Bonding Flashcards
What is ionic bonding
The transfer of electrons between metals and non-metals
Structure of ionic compounds
Ionic compounds have a giant ionic lattice structure
Properties of ionic compounds (4)
-high melting / boiling points
-electrical conductors when molten or dissolved
-usually soluble in water and solid at room temperature
-brittle and shatter easily when given a sharp blow
Why do ionic compounds have high melting points
The giant ionic lattices are held together by many strong electrostatic forces of attraction between oppositely charged ions which require lots of energy to overcome
Why can ionic compounds conduct electricity when molten or dissolved
This is because when molten or dissolved, the ions are free to move (delocalised) and can carry a charge whereas in a solid, the ions are held in fixed positions by the strong ionic bonds
Why are ionic compounds soluble in water
Water molecules are polar and can therefore pull the ions away from the lattice, causing it to dissolve
Why are ionic compounds brittle and shatter easily when given a sharp blow
The blow may distort the ions resulting in like charges to coming into contact with each other, causing them to repel (so the compound shatters)
What is covalent bonding
The sharing of electrons between non-metals
Structure and bonding in simple covalent compounds (small molecules/ simple molecular compounds)
Atoms in the molecules are held together by strong covalent bonds but the molecules are held together by weak intermolecular forces of attraction
Properties of simple covalent compounds (3)
- can’t conduct electricity
- low melting point
- soluble in water
Why can’t simple covalent compounds conduct electricity
There are no delocalised electrons to carry the charge as they have full outer shells
Why do you simple covalent compounds have a low melting point
The many weak intermolecular forces of attraction between the molecules requires little energy to overcome
Examples of giant covalent structures
-diamond
-graphite
-silicone dioxide
Structure and bonding in graphite
Graphite is a giant covalent molecule. Each carbon atom forms covalent bonds with 3 other carbon atoms leaving one delocalised electron per atom. Graphite has a hexagonal structure and has layers (one layer is called graphene)
Properties of graphite (5) and explanations
-soft and slippery: graphite has layers which have weak intermolecular forces of attraction between them which are easily broken allowing the layers to slide over one another
-high melting point: it’s a giant covalent structure and has many strong covalent bonds between the atoms which require a lot of energy to overcome
-can conduct electricity: has delocalised electrons which a pre free to carry the charge
- low-density: layers are far apart compared to the length of the covalent bonds
-insoluble: covalent bonds in the layers are too difficult to break
Structure and bonding in diamond
Diamond has a giant covalent structure, each carbon atom forms covalent bonds with 4 other carbon atoms leaving no delocalised electrons per atom. Diamond has a tetrahedral structure
Properties of diamond (3) and explanations
Hard and strong/ rigid due to tetrahedral arrangement
Can’t conduct electricity as there’s no delocalised electrons to carry the charge
High melting point as it’s a giant covalent structure, the many strong covalent bonds require a lot of energy to overcome
Definition of dative covalent bonding
When an atom with a lone pair of electrons donates the lone pair to an electron deficient atom
Define metallic bonding
Metals consist of a lattice of positive ions existing in a sea of delocalised electrons
Properties of metals (4) and explanations
- Good electrical and thermal conductors
- strong, the more electrons (more electrostatic attraction) and the smaller the ion (nucleus closer to outer shell electron), the stronger the metal is
- malleable and ductile as metal ion is still in the same environment after a small distortion
- high melting points due to many strong electrostatic forces of attraction between positively charged ions and delocalised electrons
Definition of electronegativity
The power of an atom to attract a bonding pair or electrons in a covalent bond
Factors affecting electronegativity and reasons (3)
- nuclear charge (increases across a period as number of protons increases while shielding stays the same)
- electron shielding
- distance between nucleus and outer shell electrons (the smaller the atom, the closer the nucleus is to the shared pair of electrons which increases electronegativity)
What’s the relationship between electronegativity and bond polarity
As electronegativity difference between atoms increases, bond polarity also increases (if atoms have the same electronegativities, the molecule will be non-polar as there’s no overall dipole moment)
Why aren’t symmetrical molecules polar
Symmetrical molecules can have polar bonds but the dipoles cancel each other out as they’re in equal and opposite directions and therefore the overall molecule is non polar as there’s no net dipole moment
Define electron pair repulsion theory
Each pair of electrons around an atom will repel all other electron pairs, so pairs of electrons will take up positions as far apart as possible to minimise repulsion
Lone pairs vs bonding pairs (repulsion) and how they affect bond angles
Lone pairs repel more than bonding pairs (having lone pairs on an atom decreases the bond angle by 2.5 degrees)
Shapes of molecules- linear with bond angle
2 bond pairs
180 degrees
Shapes of molecules- bent and bond angle
2 bond pairs
2 lone pairs
104.5 degrees
Shapes of molecules- trigonal planar and bond angle
3 bond pairs
120 degrees
Shapes of molecules- trigonal pyramidal and bond angle
3 bond pairs
1 lone pair
107 degrees
Shapes of molecules- tetrahedral and bond angle
4 bond pairs
109.5 degrees
Shapes of molecules- square planar and bond angle
4 bond pairs
2 lone pairs
90 degrees
Shapes of molecules- trigonal bipyramidal and bond angle
5 bond pairs
90 and 120 degrees
Shapes of molecules- octahedral and bond angle
6 bond pairs
90 degrees
What are the 3 types of intermolecular forces
Van deer waals
Dipole-dipole
Hydrogen bonding
Explain how van der waals forces arise
AKA London dispersion forces (weakest IMF)
Uneven distribution of electrons
One side of molecule slightly more electronegative than the other which induces a temporary dipole in the neighbouring molecule
This creates a temporary van der waals force between the electronegative side of one molecule and the electropositive side of another molecule
Explain dipole-dipole forces
(Only in polar molecules)
Dipole-dipole forces form between all molecules with an electronegativity difference
Explain hydrogen bonding
(Strongest IMF)
occur in molecules that have a significant electronegativity difference so can only form between a hydrogen and an oxygen/nitrogen/fluorine
