AIEEE-2012 Structure

AIEEE 2012

AIEEE 2012 – All India Engineering Entrance Examination is the gateway for admission into 1st semester of first year’s professional degree courses i.e. engineering, pharmacy and architecture offered through various AIEEE 2012 Colleges across the country. Student who successfully completes 12 year of school education in science discipline from any school board are eligible to write AIEEE examination. The minimum age is 17 years.

AIEEE 2012 shall be an Online Test as well as Paper Pencil Based Test. AIEEE 2012 Online Test shall be held for 1 lakh candidates only, according to CBSE AIEEE 2011 Notification. Candidates willing to appear for AIEEE 2012 online test must need to submit there application form online. However, it is said in AIEEE 2011 official notification that depending on the success of AIEEE 2011 Online Test, the AIEEE Board will offer Online Test for more number of aspirants for AIEEE 2012 and beyond. The AIEEE 2012 Online Test shall be available in AIEEE 2012 Paper I only.

The administering authority for AIEEE 2012 is AIEEE Board of Central Board of Secondary Education (CBSE), New Delhi. Student who appears in All India Engineering Entrance Examination (AIEEE) 2012 and finds place in the merit list prepared by AIEEE 2012 Board are called for counseling for final offer of admission. During counseling students are provided with option to choose from range of study programs of interest.

AIEEE 2012 Exam Dates
The AIEEE 2012 shall be held sometime in April / May 2012.
In the year 2011, the AIEEE Board has rescheduled AIEEE 2011 Exam Dates From April (24.04.2011) to May (01.05.2011) (Sunday) because of Easter. AIEEE 2011 has 2 papers : Paper 1 meant for aspirants interested in B.E / B.Tech Courses. Candidates willing to pursue B.Arch and related courses shall need to appear in both Paper 1 and Paper 2.

AIEEE 2012 Exam Pattern

AIEEE 2012 Paper I: This paper shall include questions from physics, chemistry and mathematics.

AIEEE 2012 Paper II: This paper exclusively meant for architecture and design courses.

Important: AIEEE 2012 Paper I shall be available in 2 modules: AIEEE 2012 Online Test and AIEEE 2012 Paper Pencil based test. The AIEEE 2012 Online Test shall be available for candidates interested in B.E & B.Tech courses only.

Preparing for AIEEE 2012

AIEEE 2012 is a competitive examination. The success in AIEEE 2012 requires systematic approach; one has to be regular in his / her preparation. There is no shortage of study material for AIEEE 2012; you can find huge number of books in different bookstalls in your locality. When choosing the study material for AIEEE 2012; do your analysis, be conscious, buy quality books which can offer you good number of solved and unsolved questions on every topic which is offered in AIEEE 2012 test. When you buy a book you must compare its content with the syllabus of AIEEE 2012.

The ideal time for preparation of All India Engineering Entrance Examination (AIEEE) is after passing of class 10th board examination. Though AIEEE 2012 follows syllabus pattern from class 11th and 12th of CBSE board school syllabus but almost 85% syllabus matches with any state board syllabus of class 11th & 12th.

Read More

Atomic Structure

Atomic Structure

The concept of the atom was created by early Greek philosophers who believed that all matter was composed of indivisible particles. They called these particles atomos, meaning “uncuttable.” It wasn't until the early nineteenth century that John Dalton formulated a theory based on scientific investigation that characterized the nature of atoms. Further discoveries in the nineteenth and twentieth centuries led to the knowledge that atoms possess an internal structure of smaller subatomic particles.

Subatomic particles. The major subatomic particles were found to be protons, electrons, and neutrons. Protons are positively charged particles that have weight.Electrons are negatively charged particles of little weight, while neutrons are just slightly heavier than protons but have no charge. Investigations revealed that protons and neutrons are located in the central core, or nucleus, of the atom, while electrons exist outside of the nucleus in areas of high probability called orbits, or shells. Orbits are further divided into more precise regions of electron probability called orbitals, or subshells.

Niels Bohr proposed the concept of the solar-system atom, in which the nucleus of the atom is like the sun and the electrons are like the planets, revolving in circular orbits. The farther an orbit is from the nucleus, the larger the orbit becomes and the more electrons it can hold.

Because all atoms are electrically neutral, the number of protons and electrons must be equal. Neutrons add weight but no charge to an atom, so additional neutrons do not change an element but merely convert it to one of its isotopic forms. Theatomic number ( Z) of an atom is equal to the number of protons in the nucleus or the number of electrons in its orbits. The atomic mass ( A) is equal to the sum of the protons and neutrons in the atom. (A proton and neutron each have a mass of 1 atomic mass unit, while an electron has virtually no mass.)

Atoms are capable of both losing and gaining electrons to achieve a stable state. If an atom loses one or more electrons, it becomes a positively charged ion called acation. If an atom gains one or more electrons, it becomes a negatively charged ion called an anion. The charge on an ion is equal to the number of electrons lost or gained.

Orbits and orbitals. Electrons fill orbits in an organized fashion based on energy factors. The order of electron fill-in, called the aufbau buildup, is 1 s, 2 s, 2 p, 3 s, 3 p, 4 s, …, where the numerals represent the principal quantum number of the orbit, and the lowercase letters represent the orbitals within a given orbit. The numbering begins with 1 for the orbit closest to the nucleus of the atom. The lower the orbit number, the smaller the orbit size and fewer electrons the orbit can hold.

The first principal orbit is large enough to hold just two electrons in an s orbital. The second principal orbit is large enough to contain one s and three p orbitals, while the third principal orbit, which is larger still, contains an s orbital, three p orbitals, and five d orbitals. When electrons are added to equivalent orbitals, which are orbitals of the same principal level and type, one electron must occupy each equivalent orbital before any of these orbitals can contain two electrons. Thus carbon, Z = 6, has six electrons distributed in these orbitals:

The orbitals can also be shown in the following fashion. In this diagram, the arrows represent electrons. Notice that single electrons are filling the 2 p orbitals one at a time and not pairing first in 2 p_x .

For two electrons to occupy the same orbital, they must have opposite spins, or paired spins, which generate orbital stability by creating opposite magnetic poles.Between equivalent orbitals, the spins of the electrons must be parallel, that is, spinning in the same direction, for the orbitals to be stable. Parallel spins create the same magnetic pole, causing repulsion between the orbitals. This repulsion gives the orbitals maximum separation and the greatest stability.

Orbitals within a given orbit have different shapes and sizes. The s orbitals are spherical, while the p orbitals are hourglass shaped. The s orbital is smaller than thep orbital.

Read More

Three‐Dimensional Shapes of Molecules

Three‐Dimensional Shapes of Molecules

The overall shape of an organic molecule is fixed by the shape of the central carbon atoms, which compose the backbone of the molecule. The shape of this backbone is determined by the types of hybrid orbitals making up the bonds between the central carbon atoms. If the central carbon atoms are sp³ hybridized, the molecule will possess a tetrahedral shape. Central carbon atoms that are sp² hybridized lead to trigonal-planar shapes, while sp hybridization produces linear molecules. Three-dimensional representations of methane ( sp³ hybridization), ethene ( sp¹hybridization), and ethyne ( sp hybridization) molecules are shown in Figure 1 . 

Figure 1

Read More

Hybridization of Atomic Orbitals

Hybridization of Atomic Orbitals

Physical studies of the simplest organic compound, methane (CH₄), have shown the following:

• all of the carbon-hydrogen bond lengths are equal


• all of the hydrogen-carbon-hydrogen bond angles are equal


• all of the bond angles are approximately 110°


• all of the bonds are covalent

The ground state, or unexcited state, of the carbon atom ( Z = 6) has the following electron configuration.

Covalent bonds are formed by the sharing of electrons, so ground-state carbon cannot bond because it has only two half-filled orbitals available for bond formation. Adding energy to the system promotes a 2 s electron to a 2 p orbital, with the resulting generation of an excited state. The excited state has four half-filled orbitals, each capable of forming a covalent bond. However, these bonds would not all be of the same length because atomic 5 orbitals are shorter than atomic porbitals.

To achieve equal bond lengths, all the orbitals would have to be the same type. The creation of identical orbitals occurs in nature by a hybridization process.Hybridization is an internal linear combination of atomic orbitals, in which the wave functions of the atomic s and p orbitals are added together to generate new hybrid wave functions. When four atomic orbitals are added together, four hybrid orbitals form. Each of these hybrid orbitals has one part s character and three parts pcharacter and, therefore, are called sp³ hybrid orbitals.

In the hybridization process, all bond lengths become equal. Bond angles can be explained by the valence-shell electron-pair repulsion theory (VSEPR theory).According to this theory, electron pairs repel each other; therefore, the electron pairs that are in bonds or in lone pairs in orbitals around an atom generally separate from each other as much as possible. Thus, for methane, with four single bonds around a single carbon, the maximum angle of repulsion is the tetra-hedral angle, which is 109°28″, or approximately 110°.

In a similar fashion, the atomic orbitals of carbon can hybridize to form sp² hybrid orbitals. In this case, the atomic orbitals that undergo linear combination are one sand two p orbitals. This combination leads to the generation of three equivalent sp²hybrid orbitals. The third p orbital remains an unhybridized atomic orbital. Because the three hybrid orbitals lie in one plane, the VSEPR theory predicts that the orbitals are separated by 120° angles. The unhybridized atomic p orbital lies at a 90° angle to the plane. This configuration allows for the maximum separation of all orbitals.

Last, the atomic orbitals of carbon can hybridize by the linear combination of one sand one p orbital. This process forms two equivalent sp hybrid orbitals. The remaining two atomic p orbitals remain unhybridized. Because the two sp hybrid orbitals are in a plane, they must be separated by 180°. The atomic p orbitals exist at right angles to each other, one in the plane of the hybridized orbitals and the other at a right angle to the plane.

The type of hybrid orbital in any given carbon compound can be easily predicted with the hybrid orbital number rule.

A hybrid orbital number of 2 indicates sp hybridization, a value of 3 indicates sp²hybridization, and a value of 4 indicates sp³ hybridization. For example, in ethene (C₂H₄), the hybrid orbital number for the carbon atoms is 3, indicating sp²hybridization.

All the carbon-hydrogen bonds are σ, while one bond in the double bond is σ and the other is π.

Thus, the carbons have sp² hybrid orbitals.

Using the hybrid orbital number rule, it can be seen that the methylcarbocation contains sp² hybridization, while the methylcar-banion is sp³ hybridized.

Read More