Unit 01

General Chemistry

Q-1What is effective nuclear charge?

The effective nuclear charge is the net charge an electron experiences in an atom with multiple electrons. The effective nuclear charge may be approximated by the equation:

Zeff = Z – S

Where Z is the atomic number and S is the number of shielding electrons.
Higher energy electrons can have other lower energy electrons between the electron and the nucleus, effectively lowering the positive charge experienced by the high energy electron.

The shielding effect is the name given to the balance between the attraction between valence electrons and protons and the repulsion between valence and inner electrons. The shielding effect explains the trend in atomic size on the periodic table and also why valence electrons are readily removed from an atom.

Q-2Draw the periodic table.

Q-3Explain shape of s orbital.

Shape of s orbital:

The s orbital is spherical having the nucleus at its center which in two dimensions and can be seen as a circle. The s-orbitals are spherically symmetric having the probability of finding the electron at a given distance equal in all the directions. The size of the s orbital is also found to increase with the increase in the value of the principal quantum number, thus, 4s > 3s> 2s > 1s.

Q-4Explain shape of p orbital.

Each p orbital consists of two sections that is known as lobes which lie on either side of the plane passing through the center that is nucleus. The three p orbitals differ in the way lobes are oriented whereas they are identical in terms of size shape and energy. The lobes lie along one of the x, y or z-axes, these three orbitals are given the designations 2px, 2py, and 2pz. Thus, we can say that there are three p orbitals whose axes are mutually perpendicular. Similar to s orbitals, size, and energy of p orbitals increases with an increase in the principal quantum number (4p > 3p > 2p).

Q-5What are atomic radii?

Atomic radius is the radius of spherical atoms. Nonbonding atoms have a larger, more undefined radius, so when atomic radius is discussed as a periodic trend, what's usually meant is bonding atomic radius. These are the radius of atoms that are chemically bonded to one another. So, if the bond between two Cl atoms in Cl2 is 1.99 angstroms, we report chlorine's bonding atomic radius as about 0.99 angstroms. Further these values to estimate bond lengths between different elements in molecules.

Q-6What are ionic radii?

Ionic radii are the radii of ions of elements. These distances are based on distances between ions in ionic compounds. Cations, or positively charged ions, are smaller than their "parent" atoms. This is because cations are formed when those outermost orbitals are vacated of electrons. This also decreases electron-electron repulsions. Therefore, the resulting ions are smaller as there are not as many occupied orbitals and the effective nuclear charge affecting the remaining electrons increases, pulling electrons in more closely.

Q-7Explain Ionization energies.

Ionization energy is the quantity of energy that an isolated, gaseous atom in the ground electronic state must absorb to discharge an electron, resulting in a cation.

H(g)→H+(g)+e−(1)(1)H(g)→H+(g)+e−

This energy is usually expressed in kJ/mol.

When considering an initially neutral atom, expelling the first electron will require less energy than expelling the second, the second will require less energy than the third, and so on. Each successive electron requires more energy to be released. This is because after the first electron is lost, the overall charge of the atom becomes positive, and the negative forces of the electron will be attracted to the positive charge of the newly formed ion. The more electrons that are lost, the more positive this ion will be, the harder it is to separate the electrons from the atom.

Q-8What is electron affinity?

Electron affinity is defined as the amount of energy released on the addition of electron in neutral atom to form an anion. The electron affinity is the potential energy change of the atom when an electron is added to a neutral gaseous atom to form a negative ion. So, the more negative the electron affinity the more favorable the electron addition process is. Not all elements form stable negative ions in which case the electron affinity is zero or even positive.

Electron affinity increases going left to right across a period. The overall trend across a period occurs because of increased nuclear attraction.

Going down the group the electron affinity should decrease since the electron is being added increasingly further away from the atom. Less tightly bound and therefore closer in energy to a free electron.

Electron affinity=1/Atomic Size

1. Atomic size: If the atomic size is small, then there will be greater electron gain enthalpy because the effective nuclear forces will be greater in the smaller atoms and the electrons will be held firmly.

2. Nuclear charge: The greater the nuclear charge more will be the value for electron gain enthalpy because an increase in nuclear charge will increase the effective nuclear force on valence electrons.

Q-9Explain oxidation states.

Some elements in the periodic table have only one oxidation number or two oxidation numbers. But some have lot of oxidation numbers. Oxidation number of elements in a compound can be positive or negative or may be zero.

In sodium compounds, sodium only forms +1 oxidation number.

But some types of atoms such as chlorine form various oxidation numbers like -1, 0, +1, +3, +5, +7 oxidation numbers in compounds.

Rules for the determination of Oxidation States:

• The oxidation state of an uncombined element is zero. This applies regardless of the structure of the element: Xe, Cl2, and large structures of carbon or silicon each have an oxidation state of zero.
• The sum of the oxidation states of all the atoms or ions in a neutral compound is zero.
• The sum of the oxidation states of all the atoms in an ion is equal to the charge on the ion.
• The more electronegative element in a substance is assigned a negative oxidation state. The less electronegative element is assigned a positive oxidation state. Remember that electronegativity is greatest at the top-right of the periodic table and decreases toward the bottom-left.

Q-10Explain Hard acid and bases.

Hard acids consist of small highly charged cations and molecules in which a high positive charge can be induced on the central atom.

Examples of Hard Acids: H+, Li+, K+, Ca2+, Al3+, Sn4+, BF3, BCl3, CO2, RCO+, SO3, RMgX, VO2+, AlCl3

Hard bases are highly electronegative and of low polarizability.

Examples of Hard Bases: F-, OH-, NH3, N2H4, ROH, H2O, SO42-, PO43-

Hard bases react more readily to form stable compounds and complexes with hard acids.

Q- 11 Enlist the characteristics of HSAB.

Hard acid soft acid:

Hard base soft base:

Q-12Explain molecular geometry of H3O+.

H3O + is pyramidal in shape.

In H3O+, the central atom (oxygen) has three bond pairs and one lone pair and hence, it is sp3 hybridized.

You must know that ideal geometry for a sp3 hybridized central atom molecule is Tetrahedral. So, the molecule should have Tetrahedral geometry in ideal condition ( a condition of all electron pairs on central atom being bond pairs only ).

But this is not an ideal case. Here, three electron pairs are bond pairs while one electron pair is lone pair creating distortion in molecule leading to a pyramidal shape of H3O+.

Q-13Explain Kossel postulate of ionic bonding.

1. In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases. Therefore, one or small number of electrons are easily gained and transferred to attain the stable noble gas configuration.