S1-L8 &9: Covalent Bonding Flashcards
Types of bonds
- 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
Define and describe Covalent Bonds
- 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
Explain how electrons are shared in covalent bonds
- 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
Outline and briefly explain how molecular orbitals work (refer to figure 1)
- 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
Analyse figure 2 to see how a single covalent bond between Cl-Cl is formed
-figure 2 diagram
Outline Cl’s e- configuration and how may a bonding molecular orbital form (figure 3)
- e- configuration: 1s2, 2s2, 2p6, 3s2, 3p5
- 2 p-orbitals could may overlap to form a bonding molecular orbital
How do double bonds (figure 4) and triple bonds (figure 5) work?
- double bonds involve sharing of 2 e- pairs
- triple bond involves sharing of 3 e- pairs
The orbital picture
-figure 6
What are “Lewis Structures”?
- method to describe covalent bonding in polyatomic molecules
- show valence e-‘s as dots/ crosses
State the method to follow to draw Lewis structures
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
Lewis Periodic table showing outer shell (valence) electrons
-figure 7-analyse it
EXAMPLE: What is the Lewis Structure of CO2?
- 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
Outline the factors which favour covalent bonding
- 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
Figure 8 shows how average electronegativity and difference in electronegativity causes ionic/metallic or covalent bonding. Also some key information
- 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
What are “non-polar bonds” and their effect?
- 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
Define and describe “polar bonds”
- 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
What is a “polar covalent bond”?
- bonding e-‘s shared unequally between two atoms
- partial charges on atoms
Summary of polar and non-polar covalent bonds
- 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
Electronegativities of:
C-Cl/ Cl-Cl/ +Na-Cl-
- C (3.0)-Cl (3.0)
- C (2.5)-Cl (3.0)
- +Na (0.9)-Cl- (3.0)
Covalent compound features
-Figure 10
Compare the following properties of covalent (A) and ionic (B) compounds:
1-State
2-M. point & B. point
3-Condictivity
4-Solubility
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)
What are “Giant Covalent structures”?
- group IV elements able to form up to 4 strong bonds between atoms
- ->giant structures contain many atoms similar to ionic lattice
Describe giant covalent structures (figure 11)
- 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
How are covalent compounds named?
- 2 word names
- second element has -ide ending
- each element has prefix indicating n. of atoms (not valence)
- ->E.G: N2O4 is dinitrogen oxide
What is the exception to the naming rules?
-drop mono for first element like CO2 (Carbon monoxide)
Why is the first vowel often dropped?
- 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)
Describe Dative Covalent bond formation
- 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
How do the following dative covalent bonds form?:
1-Ammonium ion (figure 12)
2-Complex ions (figure 13)
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
How is dative bonding involved in biology? (figure 14-examples)
- function of some biological dependent on binding a metal ion-containing cofactor–>porphyrins/ corrins
- ->haemoglobin/ myoglobin/ Cytochromes P450/ Vitamin B12/ Chlorophyll/ Photodynamic therapay
Outline how dative bonding is involved in medicine
- 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
Explain how “Valence Shell Electron Pair Repulsion Theory” can be used to predict the structures of simple covalently bonded molecules and ions
- 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
Molecules with only bonding pairs:
N. of bonds–> shape–> Bond angles–> Example
N. of bonds: 2/3/4/5/6
(figure 16 shows examples)
- 2 –> linear–> 180 –> BeCl2
- 3 –> trigonal planar–> 120–> AlCl3
- 4 –> tetrahedral–> 109.5 –> CH4
- 5 –> trigonal bipyramidal–> 90 & 120 –> PCl5
- 6 –> octahedral –> 90 –> SF6
Molecules with lone pairs:
What is the order of e- repulsion?
-bond pair-bond pair< lone pair-bond pair< lone pair-lone pair
How much do bond angles decrease by per lone pair?
-by 2 degrees
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
-Analyse Figure 17 and figure 18 tables which show the comparison
What are the shapes and bond angles for the following molecules?:
1-Ammonia 2-Water 3-ClF5 4-SF4 5-ClF3 6-XeF4
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)
State the rules to follow to predict the shape of molecules using the central atom
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
EXAMPLE: What is the molecular geometry of NH2- ion? (refer to figure 20)
- 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)
How would you predict the polarity of molecules?
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
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
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
Summary of lecture
- shapes of covalent molecules
- ->predict via Valence Shell Electron Pair Repulsion
- ->characteristic of bond angles
- ->polarity of molecules
