Bonding and structure part 2 Flashcards
Define electronegativity (2)
The ability of an atom to attract the bonding electrons towards itself to form a covalent bond
Describe and explain how a covalent bond is formed in Br2 (3)
- Both same atom so both have some pulling power (electronegativity)
- So the bonding electrons are shared equally between the bonding atoms
- Therefore the bonds in bromine are non-polar
Describe and explain how a covalent bond is formed in HCl (4)
- Cl is far more electronegative than H
- The bonding electrons are held closer to the Cl atom
- The electron cloud is denser over the Cl atom so there is a permanent dipole across the bond
- Therefore HCl has polar molecules
How would the bonding in HI be represented?
H———————————————————I
delta positive delta negative
xo
(bond shown slightly closer to I due to a difference in electronegativity)
On what scale are the electronegativities of different atoms shown?
The Pauling Scale
Does CH4 have polar molecules? Explain why. (2)
No- carbon and hydrogen have very similar electronegativities so the electron pair isn’t held closer to any of the atoms.
Compare the charges in NaBr and HBr (2)
NaBr=ionic, so full charges
HBr=polar covalent, so only partial charges
Which element is the most electronegative? Explain why (4)
- Fluorine
- Has the least number of shells in group 7, so lower shielding effect and smaller atomic radius
- So stronger nuclear attraction with oncoming electron
- Has the highest nuclear charge in period 2, so strongest nuclear attraction between nucleus and oncoming electrons
So fluorine gains an electron most easily
Non-polar or polar:
CF4
explain why (3)
Non-polar- although there is a difference in electronegativity, the bond pairs are repelled equally meaning symmetry cancels the dipoles out- no net dipole across the molecule.
Non-polar or polar:
NH3
explain why (3)
Polar-there is a difference in electronegativity, and a lone pair of electrons around the central atom making the molecule unsymmetrical, so the dipoles do not cancel
Non-polar or polar:
CO2
explain why (3)
Non-polar- although there is a significant difference in electronegativity between the two atoms, the linear shape mean it is symmetrical which cancels out the dipoles
Non-polar or polar:
OCl2
explain why (3)
Polar- O is more electronegative than Cl, and there is two pairs of lone electrons around the oxygen (central) atom, meaning the molecule is bent, as the bonds are not cancelled out.
Non-polar or polar:
SF6
explain why (3)
Non-polar- although there is a difference in electronegativity in a single bond, the symmetry of 90 degrees cancels out the dipoles.
Non-polar or polar:
CH3Cl
explain why (3)
Polar- the molecule is unsymmetrical because Cl is more electronegative than C, so an outward dipole moment is produced (as the charge flows from + to -)
Describe what intermolecular forces are (3)
- Attractive forces between molecules
- Much weaker than ionic or covalent bonds
- Only found in covalent structures
Name the 3 types of intermolecular forces
- Induced dipole-dipole interactions (London forces)
- Permanent dipole-dipole interactions
- Hydrogen bonds
In which molecules are London forces found?
Between all molecules, whether polar or non-polar (as long as there is 2+ atoms)
Explain how London forces arise between molecules (3)
- There is an uneven distribution of electrons in a molecule
- This causes an induced dipole at either ends of the molecule
- This dipole causes another dipole in a neighbouring molecule (delta + on one molecule attracts to delta - on another molecule)
Therefore, they can be found between 2+ atoms bonded together
In what molecules are permanent dipole-dipole interactions found?
Found between polar molecules
What intermolecular force(s) are found in hydrochloric acid? Explain why. (3)
- London forces and permanent dipole-dipole interactions
- LF’s because there may be an imbalance of electrons between the 2 atoms at any point
- PDDI’s because of a polar bond- difference in electronegativity and no symmetry so the dipoles do not cancel
Explain how permanent dipole-dipole interactions arise between molecules (3)
- Due to a permanent polar bond (no symmetry between the dipoles)
- The delta + on one molecule attracts to the delta - on a neighbouring molecule
- This produces a permanent dipole-dipole interaction between the molecules
Compare the strength of the three types of intermolecular forces
- Both London forces and permanent dipole-dipole interactions are weak, however London forces are weaker
- Hydrogen bonds are the strongest
What is a hydrogen bond?
A strong dipole-dipole interaction between molecules containing O-H, N-H or F-H bonds
Describe how a hydrogen bond is formed (2)
- A hydrogen atom attracts to the lone pair on a O, N or F atom on a neighbouring molecule (highly electronegative)
- This forms a delta positive hydrogen atom and a delta negative O, N or F atom
How are hydrogen bonds represented?
-As a straight, dashed line from the lone pair of electrons to the hydrogen atom on a neighbouring molecule
- Delta + and delta - shown correctly
Why does water have a higher boiling point than ammonia? (3)
-Water has 2 lone pairs around the oxygen atom, whereas ammonia only has 1 lone pair around the nitrogen atom
- Water can thus form 2x more hydrogen bonds than ammonia
- More hydrogen bonds= water requires far more energy than hydrogen to overcome.
Identify and explain the two anomalous properties of water caused by hydrogen bonding (5)
- Ice is less dense than liquid water. In ice, the H2O molecules are held further apart due to the lengthened H bonds, which forms an open lattice structure
- Water has relatively high melting and boiling points, because the hydrogen bonds are stronger than other intermolecular forces, so more energy is required to overcome them.
How do intermolecular forces explain simple covalent compounds sometime being soluble in water?
Water is a polar molecule, and hydrogen bonded molecules can form hydrogen bonds with water so are soluble.
How do intermolecular forces explain simple covalent compounds not conducting electricity?
Overall covalent compounds are uncharged, permanent dipoles are not strong enough.
Explain how a hydrogen bond formed
Hydrogen has a high charge density and F,N and O are very electronegative. The bond is so polarised that a weak bond forms between the hydrogen of one molecule and a lone pair on a neighbouring molecule’s F,N or O.
What effects do hydrogen bonding have on a molecule? (2)
- Soluble in water
- Higher boiling and freezing points than molecules of a similar size that don’t form hydrogen bonds
Magnesium boiling point: 650C
Chlorine boiling point: -101C
Describe the structure and bonding of these elements and explain the difference in melting points (6)
- Mg has a giant metallic lattice structure
- Electrostatic attraction between cations and delocalised electrons
- Cl has a simple covalent structure
- Cl has only induced dipole-dipole interactions (London forces) which are much weaker
- Less energy is needed to overcome induced dipole-dipole interactions than metallic bonds
Which has the highest boiling point?
A) Bromine
B) Silicon
C) Phosphorus
D) Sulfur
E) Argon
Order the rest from highest to lowest.
B- Silicon
Silicon has a giant covalent structure, as every silicon atom forms 4 covalent bonds, so requires lots of energy to overcome
Then Sulfur (S8), Phosphorus (P4), Bromine (Br2) and Argon.
Compare the boiling points of H2O, H2S and H2Se
H2O has the highest boiling point due to the hydrogen bonds which form due to the high electronegativity of O. H2Se has a higher boiling point than H2S as H2Se has more electrons in it’s atom, so stronger London forces which require more energy than H2S to overcome.
