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CHEMISTRY-FY > S1-L8 &9: Covalent Bonding > Flashcards

S1-L8 &9: Covalent Bonding Flashcards

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1
Q

Types of bonds

A
  • Chemical Bond:
  • ->Strong bonds: Ionic bonding/ covalent bonding –> dative covalent bonding/ metallic bonding
  • ->Weaker bonds: H bonding/ dipole-dipole interactions/ Van Der Waal’s bonding
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2
Q

Define and describe Covalent Bonds

A
  • e- pairs shared by both participating atoms
  • ->shared e- pairs localised in definite space (bonding orbital) between nuclei of 2 atoms
  • directional bond
  • established between atoms of same/ different non-metallic elements
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3
Q

Explain how electrons are shared in covalent bonds

A
  • shared e- pairs spend most of time between two nuclei
  • bond formed by electrostatic attraction between oppositely charged nuclei AND shared e-‘s
  • atomic orbitals overlap to form molecular orbitals
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4
Q

Outline and briefly explain how molecular orbitals work (refer to figure 1)

A
  • molecular orbitals–>are linear combinations of atomic orbitals
  • when atoms interact to form molecules overall n. of orbitals must remain same
  • 2 atomic orbitals (1s)–>2 molecular orbitals
  • e-‘s fill molecular orbitals same way as fill atomic orbitals
  • ->into lowest energy first
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5
Q

Analyse figure 2 to see how a single covalent bond between Cl-Cl is formed

A

-figure 2 diagram

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6
Q

Outline Cl’s e- configuration and how may a bonding molecular orbital form (figure 3)

A
  • e- configuration: 1s2, 2s2, 2p6, 3s2, 3p5

- 2 p-orbitals could may overlap to form a bonding molecular orbital

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7
Q

How do double bonds (figure 4) and triple bonds (figure 5) work?

A
  • double bonds involve sharing of 2 e- pairs

- triple bond involves sharing of 3 e- pairs

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8
Q

The orbital picture

A

-figure 6

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9
Q

What are “Lewis Structures”?

A
  • method to describe covalent bonding in polyatomic molecules
  • show valence e-‘s as dots/ crosses
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10
Q

State the method to follow to draw Lewis structures

A

1-Draw e- configuration of atoms involved–> usually ground state
2- count valence e-‘s for all atoms to determine total n. of e-‘s in molecule
3- use e- pairs to form single bond between atoms (bonding pairs)
4- arrange remaining e- pairs around atoms (lone pairs) to satisfy octet rule (duet rule for H)
5-if run out of e-‘s use multiple bonds to complete octet
6-use valency knowledge to help

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11
Q

Lewis Periodic table showing outer shell (valence) electrons

A

-figure 7-analyse it

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12
Q

EXAMPLE: What is the Lewis Structure of CO2?

A
  • Carbon has 4 valence e-‘s–> valency is 4
  • oxygen has 6 valence e-‘s–>valency is 2
  • ->so first form 2 single bonds between C and two O’s
  • ->then double bond as one more bond between C and each O
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13
Q

Outline the factors which favour covalent bonding

A
  • high ionisation energies
  • equal/similar electronegativity
  • small atomic size
  • n. of valence e-‘s–>gaining/ losing 4 valence e-‘s v. hard so C forms covalent bonds
  • equal e- affinities
  • high nuclear charge
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14
Q

Figure 8 shows how average electronegativity and difference in electronegativity causes ionic/metallic or covalent bonding. Also some key information

A
  • elements right hand side of periodic table
  • same/ different types of atoms with high electronegativity
  • don’t ionise
  • equal attraction to e-‘s to complete octets
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15
Q

What are “non-polar bonds” and their effect?

A
  • between two atoms of equal/ v. similar electronegativity
  • e-‘s shared equally between the two atoms
  • ->on average bonding pair located half way between two atoms
  • bond non-polar–> no dipole
  • bonding e-‘s shared equally between the two atoms
  • -> no charge on atoms
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16
Q

Define and describe “polar bonds”

A
  • form between two atoms of significantly different electronegativity
  • ->there is separation of charge between one end and other AS greater e- density around more electronegative atom
  • bond polar and has dipole
  • affects reactivity of bond AND types of intermolecular forces between molecules
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17
Q

What is a “polar covalent bond”?

A
  • bonding e-‘s shared unequally between two atoms

- partial charges on atoms

18
Q

Summary of polar and non-polar covalent bonds

A
  • no electronegativity difference between two atoms leads to pure non-polar covalent bond
  • small electronegativity difference leads to polar covalent bond
  • large electronegativity difference leads to ionic bond
19
Q

Electronegativities of:

C-Cl/ Cl-Cl/ +Na-Cl-

A
  • C (3.0)-Cl (3.0)
  • C (2.5)-Cl (3.0)
  • +Na (0.9)-Cl- (3.0)
20
Q

Covalent compound features

A

-Figure 10

21
Q

Compare the following properties of covalent (A) and ionic (B) compounds:

1-State
2-M. point & B. point
3-Condictivity
4-Solubility

A

1- gases/ liquids/ solids (A)/ crystalline solids (B)
2-depends on size + intermolecular bonding (A)/ High (B)
3-mostly poor- depends on size & e-‘s delocalisation (A)/ good when molten (B)
4-depends on intermolecular bonding (A)/ many soluble in water BUT not in non-polar liquids (B)

22
Q

What are “Giant Covalent structures”?

A
  • group IV elements able to form up to 4 strong bonds between atoms
  • ->giant structures contain many atoms similar to ionic lattice
23
Q

Describe giant covalent structures (figure 11)

A
  • usually v. strong/ hard/ non-conductive/ insoluble in all solvents
  • graphite able to conduct in one direction due to it’s structure
  • silicon is semi-conductor
24
Q

How are covalent compounds named?

A
  • 2 word names
  • second element has -ide ending
  • each element has prefix indicating n. of atoms (not valence)
  • ->E.G: N2O4 is dinitrogen oxide
25
Q

What is the exception to the naming rules?

A

-drop mono for first element like CO2 (Carbon monoxide)

26
Q

Why is the first vowel often dropped?

A
  • to avoid combination of “ao” OR “oo”
  • ->E.G: CO is Carbon monoxide (monoxide)
  • P4O10 is tetra phosphorus decoxide
  • BUT SO2 is Sulphur dioxide (dioxide)
27
Q

Describe Dative Covalent bond formation

A
  • both bonding e-‘s provided by one of linked atoms (pr ions)
  • atom acting as donor must have lone pair of e-‘s
  • lone pair: pair of e-‘s in valence shell of atom which not involved in bonding
  • atom acting as acceptor should have vacant orbital to accept e- pair donated by donor
28
Q

How do the following dative covalent bonds form?:

1-Ammonium ion (figure 12)
2-Complex ions (figure 13)

A

1-formed when ammonia reacts with H+
2-form when transition metals ion + Al3+ dissolved in water
-e- configuration of Al3+: 1s2, 2s2, 2p6
–>all orbitals in 3rd shell vacant AND available to accept e-‘s

29
Q

How is dative bonding involved in biology? (figure 14-examples)

A
  • function of some biological dependent on binding a metal ion-containing cofactor–>porphyrins/ corrins
  • ->haemoglobin/ myoglobin/ Cytochromes P450/ Vitamin B12/ Chlorophyll/ Photodynamic therapay
30
Q

Outline how dative bonding is involved in medicine

A
  • function of some drugs depends on their ability to:
  • ->act as donor in dative covalent bonds-chelation therapy
  • ->act as donor for biological donors
  • figure 15 gives some examples
31
Q

Explain how “Valence Shell Electron Pair Repulsion Theory” can be used to predict the structures of simple covalently bonded molecules and ions

A
  • 1 central atom–>surrounding atoms all approx same size
  • 3D shape of simple molecule/ ion is that which keeps repulsive forces to minimum
  • ->e- pairs stay far apart as possible
  • can predict molecule’s shape by counting e- pairs
32
Q

Molecules with only bonding pairs:

N. of bonds–> shape–> Bond angles–> Example

N. of bonds: 2/3/4/5/6
(figure 16 shows examples)

A
  • 2 –> linear–> 180 –> BeCl2
  • 3 –> trigonal planar–> 120–> AlCl3
  • 4 –> tetrahedral–> 109.5 –> CH4
  • 5 –> trigonal bipyramidal–> 90 & 120 –> PCl5
  • 6 –> octahedral –> 90 –> SF6
33
Q

Molecules with lone pairs:

What is the order of e- repulsion?

A

-bond pair-bond pair< lone pair-bond pair< lone pair-lone pair

34
Q

How much do bond angles decrease by per lone pair?

A

-by 2 degrees

35
Q

Outline how domain geometry and molecular geometry compare as the n. of e- domains increase and the the n. of bonds and lone pairs composing of these domains differ

A

-Analyse Figure 17 and figure 18 tables which show the comparison

36
Q

What are the shapes and bond angles for the following molecules?:

1-Ammonia
2-Water 
3-ClF5
4-SF4
5-ClF3
6-XeF4
A 
1-Pyramdal (107 degrees) 
2-Bent linear (104.5 degrees) 
3-Square pyramidal (<90 degrees) 
4-See-saw (90 & 120 degrees) 
5-T-shaped (90 and 120 degrees) 
6-square planar (90 and 120 degrees)
37
Q

State the rules to follow to predict the shape of molecules using the central atom

A

1-Find n. of valence e-‘s
–>for ions add 1 valence e- for anions & remove 1 for cations
2-work out n. of bonding pairs AND lone pairs by drawing lewis structure of molecule
3-if no lone pairs–>molecular geometry is on 5 basic types
4-if there are lone pairs-electronic geometry still one of 5 basic types
–>but molecular geometry different-name based on atom arrangement

38
Q

EXAMPLE: What is the molecular geometry of NH2- ion? (refer to figure 20)

A
  • Nitrogen: 1s2, 2s2, 2p3 has 5 valence e-‘s
  • ->add 1 for (-) charge so 6e-‘s
  • ->2 bonding pairs/ 2 lone pairs
  • ->electronic (domain) geometry is distorted tetrahedral
  • ->molecular shape is angular (water)
39
Q

How would you predict the polarity of molecules?

A

1-are bonds polar? If not molecule not polar
2-determine molecule shape
3-do polar bonds point same direction OR oppose one another
–>opposite dipoles cancel

40
Q

Summary of covalent bond types:

1- N. of shared e-‘s
2-Electronegativity difference between atoms
3-N. of e-‘s shared by each atom

A

1-single–> 1 pair

  • double–> 2 pairs
  • triple–> 3 pairs

2-non-polar–> equal/ v. similar electronegativity
-polar–> unequal e-‘s sharing due to different electronegativities

3- normal covalent bond–> equal number from each atom
-dative covalent bond–> pair provided by one atom

41
Q

Summary of lecture

A
  • shapes of covalent molecules
  • ->predict via Valence Shell Electron Pair Repulsion
  • ->characteristic of bond angles
  • ->polarity of molecules

CHEMISTRY-FY (11 decks)

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