2.2.2 Bonding and structure Flashcards
Define ionic bonding
What does a dot and cross diagram show ?
the electrostatic attraction between positive and negative ions
where the electrons in a bond come from
Define metallic bonding
the** strong electrostatic attraction **between a **lattice of positive ions **and a sea of delocalised electrons
What is a giant ionic lattice ?
regular structure of oppositely charged ions that are strongly attracted in all directions
Explain how the structure of ionic compounds impacts melting and boiling points ?
High melting/boiling points
* giant ionic lattice is held my many electrostatic forces of attraction
* these need a lot of energy to be overcome
* so high melting/boiling points
Ionic compounds are typically solid at room temp
Strength depends on size and charge
**-Smaller ion/higher charge = stronger ionic bonds
**
Explain how the structure of ionic compounds impacts electrical conductivity
- Do not conduct when solid as ions are held in fixed positions so no charge is conducted
- Conduct when molten or dissolved as ions are **mobile **and can carry a charge
Explain how the structure of ionic compounds impacts solubility
What is the exception to this rule ?
Soluble in polar solvents such as water
* The slight negative oxygen will be attracted to positive ions
* The slight positive hydrogens will be attracted to the negative ions
Ions are attracted and surrounded.
This causes the ionic lattice to break down and the ionic compound dissolves
If the ions have large charges the ionic attraction may be too strong for the water to break down the lattice structure (solubility decreases as ionic charge icreases)
Define covalent bonding
strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
It is the overlap of atomic orbitals - The attraction is localised, so only goes in two directions
What does the dot and cross diagram show in covalent bonding ?
shows the electrons transferred
Why is BF^3 an exception to the covalent bonding rules ?
Boron is **electron deficient **as it only has 6 electrons in its outer shell
This is allowed because:
* as long as the central atom maximises the covalent bonds it can form it will be stable
* Boron is in group 3
* So it forms a maximum of 3 covalent bonds
Why is SF^6 (sulfur hexofluoride) an exception to the covalent bonding rules ?
It is has an expanded octet
* outer shell has 12 electrons
* It is in period 3 so has access to the third shell
* This can hold 18 electrons, and the 6 unpaired are available for bonding
What are the three types of covalent bond ?
- single
- Double/triple
- Dative (coordinate)
Define dative bond
a covalent bond in which the shared pair of electrons comes from only one of the bonding atoms
-The shared pair was originally a LONE PAIR
- Movement of the lone pair is shown by an arrow
Define average bond enthalpy
measurement of covalent bond strength
* larger value = stronger covalent bond
What are the 4 rules in electron pair repulsion theory ?
- **electron pairs **in the valence shell repel as far as possible
- The greater the number of electron pairs = the smaller the bond angles
- multiple bonds have the same repulsion to single bonds
- Lone pairs repel more strongly than bonded pairs (which reduced the bond angles)
What do the 3 wedges represent in molecule shape diagrams ?
- solid line = bond on the plane
- solid wedge = comes out of the plane
- dotted wedge = goes into the plane
Give the bonded/lone pairs and the angle in linear molecules
2 bonded pairs
No lone pairs
180 degrees
eg Ab^2
Give the bonded/lone pairs and angle in trigonal planar molecules
3 bonded pairs
No lone pairs
120 degrees
eg AB^3
Give the bonded/lone pairs and angle in a tetrahedral molecule
4 bonded pairs
No lone pairs
109.5 degrees
eg AB^4
Shape is now 3D as one bond goes into the plane and one goes out
Give the bonded/lone pairs and angle in a trigonal bipyramidal molecule
5 bonded pairs
No lone pairs
There are now two bond angles
* 3 bonds on the same plane = 90 degree
* 2 bonds going in and out = 120 degree
AB^5
Give the bonded/lone pairs and angles in an octahedral molecule
6 bonded pairs
No lone pairs
90 degree angles
Why do lone pairs have greater repulsion ?
How much do they reduce the bond angle by ?
Show the relative repulsion of the different pairs
slightly closer to the central atom and occupies more space
-reduce bond angle by 2.5 degrees
Bonded-bonded < bonded-lone < lone-lone
Give the bonding/lone pairs and angle of a pyramidal molecule
4 bonding pairs
1 lone pair
-lone pairs repel more than the bonding pairs, so reduce the bond angles by 2.5
Angle = 109.5 - 2.5 = 107
examples = CH^4 (methane), NH^4 (ammonia)
Give the bonding/lone pairs and angle of a non-linear molecule
2 bonding pairs
2 lone pairs which reduces the bond angles by 5 degrees
Angle = 109.5 - 5 = 104.5 degrees
Example = H^2O (water)
How should you explain the shape of a molecule in an exam question ?
- state the number of bonding/lone pairs
- state that electron pairs repel
- if no lone pairs = electrons repel equally
- if lone pairs = lone pairs repel more than bonding pairs
- state the actual shape and bond angle
Define electronegativity
ability of an atom to attract the bonding electrons in a covalent bond
What does a high pauling value show ?
a high electronegativity (increases up and across the periodic table)
* fluorine, oxygen and nitrogen are the most electronegative elements
* Group one metals are the least electronegative
What is the electronagativity difference of an ionic bond ?
high
as both of the ions have full charges
- one bonded atom will have a much** greater attraction f**or the shared pair than the other
- so will attract these electrons causing them to be **transfered **(and so forms an ionic bond)
What is the electronegativity difference of a polar covalent bond ?
moderate
as both the atoms have a partial charge
What is the electronegativity difference of a non-polar covalent bond ?
low
as the atoms are uncharged
How does electronegativity change across a period ?
increases
* as the number of protons increases
* and the atomic radius decreases
* so the attraction of the bonding atoms will be greater
How does electronegativity change down a group ?
decreases
* distance between the nucleus and outer electrons increases
* shielding of inner electrons increases
* so the attraction of the bonding atoms will be reduced
Define non-polar bond
bonded electron pair is shared equally between the bonded atoms
* occurs when bonded atoms are THE SAME or have the SAME ELECTRONEGATIVITY
eg H-H or Cl-Cl
Define polar bond
bonded electron pair is **shared unequally **between the bonded atoms
* occurs when bonded atoms have **DIFFERENT electronegativity **values
eg H-Cl
- chlorine is more electronegative
- so has a greater attraction for the bonded pair
- causing polar bond (there is a slight negative and slight positive charge - polarised) - shown by a delta sign
Define dipole
the separation of opposite charges
Define permanent dipole
the difference in electronegativity of two atoms in a polar bond
(dipole is unchanging)
What type of dipole is produced when atoms are arranged:
* symmetrically
* unevenly
Symmetrical
- dipoles cancel
- molecule has NO overall dipole so is NON-POLAR
Uneven
-dipoles don’t fully cancel
- molecule has an overall dipole
- so is POLAR
Define intermolecular forces
weak attractive forces between dipoles of DIFFERENT molecules
(less energy needed to break then compared to ionic or covalent)
Molecule = a small group of covalently bonded atoms
What is an induced dipole-dipole interaction ?
aka London forces
weakest intermolecular force
attraction between induced dipoles
- exist between ALL MOLECULES
How do induced dipole-dipole interactions arise ?
- electrons are constantly moving - electron density shifts
- at any moment** electron density** could be** uneven**
- forming an instantaneous dipole
- Electron density on one end repels, causing an **induced dipole **on the neighbour molecule
- This causes a chain reaction (induced dipole induces further dipoles)
They are TEMPORARY
How does the number of electrons in an induced dipole-dipole interaction impact is properties ?
- more electrons
- bigger induced dipoles (as there is a greater difference in electronegativity)
- greater dipole-dipole interactions
- stronger attractive forces between molecules
so….
* more energy needed to overcome the intermolecular forces, increasing the boiling points
Larger surface area also increases forces as there is a more exposed electron cloud
What is a permanent dipole-dipole interaction ?
weak electrostatic forces of attraction between molecules with a **permanent dipole **(significant electronegativity difference - polar molecules)
These happen in addition to induced dipole-dipole forces
- Extra energy is needed to break this force, so molecules with both induced AND permanent interactions have a higher boiling point
What is hydrogen bonding ?
specific type of permanent dipole-dipole interaction between molecules containing:
* an electronegative element with a long pair of electrons (oxygen, nitrogen, fluorine)
* a **hydrogen **bonded to the en. element
How does a hydrogen bond form ?
Strongest intermolecular force
- a **dipole forms **between the electronegative element and the hydrogen
*This bond is so polorised that a HYDROGEN BOND forms between the hydrogen and the lone pair
molecules with hydrogen bonding usually have -OH or -NH groups
Why does water have a higher boiling point than expected ?
- large number of hydrogen bonds
- a lot more energy is needed to break these bonds
- increasing the melting/boiling point
Why does water have a high surface tension ?
- many hydrogen bonds
- give water a tight structure
- this creates surface tension
Why is ice less dense than water?
- hydrogen bonds in ice hold the molecules in the open lattice further apart
- hydrogen bonds are long, so volume increases and density decreases
- so the ice floats
when ice melts, the hydrogen bonds are broken and the ice lattice collapses - molecules move closer together so water is more dense
Define simple molecular lattice
covalently bonded molecules attracted by weak intermolecular forces
The atoms WITHIN each molecule are held by strong covalent bonds
Why do simple molecules have low melting and boiling points ?
- intermolecular forces are weak
- minimal energy to break
- low melting and boiling points (liquids and gases at room temp)
larger molecule = greater intermolecular forces
Why do simple molecules not conduct electricity ?
- no charged particles
- so cannot carry charge
How does solubility vary in simple molecules?
In polar solvents…
Non-polar molecule
* only form weak induced dipole-dipole interactions
* the intermolecular forces in the polar solvent are too strong to be broken
* so molecule is insoluble in polar solvents
Polar molecule
* can form stronger hydrogen bonds
* these interact with the dipoles in the water
* so can dissolve
Polar molecules = soluble in polar solvents (water, ethanol)
Non-polar molecules = soluble in non-polar solvents (alkanes and aromantics)
