Unit - 5

Periodic properties

Q1) Define effective Nuclear charge?

A1) The term effective Nuclear charge is the net total charge that an electron experiences in an atom in the presence multiple electrons. The effective nuclear charge may be estimated by the equation:

Zeff = Z - S

Where Z is the atomic number and S is the number of shielding electrons.

Electrons of higher can possess different electrons of lower energy between the electron and the nucleus, thereby lowering the positive charge experienced by the high energy electron.

The shielding effect is named as it gives a balance between the repulsion between valence and inner electrons the attraction of valence electrons. The shielding effect also explains the trend in atomic size followed in the periodic table and also explains why valence electrons are readily removed from an atom.

Q2) Which orbital has the highest energy?

A2) The 1s orbital has the highest energy. the energy of an electron is the energy it would require to rip it out of the atom’s electronic cloud So the electrons that are closest to the nucleus are the ones that are most attracted to it:.1s is the orbital that is closest to the atom’s nucleus

Q3) What is Hund rule?

A3) It states that:

1. In a sublevel, each orbital is singly occupied before it is doubly occupied.

2. The electrons present in singly occupied orbitals possess identical spin.

The electrons enter an empty orbital before pairing up. The electrons repel each other as they are negatively charged. The electrons do not share orbitals to reduce repulsion.

When we consider the second rule, the spins of unpaired electrons in singly occupied orbitals are the same. The initial electrons spin in the sub-level decides what the spin of the other electrons would be. For instance, a carbon atom’s electron configuration would be 1s22s22p2. The same orbital will be occupied by the two 2s electrons although different orbitals will be occupied by the two 2p electrons in reference to Hund’s rule.

Q4) What kind of orbital must and electron with the principal quantum number n=2 occupy?

A4) Each such orbital can be occupied by a maximum of two electrons, each with its own spin quantum number s. The simple names s orbital, p orbital, d orbital and f orbital refer to orbitals with angular momentum quantum number ℓ = 0, 1, 2 and 3 respectively.

Q5) Why does the ionic radius increase as you move down a group?

A5) As you move across a row of the periodic table, the ionic radius decreases for metals forming cations, as the metals lose their outer electron orbitals. The ionic radius increases for non-metals as the effective nuclear charge decreases due to the number of electrons exceeding the number of protons. Ionic radius increases as you move from top to bottom on the periodic table

Q6) What is atomic and ionic radius?

A6) Atomic radii or size can be defined in the most simplified manner as the radii (or half the "width") of the spherical atoms. Nonbonding atoms have a much larger, more indefinite or "fuzzy" radius, therefore when atomic radius is described as a periodic trend, the thing that comes to mind is bonding atomic radius.

These are the radii of atoms that are chemically bonded to each another. So, if the bond between two Cl atoms in Cl2 is 1.99 angstroms, resulting in chlorine's bonding atomic radius as about 0.99 angstrom

Ionic size Cations, or positively charged ions, are smaller in size than their "parent" atoms. This is because cations are formed when the outermost orbitals become empty due to absence of electrons. This also decreases electron-electron repulsions. Therefore, the resulting ions are smaller as there are not as many orbitals that are occupied and the effective nuclear charge affecting the remaining electrons increases, pulling electrons in more closely.

Anions, or negatively charged ions, are larger in size than their "parent" atoms. This is because of the addition of electrons form these ions, increasing electron-electron repulsions, making the electrons spread out more effectively. Also, effective nuclear charge felt by the outermost electrons decrease. For ions that carry the same charge (i.e. from parent elements of the same group), size increases from top to bottom as we go down a group.

Q7) Explain Ionization energy?

A7) The ionization energy of a chemical category (i.e., an atom or molecule) is the energy required to remove electrons from gaseous ions or atoms. This property is also referred to as the ionization potential and is measured in volts. In chemistry, it often refers to one mole of a substance. (molar ionization energy or enthalpy) and is reported in kJ/mol. In atomic physics, the ionization energy is typically measured in the unit electron volt (eV). Atoms or molecules which are larger have lower ionization energy, while smaller molecules or atoms tend to have higher ionization energies.

The ionization energy is not the same for electrons of different molecular or ionic orbitals. More specifically, the nth ionization energy is the energy required to eliminate the nth electron after the first n-1 electrons are removed. It is perhaps a measure of the tendency of an ion or atom to surrender an electron or the strength of the electron binding.

Q8) Which element has highest ionization energy?

A8) It is because of the shielding effect that the ionization energy decreases from top to bottom within a group. From this trend, Cesium is said to have the lowest ionization energy and Fluorine is said to have the highest ionization energy (with the exception of Helium and Neon).

Q9) Differentiate between electronegativity and electron affinity?

A9) Electron affinity reflects the ability of an atom to accept an electron. It is the energy change which occurs when an electron is added to a gaseous atom. Atoms with strong effective nuclear charge have greater electron affinity. Electronegativity is the property of an atom which increases with its tendency to attract the electrons of a bond. when two bonded atoms that have similar electronegativity values as each other, they share electrons equally in a covalent bond. Usually, the electrons present in a chemical bond are more attracted to one atom (the more electronegative one) than to the other one. This results in a polar covalent bond. If the electronegativity values are very different, the electrons do not show any sharing. One atom essentially takes the bond electrons from the other atom, forming an ionic bond.

Q10) Explain the types of polarization.

A10) Electronic polarization, αe, describes the shift of the cloud of bound electrons with reference to the nucleus under an applied electric field. The atom scatter, and the centre of the atom’s negative charge eventually does not coincide with the position of the nucleus, resulting in an electric dipole moment. The electric dipole moment of each atom is described by p=αeE*

Distortion polarization αd (also often referred to as ionic polarization) relates to the distortion of the position of the nuclei by the applied field, resulting in stretching or compressing the bond length, depending on the relative orientation between the electric field and ionic bond. The molecule is stretched and bent by the applied field and accordingly changes its dipole moment. Nonpolar molecules may acquire an induced dipole moment in an electric field as a result due to the distortion of the electric field that causes disturbance in electronic distributions and nuclear positions.

Q11) How are Acids different from bases?

A11) Acids: Acids become less acidic when mixed with bases they taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) to red.

Bases: feel slippery, change litmus blue, and become less basic when mixed with acids.

While Boyle and others failed to explain why acids and bases behave the way they do, the first reliable definition of acids and bases was not be proposed until 200 years later.

Q12) Define the Arrhenius Principle?

A12) Arrhenius theory, introduced in 1887 by the Swedish scientist Svante Arrhenius, that acids are substances that dissociate in water to yield electrically charged atoms or molecules, called ions, one of which is a hydrogen ion (H+), and that bases ionize in water to yield hydroxide ions (OH−). It is now known that the hydrogen ion cannot exist alone in water solution; rather, it exists in a combined state with a water molecule, as the hydronium ion (H3O+).

Q13) Who introduced the HSAB theory? Why?

A13) Ralph Pearson introduced the Hard Soft [Lewis] Acid Base (HSAB) principle in the early 1960s and made an attempt to unite inorganic and organic reaction chemistry. The impact of the new concept was immediate, however with the HSAB principle had rather fallen apart. HSAB is widely used in chemistry for explaining stability of compounds, reaction mechanisms and pathways. It assigns the terms 'hard' or 'soft', and 'acid' or 'base' to chemical species. 'Hard' applies to species which are small, have high charge states (the charge criterion applies mainly to acids, to a lesser extent to bases), and are weakly polarizable. 'Soft' applies to species which are big, have low charge states and are strongly polarizable. The theory is used in contexts where a qualitative, rather than quantitative description would help in understanding the predominant factors which drive chemical properties and reactions

Q14) Explain hybridization with an e.g?

A14) Hybridization is a process where the atomic orbitals that have different shape and energy inter mingle together to form a set of same number of orbitals which have same shape and energy. happens when atomic orbitals mix to form new atomic orbitals. The new orbitals that is formed have the same total electron capacity as the old ones. The property and energy of the newly formed, hybridized orbitals are an 'average' of the original orbitals that are unhybridized orbitals.