Bonding and Structure Flashcards
Define a Covalent Bond
A covalent bond is an electrostatic force between a shared paired of negative electrons and the positive nuclei of two atoms
Compare and contrast these two different types of bonds
- Both include two orbitals overlapping
- Sigma involves s and p whilst Pi involves p orbitals only
- Sigma overlaps two lobes whilst pi overlaps 4 lobes
Define ‘Electronegativity’
Electronegativity is a measure of a tendency of an atom to attract a bonding pair of electrons in a covalent bond
Explain the factor affecting electronegativity
- Nuclear Charge
-Atomic radius
Give the electronegativity difference ranges for different types of bonds
Non-polar covalent bond <0.4
Polar Covalent bond 0.4 - 1.7
Ionic bond > 1.7
Explain why a molecule can contain polar bonds but be non polar
- The shape of the molecule and where the polar bonds are whether the positive and negative dipoles coincide
- If the shape is symmetrical and the polar bonds cancel out as they coincide
Define an Ionic Bond
Ionic Bonding is the Electrostatic attraction between oppositely charged positive and negatively charged ions.
Describe an ionic compound and its properties
- Ionic lattice structure held together by electrostatic forces by oppositely charged ions for the maximum number of attraction
- High melting and boiling points due to strong electrostatic forces
- Brittle due to layers being able to be distorted
- Ions move when dissolved in a solution
Explain the bond lengths between C-C and C=C
C=C is shorter
C=C has a higher electron density
This leads to greater attraction between positive nuclei and negative electron
Therefore C=C has a greater bond strength
State the 2 factors affecting the strength of an ionic bond
- The greater the charge the stronger the ionic bond
- The smaller the ionic radius, the stronger the ionic bond
Define isoelectronic elements and its factor affecting its ionic radius
- a group of elements that have an identical electron configuration
- more protons result in a greater nuclear charge reducing ion radius
Describe the Valence Shell Electron Pair Repulsion Theory
Electrons pairs are negative and repel each other
They will space themselves out from one another to maximise separation and minimise repulsion
The basic shape of a molecule is the number of electrons around an central atom
State the shape and the bonding angle for 2 election bonding pairs
Linear
180°
State the shape and the bonding angle for 3 election bonding pairs
Trigonal Planar
120°
State the shape and the bonding angle for 4 electron bonding pairs
Tetrahedral
109.5°
State the shape and the bonding angle for 5 election bonding pairs
Trigonal Bipyramidal
120°, 90°, 180°
State the shape and the bonding angle for 6 election bonding pairs
Octahedral
90°, 180°
Describe the affect of lone pairs on its shape
Lone pairs cause greater repulsion and will push bonding pairs closer together, reducing bond angle by 2.5°
State the three type of intermolecular bonding
- London forces (instantaneous dipole- induced dipole)
- Permanent Dipole
- Hydrogen bonding
Describe the conditions for hydrogen bonding
- A delta positive hydrogen atom bonded to a small highly electronegative atom(O,F,N)
- Another highly electronegative atom on a separate molecule which has a lone pair
Explain the properties of H20, NH3 and HF
The anomalously high boiling points of H2O, NH3 and HF are caused by the hydrogen bonding between these molecules in addition to their London forces. The additional forces require more energy to break and so have higher boiling points
Explain why water has a relatively high boiling and temperature
- Presence of hydrogen bonds which require additional energy to break
- Water has 2 lone pairs on the oxygen atom which can form 2 hydrogen bonds
- Each water molecule can form up to 4 hydrogen bonds(oxygen - 2, each hydrogen - 1)
- Also the addition of dipole - dipole forces
Explain why ice has a lower density in its liquid form
When the hydrogen bonds are broken as the ice melts, the water molecules get closer together. That means that ice is less dense than water, and so will float on the water.
Describe the trend in alkane when the chain length increases
- Boiling temperature increases
- More electrons present
- Greater London forces (inducing dipole)
Describe the effect of branching on alkanes
- Lower boiling temperature
- Reduced surface Area contact with other molecules
- Lowers London forces
Explain why alcohols have a higher boiling point than alkanes with the similar number of electrons
- Presence of hydrogen bonding on oxygen atom in addition to London forces
- Alkanes have London forces only
Describe the trend of hydrogen halides
- HF forms hydrogen bonds whilst other form dipole-dipole or London forces
- Thisresults in a slight jump in energy between HF and HCl
-After Cl it increases due to London forces getting stronger and greater permanent dipole forces
Explain why water can dissolve some ionic compounds
The ionic bonds between cation and anion are broken.
Ion-dipole forces between oxygen and cation and hydrogen and anion are formed.
Explain if water can dissolve simple alcohols
- Water and simple alcohols can form hydrogen bond with each other
-The oxygen atom on ethanol has a lone pair which is available for a hydrogen bond
Explain if water can dissolve polar molecules like halogenoalkanes
- Water cannot form hydrogen bonds with halogenoalkanes
- Halogenoalkanes can only form dipole-dipole interactions
- Forming two immiscible(seperate) layers
Describe the two conditions for dissolving compounds
- The solute must separate from one another, and become surrounded by solvent particles
- The forces of attraction between the solvent and solute particles must be sufficiently strong to overcome the solvent-solvent and solute-solute forces
Define ‘Metallic Bonding’
The electrostatic attraction between delocalised electrons and metal ions.
State two simple covalent structures
- Iodine(I2)
- Water(H20)
State three places where giant lattice locations are found
- Ionic Solids(giant ionic lattices)
- Giant Covalent (diamond, silicon dioxide, graphite and graphene)
- Metal solids(giant metallic lattices)
Describe the properties and uses of Diamond
4 strong covalent bonds
Very Hard - Abrasives, Drills
High melting and boiling point - conducts thermal energy very well
Refracts light - Jewellery
Describe the properties and uses of Buckminster Fullerene
Roughly spherical with each carbon atom bonded to nearest 3 atoms
Very Soluble - mascara and printing ink
Delocalised electrons - Good insulators
Weak intermolecular forces - drug delivery
Describe the properties and uses of Graphite
Hexagonal arrangement in each layer, arranged in layers, each carbon bonded to 3 others
Delocalised electrons conducting electricity - Electrodes
Strong covalent bonds creating good tensile strength- Composites
Describe the properties and uses of Graphene
Singular hexangular layer of carbon boned to 3 atoms
Delocalised Electrons - good conductors
Lightweight and strong - Aircrafts or cars
Transparent, flexible, conductive - Touchscreens
