Molecular Geometry and Chemical bonding Theory

Question 1

Molecular Geometry

Answer 1

• the general shape of a molecule, as determined by the relative positions of the atomic nuclei
• 3D shape plays an important part in determining the molecules chemistry
• a change in a single site in a molecule can make a difference in whether a particular reaction occurs.

Question 2

The Valence Shell Electron Pair Repulsion (VSEPR)

Answer 2

• Predicts the shape of molecules and ions by assuming that the valence shell electron pairs are arranged about each atom so that the electron pairs are kept as far away as possible from one another, to minimize repulsion.
• Minimizes the energy of repulsion and represents the lowest energyu configuration of the molecule
• The number of valence electron pairs and their nature make the terminal atoms adopt specific orientations around the central atom.

Question 3

How are linear arrangement configured?

Answer 3

• 2 electron pairs
• bond angle is 180º

Question 4

How are trigonal planer configured?

Answer 4

• 3 pairs on the central atom
• electron pairs are farthest apart when they lie in the same plan on point to the corners of an equlateral triangle
• angle is 120º

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Question 5

How are tetrahedral configured?

Answer 5

• 4 electron pairs around the central atom
• Angle should be 90º , but molecules are 3D so it is usually 109.5º
• electron pairs are furthest apart when they extend towards the corners of a regular tetrahedron

Question 6

How are trigonal bipyramidal configured?

Answer 6

• 5 electron pairs
• 120º

Question 7

How are octahedral configured?

Answer 7

• 6 elecron pairs
• 90º

Question 8

Electron group

Answer 8

• Any collection of valence electrons, localized in a region around a central atom, which repels other groups of valence electrons.
• Electon groups can be
  
  • Single unpaired electron
  • lone pair electrons
  • One bonding pair of electrons in a single covalent bond
  • Two bonding pair of electons in a double covalent bond.
  • Three bonding pairs of electrons in a triple covalent bond.

Question 9

What are the 5 common types of electron group configurations?

Answer 9

• There are usually 5 geometry types
• The mutual repulsion among the electron groupd causes an orientaion of the group which is called electron group geometry.
  
  • 2 electron group: Linear
  • 3 electron group: trigonal planar
  • 4 electron group: tetrahedral
  • 5 electron group: trigonal bipyramiidal
  • 6 electron group: octahedral

Question 10

Bond angles and the effects of lone pairs

Answer 10

• Bond angles for lone pair of electons spread out more than bonding pairs of electrons
• Lone pairs take up more space and so reduce the angle of the atoms
• reasons
  
  • bonding pairs of electrons are attracted to two nuclei simultaneously, the charge cloud associated with them is pulled into a compact shape
  • lone pairs of electrons are associated with one nucleus. repulsion of one lone pair is greater than the repulsion between bonding pairs.

Question 11

Dipole Moment

Answer 11

• Quantitative measure of charge seperation in a molecule
• is a product of the magnitude of the charge and distance that seperate the centers of positive and negative charge.
• debyes (D):
  
  • traditional measurement
  • 3.34 x 10^-30 C.m
• SI unit: coulomb-meters (C.m)
• Use (+–>) to show the bond dipole, the dipole moments arrow points from the positive partial charge toward the negative partial charge.

Question 12

Bond Dipole

Answer 12

a seperation of positive and negative charge centers in an individual bond.

Question 13

Molecular dipole

Answer 13

Charge seperation in the molecule as a sholw, considering all the bonds.

Question 14

Non-Polar

Answer 14

Vectors may have equal magnitude but opposite direction caused cancellation of effects.

Question 15

Symmetrical structures

Answer 15

may cause the bond polarities to cancel leading to a zero dipole movement.

Question 16

Valance Bond Theory

Answer 16

• Explains how bonding occurs based on quantom mechanics
• a covalent bond is formed by pairing of two electrons with opposing spins in the region of overlap of atomic orbitals. this overlap has a high electron charge density
• The greater the amount of orbital overlap, the stronger the bond
• Optimal orbital overlap is needed for each bond for maximum bond strength at a particular internuclear distance.
• if two atoms are too close, then repulsion of atomic nuclei compromised the attraction between electrons and nuclei and the bond becomes weak and unstable.

Question 17

Hybrid Orbitals

Answer 17

• Orbitals used to describe bonding. these are obtained by combing atomic orbitals of the isolated atom.
• the orbitals that contain an unpaired electron are those that will overlap to form bonds.

Question 18

Steps in predicting a probable hybridization scheme.

Answer 18

1. Write a plausible lewis structure of the molecule or ion
2. Use the VSEPR method to predict the electon group geometry of the central atom.
3. Selecting the hybridization scheme that corresponds to the VSEPR prediction.
4. Put the valence electrons of the central atom in hybrid orbitals. if possible singly.
5. Finally overlap the singly occupied orbitals of the othe atoms with the singly occupied hybrid orbitals of the central atom.

Question 19

Description of Multiple Bonding

Answer 19

One hybrid orbital is needed for each bond (whether a single or multiple bond) and for each lone pair

Question 20

Sigma Bond (σ)

Answer 20

• head-to-head or end-to-end overlap of orbitals regardless of orbital type
• all single covalent bonds are σ bonds

Question 21

Pi bonds (π)

Answer 21

• side-by-side parallel overlap of atomic orbitals
• electron distribution is above and below the bond axis
• a pi bond is formed only if there is a also a sigma bond.

Question 22

Double bond

Answer 22

one sigma and one pi bond

Question 23

Triple bond

Answer 23

• one sigma and two pi bonds
• each carbon has two regions of high electron density - so sp hybrididized and linear geometry.

Question 24

Isomers

Answer 24

• same molecular formula but different arrangement of atoms
• Double bond:
  
  • less flexible, rotation is hard
• Multiple bond: more reactive
  
  • more reactive and stronger than single bonds
  • a lot of energy is required to break up a pi bond os cis- and trans- isomers are not easily interconverted.

Question 25

Molecular orbital theory

Answer 25

• combination of atomic orbitals on different atoms forms molecular orbitals so that electrons in them belong to the molecule as a whole
• Two atomic orbitals combine to form two molecular orbitals. So combining n atomic orbitals give n molecular orbitals.
• Electrons seek the lowest energy molecular orbital that is available.
• A maximum of two electrons can be present in a molecular orbital(Pauli exclusion principle).
• Electrons enter molecular orbitals of identical energies singly with parallel spins before they pair up(Hund’s rule)

Question 26

Molecular Orbital

Answer 26

• a wave function whose square gives the probablity of finding an electron within a given region of space in a molecule.
• are mathematical equations that describe the regions in a molecule where there is a high probability of finding electrons.
• Like atomic orbitals, molecular orbitals have specific energy levels and can be occupied by a maximum of two electrons with opposite spin.

Question 27

Additive combination or bonding molecular orbital (σ)

Answer 27

• molecular orbitals that are concentrated in regions between nuclei
• is lower in energy than the separated atomic orbitals
• any electrons it contains spend most of their time in the region between the two nuclei, helping to bond the atoms together.
• it has a high electron probablity or electron charge density between the bonded atoms.

Question 28

Subtractive combnation or antibonding molecular orbital (σ*) sigma star

Answer 28

• Molecular orbitals with zero value in the region between two nuclei
• is higher in energy than isolated atomic orbitals
• an electrons that is contains cannot occupy the central region between nucleai and so cannot conribute to bonding
• high electron probability away from the region between the bonded atoms.

Question 29

Bond order

Answer 29

• Is one-half the difference between the number of electrons in the bonding molecular orbitals and the number in the antibonding orbitals.
• The number of bonds that exists between two atoms.
• The number of electron pairs shared between atoms.
• A way to judge stability. E.g. bond order for He₂ is 0, that is why it is unstable.
• Fractions are possible.

Question 30

Factors that determine orbital intraction

Answer 30

1. Energy difference between the interacting orbitals, should be equal or close
2. Magnitude or the overlap must be large

Question 31

Paramagnetic

Answer 31

• attracted to a magnet
• unpaired electrons are attracted by magnetic fields
• More unpaired electrons, the stronger the paramagnetic attraction

Question 32

Diamagnetic

Answer 32

• repelled by magnetic fields
• paired electrons with opposite spin