SNS - General Chemistry - Bonding And Chemical Interactions Flashcards
Octet rule
States that atoms lose, gain and share electrons until they are surrounded by eight valence electrons. The noble gases (except helium) have eight valence electrons and are therefore quite stable
Molecular Geometry
Shape of the molecule that can be predicted using the Lewis dot diagram
Molecular Geometry Valence Shell Electron Pair Repulsion Theory
States that the geometry of the electrons surrounding an atom will be an arrangement such that the electrons are as far away from each other as possible
Molecular Geometry No Non-bonding Pairs AX
Linear, 180⁰
Molecular Geometry No Non-bonding Pairs AX2
Linear, 180⁰
Molecular Geometry No Non-bonding Pairs AX3
Trigonal planar, 120⁰
Molecular Geometry No Non-bonding Pairs AX4
Tetrahedral, 109.5⁰
Molecular Geometry No Non-bonding Pairs AX5
Trigonal bipyramidal, 120⁰, 90⁰
Molecular Geometry No Non-bonding Pairs AX6
Octahedral, 90⁰
Molecular Geometry Derived Structures
Have non-bonding pairs and are variations of the basic shapes without non-bonding pairs Non-bonding pairs take up slightly more space than do atoms. The trigonal planar structure AX3 is bent with a non-bonding electron pair in place of an atom
Molecular Geometry Derived Structures AE
Single atom
Molecular Geometry Derived Structures AXE
Linear Diatomic, 180⁰
Molecular Geometry Derived Structures AX2E
Bent, 120⁰
Molecular Geometry Derived Structures AX3E
Triangular Pyramidal, 109.5⁰
Molecular Geometry Derived Structures AX2E2
Bent, 109.5⁰
Molecular Geometry Derived Structures AX4E
Seesaw, 120⁰, 90⁰
Molecular Geometry Derived Structures AX3E2
T-shaped, 90⁰
Molecular Geometry Derived Structures AX2E3
Linear, 180⁰
Molecular Geometry Derived Structures AX5E
Square pyramidal, 90⁰
Molecular Geometry Derived Structures AX4E2
Square planar, 90⁰
Molecular Geometry H2O
Oxygen has two non-bonding pairs Structure: derived structure AX2E2 Bent, 109.5⁰
Molecular Geometry SO2
Sulphur has one non bonding pair Structure: derived structure AX2E Bent, 120⁰
Molecular Geometry CO2
Carbon has no non-bonding pairs Structure: AX2 Linear, 180⁰
Molecular Geometry BeCl2
Be has no non-bonding pairs Structure: AX2 Linear, 180⁰
Molecular Geometry NH3
Nitrogen has one non-bonding pairs Structure: derived structure AX3E Trigonal pyramidal, 109.5⁰
Molecular Geometry BF3
Boron has no non-bonding pairs Structure: AX3 Trigonal Planar, 120⁰
Molecular Geometry BH3
Boron has no non-bonding pairs Structure: AX3 Trigonal Planar, 120⁰
Molecular Geometry CH4
Carbon has no non-bonding pairs Structure: AX4 Tetrahedral, 109.5⁰
Molecular Geometry CCl4
Carbon has no non-bonding pairs Structure: AX4 Tetrahedral, 109.5⁰
Bond Types
Ionic
When two atoms with a large difference in electronegativity (>1.7) react, there is a complete transfer of electrons from the less electronegative to the more electronegative atom (becoming a cation and anion respectively)
In general, the elements of groups II and II (low electronegativities) bond ionically to elements of group VII (high electronegativities). The resultant electrostatic force is called an ionic bond
Bond Types Metalliic
- Solids 2. Positively charged ions in a sea of electrons 3. Variable hardness 4. Variable melting points 5. Good conductors of electricity 6. Ductile and malleable with high deformity
Bond Types Covalent
- Solids, liquids, gases 2. Poor conductors of electricity and heat 3. Polar or non-polar
Bond Types Covalent Polar
- Unequal sharing of bonding electron pairs 2. Electronegativity difference of 0.4-0.7 3. Dipoles
Bond Types Covalent Non-Polar
- Equal sharing of bonding electron pairs 2. Same electronegativity 3. Diatomic molecules (O2, H2, Cl2)
The Octet Rule
-Exceptions
- Hydrogen - can have only two valence electrons
- Lithium and Beryllium - Bond to attain two and four valence electrons respectively
- Boron - Bonds to attain six valence electrons
- Elements beyond the second row such as phosphorus and sulphur which can expant their octets by incorporating d-orbitals
Bond Types
Ionic
Properties
High melting and boiling points due to strong electrostatic forces
Can conduct electrivity in the liquid and aqueous, but not solid, states due to crystal lattice structure in solid state
Bond Types
Covalent
Bond Length
Average distance between the two nuclei of atoms involved in the bond
As the number of shared electron pairs increases, the two atoms are pulled closer together and bond length decreases
Bond Types
Covalent
Bond Energy
Energy required to separate two bonded atoms
Increases as the number of shared electrons increases
Formal Charges
Difference between the number of electrons assigned to an atom in a Lewis model and the number of valence electrons in the free atom
= V - 1/2 Nbonding - Nnon-bonding
Where V is valence electrons in the free atom, Nbonding is the number of bonding electrons and Nnon-bonding is the number of non-bonding electrons
Calculate the formal charge on the central N atom in [NH4]+
V = 5
In NH4, N has four bonds (ie 8 bonding electrons and no non-bonding electrons)
= 5 - (0.5 x 8) - 0
= +1
Write Resonance Structures For [NCO]-
- C is the least electronegative atom so is placed at the centre - N C O
- N has 5 valence electrons, C four and O six. The species itself has a charge of -1. Total valence electrons = 5 + 4 + 6 + 1 = 16
- Draw single bonds between C and the surrounding atoms: N:C:O
- Complete the octets of N and O with the remaining 16 - 4 = 12 electrons: <strong><span>..</span></strong> <strong>..</strong> ** :N:** C**:O:** .. <strong> .. </strong>
- The C octet is incomplete. There ae three ways in which double and triple bonds can be formed to complete the octet: N-C≡O, N=C=O, N≡C-O. These are the three resonance structures for NCO-
- Assign formal charges to each atom in each resonance structure. The most stable is N≡C-O since the negative formal charge is on the most electronegative atom O
Resonance
Guidelines
- A Lewis structure with small or no formal charges is preferred over a Lewis structure with large formal charges
- A Lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which these are placed on less electronegative atoms
Dipole Moment
µ = q x r
Where q = the charge amgnitude and r = the distance between the tro partial charges
Denoted by an arrow pointing from the positive to the negative charge
Bond Types
Covalent Bond
Co-ordinate Bond
Bond in which both electrons are contributed by one of the atoms in the molecule: the shared electron pair comes from a lone pair. Once formed it is indistinguisable from any other covalent bond
Intermolecular Forces
- Hydrogen Bonds
- Dipole-Dipole interactions
- Dispersion forces
Intermolecular Forces
Dipole-Dipole Interactions
Present in solid and liquid phases but not gas - molecules generally much farther apart
Intermolecular Forces
Hydrogen Bonding
Specific, unusually strong form of dipole-dipole interaction which may be either inter or intramolecular
When hydrogen is bound to a highly electronegative atom such as fluorine, oxygen or nitrogen, the hydrogen atom carries very little of the electron density of the covalent bond. The positively charged hydrogen interacts with the partial negative charge on nearby electronegative atoms
Intermolecular Forces
Dispersion Forces
The bonding electrons in covalent bonds will, at any particular point in time, be located randomly throughout the orbital. This permits unequal sharing of electrons causing rapid polarisation and counterpolarisation of the electron cloud and the formation of short-lived dipoles. These can interact with the electron clouds of neighbouring molecules inducing the formation of more dipoles
Generally weaker than the other intermolecular forces and don’t extend over long distances. Strength depends on how easily electrons within a given substance can move (ie be polarised). Large molecules in which electrons are far from the nucleus are relatively easy to polarise and therefore possess greater dispersion forces
