the right stairway that runs diagonally across the periodic table.
As the word nonmetal applies implies, these elements do not show metallic
properties, in the solid they are brittle as opposed to ductile, insulators
rather than conductors.
Most of the nonmetals, particularly those in Groups 15 to 17 of the periodic
table are molecular in nature. The noble gases (Group 18) consists of
individual atoms attracted to each other by weak dispersion forces. Carbon in
Group 14 has a network covalent structure.
The Elements and Their Preparation (N, P, O, S, F, Cl, Br, I)
Molecular – N2, O2, F2, Cl2 – those of low molar mass are gases at room
temperature and atmospheric pressure.
Stronger dispersion forces cause the nonmetals of higher molar mass to be
either liquids (Br2) or solids (I2, P4, S8).
Physical Properties of Nonmetals
1. Physical State: Most of the non-metals exist in two of the three states
of matter at room temperature: gases (oxygen) and solids (carbon).
Only bromine exists as a liquid at room temperature.
2. Non-Malleable and Ductile: Non-metals are very brittle, and cannot be
rolled into wires or pounded into sheets.
3. Conduction: They are poor conductors of heat and electricity.
4. Luster: These have no metallic luster and do not reflect light.
5. Melting and Boiling Points: The melting points of non-metals are
generally lower than metals, but are highly variable.
6. Seven non-metals exist under standard conditions as diatomic
molecules: H2(g) , N2(g) , O2(g) , F2(g) , Cl2(g) , Br2(l) , I2(s) .
Properties of Nonmetallic Elements
Nitrogen
Phosphorus
Oxygen
Sulfur
Fluorine
Chlorine
Bromine
Iodine
2s22p3
3s23p3
2s22p4
3s23p4
2s22p5
3s23p5
4s24p5
5s25p5
Molecular formula
N2
P4
O2
S8
F2
Cl2
Br2
I2
Molar mass (g/mol)
28
124
32
257
38
71
160
254
State (25°C, 1 atm)
gas
solid
gas
solid
gas
gas
liquid
solid
Melting point (°C)
-210
44
-218
119
-220
-101
-7
114
Boiling point (°C)
-196
280
-183
444
-188
-34
59
184
Bond Enthalpy*
(kJ/mol)
941
200
498
226
153
243
193
151
-
-
-
-
+2.889V
+1.360V
+1.077V
+0.534
Outer electron
configuration
E°red
Chemical Reactivity
Of the eight nonmetals, nitrogen is by far the least reactive. Its inertness is
due to the strength of the triple bond holding the N2 molecule together
(N ≡ N 941 kJ/mol). This same factor explains why virtually all explosives are
compounds of nitrogen.
Fluorine is the most reactive of all elements, in part because of the weakness
of the F – F bond (F – F = 153 kJ/mol), but mostly it is such a powerful
oxidizing agent (E°red = +2.889V). Fluorine combines with every element in the
periodic table except He and Ne.
Chlorine is somewhat less reactive than fluorine. Although it reacts with
nearly all metals, heating is often required. This reflects the relatively strong
bond in the Cl2 molecule (Cl – Cl = 243 kJ/mol)
K-expression: Equilibrium Constant Expression
Equilibrium – a state of dynamic balance in which rate of forward and reverse
reactions are equal, the system does not change with time.
𝑎𝐴 𝑔 + 𝑏𝐵 𝑔 ↔ 𝑐𝐶 𝑔 + 𝑑𝐷(𝑔)
Where A, B, C and D represent different substances and a,b,c and d are their
coefficients in the balanced equation,
𝑃𝐶
𝐾=
𝑃𝐴
𝑐
𝑎
𝑃𝐷
𝑃𝐵
𝑑
𝑏
Where PC, PD, PA and PB are the partial pressures of the four gases at
equilibrium. These partial pressures must be expressed in atmospheres
K-expression:
• Gases enter as partial pressures in atmospheres
• Aqueous species enter as concentrations in molarity
• Water is not included
• Pure liquids and solids are not included in the expression
• Numerator – products (right side of the equation)
• Denominator – reactants (left side of the equation)
• Each partial pressure is raised to a power equal to its coefficient in
the balanced equation
Problem 8-1. Consider the reaction between chlorine gas and water.
𝐶𝑙2 𝑔 + 𝐻2 𝑂 ↔ 𝐶𝑙 − 𝑎𝑞 + 𝐻 + 𝑎𝑞 + 𝐻𝐶𝑙𝑂(𝑎𝑞)
a. Write the expression for the equilibrium constant K
𝐻𝐶𝑙𝑂 𝐶𝑙 − 𝐻 +
𝐾=
𝑃𝐶𝑙2
b. At the temperature of the reaction, K = 2.7 x 10-5. What are the
concentration of HClO in equilibrium with chlorine gas at 1.0 atm? 0.030M
Problem 8-2. Ammonium Chloride is sometimes used as flux in soldering
because it decomposes on heating:
𝑁𝐻4 𝐶𝑙 𝑠 ↔ 𝑁𝐻3 𝑔 + 𝐻𝐶𝑙(𝑔)
The HCl formed removes oxide films from metals to be soldered. In certain
equilibrium system at 400°C, 22.6 g of NH4Cl is present, the partial pressures of
NH3 and HCl are 2.5 atm, and 4.8 atm, respectively. Calculate K at 400 °C.
K =12
Hydrogen Compound of Nonmetals
• Ammonia, NH3 – used to make fertilizers and host of different nitrogen
compounds, notably nitric acid, HNO3. The NH3 molecule act as a BronstedLowry base in water, accepting proton from a water molecule. Ammonia
can also act as a Lewis base when it reacts with a metal cation to form
complex ion.
• Hydrogen Sulfide, H2S – in water solution, hydrogen sulfide acts as a
Bronsted-Lowry acid, it can donate a proton to a water molecule.
• Hydrogen Peroxide, H2O2 – in hydrogen peroxide, oxygen has an oxidation
number of -1, intermediate between the extremes for the element 0, and 2. This means that H2O2 act as an either an oxidizing agent, in which case it
is reduced to H2O, or as a reducing agent, when it is oxidized to O2.
• Hydrogen Fluoride, HF and Hydrogen Chloride, HCl – most common
hydrogen halides
Problem 8-3. Write the equation of the reaction between hydrogen and
phosphorus.
P2(g) + 3H2(g) → 2PH3(g)
Problem 8-4. Write the equation of the reaction between hydrogen and
bromine.
H2(g) +Br2(g) → 2HBr(g)
Lewis Structure – an electronic structure of a molecule or ion in which
electrons are shown by dashes or dots (electron pairs).
There are two kinds of electron pairs.
1. A pair of electrons shared between two atoms is a covalent bond,
ordinarily shown as a straight line between bonded atoms.
2. An unshared electron pair, owned entirely by one atom, is shown as a pair
of dots on that atom.
Covalent bond – a chemical link between two atoms produced by sharing
electrons in the region between atoms.
Single bond – a single electron pair is shared between two bonded atoms.
Double bond – occurs when bonded atoms can share two electron pairs.
Triple bond – three pairs of electrons are shared.
Octet rule – bonded atoms tend to have a share in eight valence electrons.
Writing Lewis Structures
1. Draw a skeleton of the species joining atoms by single bonds
• The central atom is usually written first in the formula.
• The terminal atoms are most often hydrogen, oxygen and halogens.
2. Count the number of valence electrons
• For molecule – add the number of valence electrons of all atoms present
• For polyatomic anion – add the number of valence electrons of each atom plus one
electron for each unit of negative charge
• For polyatomic cation - add the number of valence electrons of each atom and
subtract one electron for each unit of negative charge
3. Count the number of valence electrons available for distribution
• AE = VE – 2(number of bonds in the skeleton)
4. Count the number of electrons required to fill out an octet for each
atom (except H) in the skeleton
• If AE = NE , your skeleton is correct. Distribute available electrons as unshared
pairs satisfying the octet rule
• If AE < NE, modify your skeleton by changing single bonds to double or triple
bonds
*Hydrogen and halogens never form double bonds
Draw the Lewis structures of the following:
OCl- , C2H6 , NCl3, NO2- , N2O
Oxygen compound of Nonmetals
Molecular Structure of Nonmetal Oxides
The Lewis structure of the oxides of
nitrogen are shown in the figure.
For example:
Acid Strength – the acid equilibrium constants of the oxoacids of the halogens
are listed below
Notice that the value of Ka increase with
• Increasing oxidation number of the central atom (HClO<HClO2< HClO3< HClO4)
• Increasing electronegativity of the central atom (HIO<HBrO< HClO)
Electronegativity – ability of an atom to attract itself the electron forming a covalent bond
Problem 8-5. Consider sulfurous acid, H2SO3
a. Show its Lewis structure and that of the HSO3- and SO32- ions.
b. How would its acid strength compare with that of H2SO4?H2TeO3?
S = 2.6, Te = 2.5
Ans: H2SO4 is a stronger acid than H2SO3
References:
Masterton, W.L., et. al(2018). Principles and Reactions: Chemistry for
Engineering Students: Quezon City
https://chem.libretexts.org/Under_Construction/Essential_Chemistry_(Curri
ki)/Unit_1%3A_Atomic_and_Molecular_Structure/1.4%3A_Electron_Config
uration_and_Orbital_Diagrams
https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Che
mistry__The_Central_Science_(Brown_et_al.)/07._Periodic_Properties_of_the_Ele
ments/7.6%3A_Metals%2C_Nonmetals%2C_and_Metalloids