Bonding Flashcards
Ionic bonding definition
Electrostatic attraction between oppositely charged ions formed by electron transfer
Ionic bonding structure
Giant ionic lattice
Crystalline solid
Ionic bonding examples
Sodium Chloride (NaCl) Magnesium Oxide (MgO)
Ionic bonding melting and boiling points
High MP and BP
Giant lattice of ions with many strong electrostatic forces between oppositely charged ions
The smaller the ion and the higher the charge, the stronger the bond
Ionic binding solubility
Generally good
Water molecules surround ions that have broken out of lattice
When water molecule hits ionic lattice, can knock ions off to surround
Aluminium oxide (Al2O3) not soluble in water as electrostatic attraction too strong
ionic bonding conductivity
Do not conduct electricity when solid as ions cannot move and are fixed in a lattice
Conduct electricity when molten or in aqueous solution as ions are free to move and carry charge
covalent bonding definition
Shared pair of electrons
covalent bonding structure
Simple molecular With intermolecular forces between molecules Mostly gases and liquids Macromolecular Giant molecular structures solids
covalent bonding examples
Simple molecular Iodine (I2) ice/water (H2O) Carbon dioxide (CO2) Methane (CH4) Macromolecular Diamond (carbon bonds) Graphite (carbon bonds) Silicon dioxide (SiO2) Silicon (silicon bonds)
covalent bonding melting and boiling points
Simple molecular
Low due to weak intermolecular forces between molecules
Macromolecular
High due to many strong covalent bonds in structure
covalent bonding solubility
Simple molecular
Generally poor
Macromolecular
insoluble
covalent bonding conductivity
Simple molecular
Poor as no ions free to conduct and electrons are fixed in place (localised)
Macromolecular
Most compounds poor as no ions and only localised electrons
Graphite good as free electrons between layers
Dative covalent bonding
Forms when shared pair of electrons in covalent bond comes from only one of the bonding atoms
Also called co-ordinate bond
Direction of arrow goes from atom providing lone pair to atom that is deficient
common examples of dative covalent bonding
NH4 +
H3O +
NH3BCl3
Diamond
Each carbon atom is bonded to four other carbon atoms by strong covalent bonds creating a giant covalent structure.
Very hard due to the many strong covalent bonds.
Does not conduct electricity as no delocalised electrons.
High melting point
Very rigid structure so vibrations can easily carry vibrations through the structure so diamond is a good thermal conductor.
Graphite
Each carbon atom is bonded to 3 other carbon atoms creating giant covalent structure.
High melting point.
Insoluble in solvents
Soft as carbon atoms form hexagonal rings with each layer held together by weak intermolecular forces. Layers can slide over each other, making it a good lubricant.
Conducts electricity due to delocalised electrons that can move freely.
metallic bonding definition
Electrostatic force of attraction between positive metal ions and delocalised electrons
metallic bonding structure
Giant metallic lattice
metallic bonding examples
All metals
Properties of metallic bonding
Shiny metal
Malleable as positive ions are all identical so planes of ions can slide over one another
Attractive forces are the same whichever ions are adjacent
metallic bonding melting and boiling points
High due to strong electrostatic forces of attraction
Large rectangular structure
metallic bonding solubility
Insoluble in water
React with water to form metal oxides
metallic bonding electrical conductivity
Good as delocalised electrons can move throughout structure
metallic bonding heat conductivity
High as particles can move and are close enough together to transfer thermal energy
Silver excellent conductor used in rear window defrosters in cars
Factors affecting strength of bonding
Number of protons: More protons Stronger bond Number of delocalised electrons per atom: More electrons Stronger bond Size of ion: Smaller ion Stronger bond
Electronegativity definition
Ability of an atom to attract a bonding pair of electrons in a covalent bond
Measured on Pauling Scale (0-4)
F, O, N and Cl are the most electronegative atoms
Factors affecting electronegativity
Distance of bonding electrons from nucleus
Can be taken to be the same as atomic radius
Nuclear charge
Greater nuclear charge
Decreases atomic radius
Increases electronegativity
Shielding
Inner electrons shield attractive power of nucleus
Trend across a period
electronegativity
Atomic radius decreases
Stronger attraction
Nuclear charge increases
Electronegativity increases
Trend down a group
electronegativity
Atomic radius increases
Shielding increases
Electronegativity decreases
Intermediate bonding
Compounds containing elements with a small electronegativity difference will be purely covalent
Compounds containing elements with large electronegativity difference (>1.7) will be ionic
Polar covalent bond forms when elements have electronegativity difference of 0.3-1.7
Polar and non polar molecules
Symmetrical molecules (all bonds identical and no lone pairs) will never be polar Individual dipoles cancel out so no net dipole, non polar
Intermolecular forces
Induced dipole-dipole forces
Permanent dipole-dipole forces
Hydrogen bonding
Van der waals forces
Also called transient forces and induced dipole-dipole interactions
Occur between all molecular substances and noble gases
Do not occur in ionic substances
Movement of electron clouds cause partial charges to change rapidly creating temporary dipoles. These partial charges exert force on nearby molecules resulting in induced dipole interaction
Factors affecting van der waals
More electrons increases chance that temporary dipole will form
Increases strength of van der waals forces between molecules
Permanent dipole-dipole forces
Occur only between polar molecules
Stronger than van der waals forces
Have permanent dipole within molecule
The delta+ end of one dipole is attracted to delta- end of dipole on another molecule
Hydrogen bonding
Occurs in compounds that have a hydrogen atom attached to N, O or F
Large electronegativity difference
Stronger than van der waals and permanent dipole-dipole
Weaker than any intramolecular interactions
