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DAT > SNS - General Chemistry - Bonding And Chemical Interactions > Flashcards

SNS - General Chemistry - Bonding And Chemical Interactions Flashcards

1
Q

Octet rule

A

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

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2
Q

Molecular Geometry

A

Shape of the molecule that can be predicted using the Lewis dot diagram

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3
Q

Molecular Geometry Valence Shell Electron Pair Repulsion Theory

A

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

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4
Q

Molecular Geometry No Non-bonding Pairs AX

A

Linear, 180⁰

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5
Q

Molecular Geometry No Non-bonding Pairs AX2

A

Linear, 180⁰

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6
Q

Molecular Geometry No Non-bonding Pairs AX3

A

Trigonal planar, 120⁰

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7
Q

Molecular Geometry No Non-bonding Pairs AX4

A

Tetrahedral, 109.5⁰

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8
Q

Molecular Geometry No Non-bonding Pairs AX5

A

Trigonal bipyramidal, 120⁰, 90⁰

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9
Q

Molecular Geometry No Non-bonding Pairs AX6

A

Octahedral, 90⁰

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10
Q

Molecular Geometry Derived Structures

A

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

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11
Q

Molecular Geometry Derived Structures AE

A

Single atom

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12
Q

Molecular Geometry Derived Structures AXE

A

Linear Diatomic, 180⁰

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13
Q

Molecular Geometry Derived Structures AX2E

A

Bent, 120⁰

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14
Q

Molecular Geometry Derived Structures AX3E

A

Triangular Pyramidal, 109.5⁰

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15
Q

Molecular Geometry Derived Structures AX2E2

A

Bent, 109.5⁰

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16
Q

Molecular Geometry Derived Structures AX4E

A

Seesaw, 120⁰, 90⁰

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17
Q

Molecular Geometry Derived Structures AX3E2

A

T-shaped, 90⁰

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18
Q

Molecular Geometry Derived Structures AX2E3

A

Linear, 180⁰

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19
Q

Molecular Geometry Derived Structures AX5E

A

Square pyramidal, 90⁰

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20
Q

Molecular Geometry Derived Structures AX4E2

A

Square planar, 90⁰

21
Q

Molecular Geometry H2O

A

Oxygen has two non-bonding pairs Structure: derived structure AX2E2 Bent, 109.5⁰

22
Q

Molecular Geometry SO2

A

Sulphur has one non bonding pair Structure: derived structure AX2E Bent, 120⁰

23
Q

Molecular Geometry CO2

A

Carbon has no non-bonding pairs Structure: AX2 Linear, 180⁰

24
Q

Molecular Geometry BeCl2

A

Be has no non-bonding pairs Structure: AX2 Linear, 180⁰

25
Q

Molecular Geometry NH3

A

Nitrogen has one non-bonding pairs Structure: derived structure AX3E Trigonal pyramidal, 109.5⁰

26
Q

Molecular Geometry BF3

A

Boron has no non-bonding pairs Structure: AX3 Trigonal Planar, 120⁰

27
Q

Molecular Geometry BH3

A

Boron has no non-bonding pairs Structure: AX3 Trigonal Planar, 120⁰

28
Q

Molecular Geometry CH4

A

Carbon has no non-bonding pairs Structure: AX4 Tetrahedral, 109.5⁰

29
Q

Molecular Geometry CCl4

A

Carbon has no non-bonding pairs Structure: AX4 Tetrahedral, 109.5⁰

30
Q

Bond Types

Ionic

A

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

31
Q

Bond Types Metalliic

A
  1. 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
32
Q

Bond Types Covalent

A
  1. Solids, liquids, gases 2. Poor conductors of electricity and heat 3. Polar or non-polar
33
Q

Bond Types Covalent Polar

A
  1. Unequal sharing of bonding electron pairs 2. Electronegativity difference of 0.4-0.7 3. Dipoles
34
Q

Bond Types Covalent Non-Polar

A
  1. Equal sharing of bonding electron pairs 2. Same electronegativity 3. Diatomic molecules (O2, H2, Cl2)
35
Q

The Octet Rule

-Exceptions

A
  1. Hydrogen - can have only two valence electrons
  2. Lithium and Beryllium - Bond to attain two and four valence electrons respectively
  3. Boron - Bonds to attain six valence electrons
  4. Elements beyond the second row such as phosphorus and sulphur which can expant their octets by incorporating d-orbitals
36
Q

Bond Types

Ionic

Properties

A

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

37
Q

Bond Types

Covalent

Bond Length

A

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

38
Q

Bond Types

Covalent

Bond Energy

A

Energy required to separate two bonded atoms

Increases as the number of shared electrons increases

39
Q

Formal Charges

A

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

40
Q

Calculate the formal charge on the central N atom in [NH4]+

A

V = 5

In NH4, N has four bonds (ie 8 bonding electrons and no non-bonding electrons)

= 5 - (0.5 x 8) - 0

= +1

41
Q

Write Resonance Structures For [NCO]-

A
  1. C is the least electronegative atom so is placed at the centre - N C O
  2. 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
  3. Draw single bonds between C and the surrounding atoms: N:C:O
  4. 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>
  5. 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-
  6. 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
42
Q

Resonance

Guidelines

A
  1. A Lewis structure with small or no formal charges is preferred over a Lewis structure with large formal charges
  2. 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
43
Q

Dipole Moment

A

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

44
Q

Bond Types

Covalent Bond

Co-ordinate Bond

A

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

45
Q

Intermolecular Forces

A
  1. Hydrogen Bonds
  2. Dipole-Dipole interactions
  3. Dispersion forces
46
Q

Intermolecular Forces

Dipole-Dipole Interactions

A

Present in solid and liquid phases but not gas - molecules generally much farther apart

47
Q

Intermolecular Forces

Hydrogen Bonding

A

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

48
Q

Intermolecular Forces

Dispersion Forces

A

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

DAT (53 decks)

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