Chemical Bonding Flashcards
Ionic bonding
The electrostatic attraction between oppositely charged ions
Giant ionic lattice
A three dimensional structure of oppositely charged ions held together by ionic bonds
Properties of ionic compounds
- High melting point and boiling point as ionic bonds are strong and have to all be broken to change states
- Conducts electricity when molten or aqueous as ions are free to move to carry charge
- Dissolves in polar substances. Six O- rip off a cation while 4 H+ rip off the anion
- Don’t dissolve in non-polar solvents
- Don’t conduct electricity as a solid
- Tend to be brittle as when the layers shift, the opposite ions will repel
Isoelectronic
To have the same number of electrons e.g. Si4+ is isoelectronic with Ne
Covalent bond
The electrostatic attraction between two nuclei and the same shared pair of electrons
Lone pair
An outer shell pair of electrons that is not involved in chemical bonding
The octet rule
The rule that only 8 electrons occupy one shell when bonding with other atoms
Dative covalent bond
A covalent bond formed when a lone pair has been provided by one of the bonding atoms
Properties of diamond
- Doesn’t conduct electricity as there are no free charge carriers available
- Giant covalent lattice
- Tetrahedral C arrangement
- Bare hard as covalent bonds are strong and a lot of energy is needed to break them all
Properties of graphite
- Giant covalent structure
- Conducts electricity as there is a layer of delocalised electrons between the C layers that can carry charge
- C arranged in layers
- Layers can slide over each other so graphite can acts as a lubricant
Allotrope
Refers to one or more forms of an elementary substance
Metallic bonding
The electrostatic attraction between cations and a sea of delocalised electrons
Properties of metals
- High melting and boiling points as there is a large electrostatic attraction between the cations and the delocalised electrons
- Good conductors as delocalised electrons can carry charge
- Even better conductors when molten as the cations can carry charge too
- Insoluble in water or non-polar substances as the metallic bonds are too strong relative to intermolecular forces
- Malleable
- Ductile
Trend in mp/bp of metals (I,II,III)
- Increases along the group as the number of electrons are increasing and the charge on the cations are too
- Decreases down a group as the outer electron is far from the nucleus and can be removed easily
Malleability
The ability to change shape without breaking
Ductility
The ability to be pulled into a thin wire
Alloys
A mixture of metals; not a compound
E.g. stainless steel = steel + chromium
Brass = copper + zinc
Use of alloys
Alloys make metals harder as the ions sizes of the metals are different, so the layers cannot slide over each other
Valence Shell Electron Pair Repulsion theory
Each bonded pair repel themselves so they are as far away from eachother
Lone pairs repel more than bonding pairs
Linear
e.g. BeCl2
2 bonding pairs 0 lone pairs
bond angle 180 degrees
Trigonal planar
e.g. BF3
3 bonding pairs 0 lone pairs
bond angle 120 degrees
Tetrahedral
e.g. CH4
4 bonding pairs 0 lone pairs
bond angle 109.5 degrees
Trigonal pyramid
e.g. NH3
3 bonding pairs 1 lone pair
bond angle 107 degrees
Bent
e.g. H2O
2 bonding pairs and 2 lone pairs
bond angle 104.5 degrees
Octahedral
e.g. SF6
6 bonding pairs 0 lone pairs
bond angle 90 degrees
The three types of intermolecular force
van der Waals’
Permanent dipole-dipole
Hydrogen bonding
How do van der Waals’ forces arise?
Electrons are constantly moving. At any instant, the distribution may not be symmetrical. This results in an instantaneous temporary dipole. This dipole induces dipoles in neighboring molecules and leads to an attraction between the opposite charges in the dipoles
Factors that affect the van der Walls’ force
- Number of electrons - the more the stronger
- The weaker the contact area, the stronger the induced dipole
- Unbranched molecules have stronger van der Waals’ forces
Trends in group VII
State progressively changes from gas to solid down the group
Reactivity decreases down the group
Mp/Bp increases down the group
Electronegativity
The ability of an atom to attract electrons to itself in a bond
Electronegativity’s of some elements
F=4 O,N=3.5 Cl=3 Mg=1.2 Na=0.9 Al=1.5
Permanent dipole
A small charge difference across a bong resulting from a difference in electro-negativities of the bonded atoms
Permanent dipole-dipole attractions occur in addition to van der Waals’ forces
Trend in electronegativity
Electronegativity increases across a period and decreases down a group
Use of electronegativity
Electronegativity can be used to predict bonding type
Electronegativity > 2.0 usually ionic
Electronegativity 1.0~2.0 usually polar
Electronegativity < 1.0 usually covalent
Relative strengths of bonds and intermolecular forces
Ionic/covalent bonds 1000
Hydrogen bonds 50
Permanent dipole-dipole attraction 10
Van der Waals’ 1
Polar covalent bond
A covalent bond with a permanent dipole
Polar molecule
A molecule with an overall dipole when you take into account any dipoles across the bond
When drawing hydrogen bonded molecules
Draw the hydrogen bond with a striped line
Draw the delta + and -
Position atoms either side of the H+ as far away from each other
Draw any lone pairs
Hydrogen bonds
The attraction between a hydrogen atom on one molecule and a lone pair of another O, N or F molecule
The reason water is not like the other Group VI hydrides
All group VI hydrides have a van der Waals’ attraction. On top of this, water has hydrogen bonding so its mp/bp are different
The reason ice floats on water
In a solid, water forms a rigid structure held apart by hydrogen bonds. When this structure melts, the rigid hydrogen bonds break and the molecules move closer together
