Chemical Properties of Elements Explained With Examples

Chemical property of elements
Every compound or any naturally-occurring substance is made out of the fundamental elements, be it on Earth, any other planet, or any material in the entire universe (except dark matter as its composition is unknown, and it does not constitute the elementary particles like protons, neutrons, and electrons). The chemical properties of an element account for its potential to undergo a chemical change or a reaction depending upon its composition.
Fast Fact
The first periodic table, which looked similar to the one used presently, was made by Dmitri Mendeleev in 1869. It was designed with regards to the order of increasing atomic weight, and hence, several physical and chemical properties of elements showed repeated occurrence afterwards, in a cyclic manner. The empty spaces were marked with question marks to indicate undiscovered elements.
An element can be distinguished by its chemical properties like the atomic number, atomic mass, electronic configuration, etc., which depend on the number of protons (positively charged ions) present in its nucleus. The term 'element' is generally referred to as a pure chemical substance consisting of atoms, with an equal number of protons and electrons. Till now, 118 elements have been discovered, out of which 98 occur naturally on our planet. For example, elements like carbon, copper, iron, hydrogen, nitrogen, silver, oxygen, etc., are used for numerous applications in our day-to-day lives.
What are the Chemical Properties of an Element?
Atomic Number
This property denotes the number of protons present in an atom's nucleus, and it is one of the fundamentals in the field of chemistry. An element and its position in the periodic table depends on its atomic number, and each element is positioned subsequently with reference to the increase in this property. When an atom is electrically neutral, the atomic number equalizes the number of electrons (negatively charged ions) present around the nucleus. When an atom is electrically charged, the atomic number is different as compared to the number of protons present.
✤ Example
Atomic number of iron
The atomic number of iron element (Fe) is 26, and this means that it is the 26th element placed in the periodic table according to the conventional order.
Atomic Mass
It represents the mass of an atom, which is measured in atomic mass units (amu). It depends on the protons and neutrons present in the elemental nucleus. The weight of each proton or neutron is about 1 amu. Hence, the atomic mass is nearly equal to the mass or nucleon number, as the weight of every proton and neutron is around 1 amu. The atomic mass of an isotope represents the number of neutrons present in the nucleus. In the periodic table, the atomic mass written below the element sign indicates the average of all the atomic mass values of all isotopes of that element. The addition of mass units of all isotopes of an element gives us the total atomic mass of that specific element.
✤ Example
Atomic mass of silver and bromine
The atomic mass of silver is 107, while that of bromine is 80.
Atomic Radius
This property can be simply defined as the distance between the nucleus and the 1st electron shield surrounding the elemental core. Atomic radius is a characteristic of an element, which shows the most exact trend, if moving along this table in a certain direction is taken into account. It is measured and represented in the unit called picometers. Thus, the ionic or atomic radii generally decrease as we move along the periods from left to right, but increase along the groups from top to bottom. Generally, the length of the ionic radii is directly proportional to the atomic number and atomic mass of the specific element, if we take into account the latter direction (moving down a group). There are various types of atomic radii like ionic radius, covalent radius, van der Waals radius, etc., as the boundary of electrons in the first shell is quite indistinct.
✤ Example
Examples of covalent radius of lithium and beryllium are Li (145 pm) and Be (105 pm).
This property was first discovered by Linus Pauling in 1932, and is simply defined as the ability of an atom to attract electrons from another atom. It mainly increases as one moves along the periodic table from left to right, and decreases along the elements from top to the bottom with reference to the groups. This property is directly proportional to the atomic number of an element, as well as the distance between the valence electrons and the atomic nucleus. This property also shows some exception in the periodic table; the elements aluminum (Al) and Silicon (Si) show a lower value of electronegativity as compared to Germanium (Ge) and Gallium (Ga). This is because of the shortening of the d-block in the table, which is situated near the elements Al and Si.
✤ Example
According to the related present-day studies, fluorine and cesium have the highest and lowest electronegativity, respectively.
Electron Affinity
When an uncharged atom accepts an electron, it gets converted into an anion, as electrons are negatively charged. In this case, a specific amount of energy is released, and the value of this energy is known as electron affinity of an element. This property mainly decreases as one moves along the periods of the periodic table. On the contrary, it shows a mixed result when going along the groups from top to bottom. For example, in the 1st period, this property decreases as one moves towards the downward direction. However, in the 3rd period, the elements having a larger nucleus exhibit higher electron affinity than the lighter elements. Metallic elements mainly show less positive values of this property than the non-metallic ones.
✤ Example
Chlorine gas has the highest value of electron affinity in the periodic table.
Electronic Configuration
It indicates the arrangement of the electrons in the form of shells or circles around the core. Each shell has a fixed energy level. Electrons away from the core have high energy levels, whereas the ones closer to the nucleus have lower levels of energy. The electronic configuration changes according to the atomic number of the element. After the shell is filled till the limit, the next one starts taking electrons or loses them according to the electron affinity of the specific element. The electronic shells are further divided into sub-shells, which are named as s, p, d, f, and g. All the elements in the periodic table show a kind of repeated pattern of occurrence after specific breaks.
✤ Example
The electronic configuration of sodium (Na) is 1s2, 2s2, 2p6, 4s.
Energy of Ionization
When an atom or a molecule gives out an electron for bonding with another atom or an element, it requires energy for this purpose. This is known as the energy of ionization and is mainly divided into several successive rounds. The 1st round is mainly concerned with measuring the energy needed while removing the outermost shell electrons, whereas the subsequent measurements are required for the second, third, fourth ... nth shell electrons. These values indicate the resistance offered by an atom, in order to give up the electrons. In noble gases, the ionization energy values are quite high, as there is a lot of resistance in removal of electrons from such elements. In general, this property shows an increase in values as we move from the left to the right of the periodic table.
✤ Example
The third row elements show a very large increase in these energies between the 1st and 2nd rounds of ionization (740 kJ/mol and 1450 kJ/mol, respectively).
Metallic Character
This property can be defined as the inherent ability of an element to show specific metallic characteristics, depending on its position, electronic configuration, atomic radii, etc. After extensive studies, it has been found out that the other chemical properties like electronegativity, electron affinity, and ionization energy have a profound influence regarding development of metallic characteristics in an element. Thus, lower the values of these properties, higher are the chances for the specific element to behave as a metal. Generally, the metallic character in the periodic table goes on decreasing if one goes on from left to right along the periods.
✤ Example
Elements like rubidium and strontium show strong metallic characters, while elements like oxygen and nitrogen show strong non-metallic characters.
Oxidation States
The property that denotes the ability of an element to take part in an oxidation reaction with minimum resistance is known as the oxidation state. It is a virtual theoretical charge of an element, which basically indicates its affinity to bond with oxygen. Many elements can have numerous oxidation states, and they are mostly represented as integers, though few also are written down as fractions. The elements having tetroxides have the highest oxidation state (+8), whereas the carbon group of elements have the lowest value of oxidation state (-4).
✤ Example
If sulfur loses 2 electrons in an oxidation reaction, this means that its oxidation state is +2, as the element was oxidized to lose those couple of electrons.
Standard Potential
It is defined as the potential of a redox reaction, when the elemental state is at equilibrium. If the potential value increase above zero, it is an oxidation reaction, otherwise it is a reduction reaction. The standard potential is measured in volts (V), and is expressed by the symbol V0. Standard potential property is very useful regarding the reactions that take place at the electrodes. This property varies with changes in temperature and pressure of the surrounding environment, and also with the concentration of the particular element that takes part in the reaction. While calculating this property, it is important to have a knowledge of the reaction that is going on at the reduction end, more than at the oxidation end.
✤ Example
In case of the formation of sodium fluoride, the potential of fluorine atom is negative as it accepts electrons, and hence, a reduction reaction takes place at the fluorine end.
The above-described parameters reflect an element's quality, which becomes evident during a chemical reaction. These properties help understand the nature and behavior of an element under varying conditions. Various fields like thermodynamics, organic chemistry, analytical chemistry, etc., need a basic knowledge about all these parameters.