BONDING Flashcards
Define ionic bonding
The electrostatic force of attraction
between oppositely charged ions
formed by electron transfer
Metal atoms ___ electron to form _ve ions
Non-metal atoms ___ electron to form _ve ions
Metal lose to form +ve
Non-metal gain to form -ve
How do smaller ions impact melting points
High melting points
In smaller ions
Because ionic bonding is stronger
Or a higher charge
Define covalent bonding
A shared pair of electrons
Define dative covalent bonding / Coordinate bindi g
A shared pair of electrons
Donated from only one of the bonding atoms
Give the 3 common dative covalent molecules
NH4 +
H3O +
NH3BF3
How is a dative covalent bond represented
An arrow
From the atom providing the lone pair
To the deficient atom
Define metallic bonding
Electron static force of attraction
Between positive metal ions
And delocalised electrons
List the 3 factors effecting the strength of a metallic bond
Proton number
Delocalised electron number (per atom)
Size of ion
How does proton number affect the strength of metallic bonding
More protons
Stronger bond
How does delocalised electron number affect the strength of metallic bonding
More delocalised electrons
Stronger bond
How does ion size affect the strength of metallic bonding
Smaller ion
Stronger bond
Give an example of a giant ionic lattice molecule
Sodium chloride
Magnesium oxide
Give an example of a simple molecular molecule
Iodine
Ice
Carbon dioxide
Water
Methane
Give an example of a macromolecular molecule
Diamond
Graphite
Silicon dioxide
Silicon
Give an example of a giant metallic lattice molecule
Magnesium
Sodium
All metals
Outline the properties of giant ionic lattices
High melting and boiling point
Soluble in water
Poor conductor when solid
Good conductor when molten/dissolved
Crystalline solid
Outline the properties of a simple molecular molecule
Low melting and boiling point
Poor solubility in water
Poor conductor when solid and molten
Mostly gas or liquid
Outline the properties of a macromolecule
High melting and boiling point
Insoluble in water
Diamond/Sand are poor conductors when solid but graphite is good
Poor conductors when molten
Solids
Describe the properties of a metallic molecule
High melting and boiling point
Insoluble in water
Good conductors when solid and molten
Shiny
Malleable
Outline a linear molecule. Give an example.
2 bonding pairs
No lone pairs
180 bond angle
CO2, CS2, HCN, BeF2
Outline a trigonal planar molecule. Give an example.
3 bonding pairs
No lone pairs
120 bond angle
BF3, AlCl3, SO3, NO3 -, CO3 2-
Outline a tetrahedral molecule. Give an example.
4 bonding pairs
No lone pairs
109.5 bond angle
SiCl4, SO4 2-, ClO4 -, NH4 -
Outline a trigonal pyramidal molecule. Give an example.
3 bonding pairs
1 lone pair
107 bond angle
NCl3, PF3, ClO3, H3O +
Outline a bent molecule. Give an example.
2 bonding pairs
2 lone pairs
104.5 bond angle
OCl2, H2S, OF2, SCl2
Outline a trigonal bipyramidal molecule. Give an example.
5 bonding pairs
No lone pairs
120 and 90 bond angles
PCl5
Outline a octrahedral molecule. Give an example.
6 bonding pairs
No lone pairs
90 bonding angle
SF6
Outline the structure for explaining the shape of a molecule
- state number of bonding pairs and lone pairs
- State electron pairs repel to a position of minimum repulsion
(3. State lone pairs repel more than bonding pairs) - State actual shape and bonding angle/s
How much do lone pairs reduce bond angles
2.5
Outline a square planar molecule. Give an example.
4 bonding pairs
2 lone pairs
Bond angle 90
XeF4
Outline the shape of BrF5
Bond angle 89
4 bond pairs
2 lone pairs
Outline the shape of I3 -
Bond angle 180
2 bonding pairs
3 lone pairs
Outline the shape of ClF3
Bond angle 89
3 bonding pairs
2 lone pairs
Outline the shape of SF4 of IF4 +
Bond angles of 119 and 89
3 bonding pairs
1 lone pair
Define electronegativity
The relative tendency
of an atom in a covalent bond
in a molecule to attract electrons
in a covalent bond to itself
List the 4 most electronegative elements
Fluorine
Oxygen
Nitrogen
Chlorine
What factors affect electronegativity
Proton number
Atomic radius
Shielding
Outline how electronegativity changes across the period
Increases
Larger proton numbers
Smaller atomic radius
Electrons closer to nucleus
Outline how electronegativity changes down a group
Decreases
More shells
More shielding
Electrons further from the nucleus
How does electronegativity effect the type of bonding in a compound
Similar electronegativity
Covalent bond
Large difference in electronegativity
Ionic bond
Outline how a permanent dipole (polar covalent) bond forms
Elements in the bond have different electronegativity
Difference of around 0.3 to 1.7
Causing an unequal distribution in electrons
And thus a charge separation
(Delta) + and (delta) -
Outline the effects of symmetry in molecules
All bonds are identical
With no lone pairs
Cannot be a polar molecule
As the any dipoles cancel out
Due to symmetry
When does van der waals occur
Between all molecular substances
And Nobel gases
Not in ionic substances
Describe how van der waals (induced dipole) occurs
Electron density fluctuates
Parts of the molecule become more/less negative
Induces neighbouring molecules to do the same (induced dipole)
The induced dipole is the opposite to the original one
Outline the factors effecting van der waals
More electrons
More likely to form
Increased van der waals
Higher boiling points
Long chain alkanes more likely
More surface area
Eg halogens increase down group due to more van der waals
When does permanent dipole-dipole occur
Between polar molecules
Outline permanent dipole dipole forces
Stronger than van der waals
Further increase boiling point
Molecules are asymmetrical
and have significant electronegativity difference between atoms
When does hydrogen bonding occur?
Between hydrogen atoms and either:
Oxygen
Nitrogen
fluorine
With a lone pair available
Outline hydrogen bonding
Strongest intermolecular bond
Causes very high boiling points (eg water is a liquid not gas)
Alcohols, carboxylic acids, proteins, amides all form hydrogen bonds
Order the intermolecular forces from strongest to weakest
Hydrogen bonding
Permanent dipole dipole
Van der waals
