bonding (A1 physical chemistry) Flashcards
molecular ions - hydroxide, nitrate, ammonium, sulfate, carbonate
OH⁻
NO₃⁻
NH₄⁺
SO₄²⁻
CO₃²⁻
swap and drop method to work out the formula of an ionic compound e.g. calcium nitrate, calcium oxide
- write the two ions with their charges
- swap the charges of the ions
- drop the charges to subscript
- if necessary, simplify to lowest whole number ratio
- Ca²⁺ and NO₃⁻
- Ca⁻ and NO₃²⁺
- Ca and (NO₃)₂
- Ca(NO₃)₂
- Ca²⁺ and O²⁻
- Ca²⁻ and O²⁺
- Ca₂ and O₂
- CaO
properties of ionic compounds and giant ionic structures
most ionic compounds dissolving water as the water molecules are polar and can attract the positive and negative ions to break up the structure
conduct electricity when molten or dissolved in solution as the ions are free to move around
have high melting points as there are many strong electrostatic forces between oppositely charged ions that require lots of energy to overcome
giant ionic structures form a regular structure with a cubic shape and giant repeating pattern (e.g. sodium chloride)
compare the structure, bonding, and properties of graphite and diamond
graphite and diamond are both giant covalent structures
in graphite, each carbon is bonded three times with a fourth delocalised electron
this means that graphite forms layers which slide easily over each other as there are weak forces between them
graphite has a low density as the layers are far apart in comparison to the covalent bond length
the delocalised electrons between the layers allow graphite to conduct electricity as they can carry a charge
graphite has a very high melting point as there are lots of strong covalent bonds, this means graphite is insoluble as the covalent bonds are too strong to break
diamond forms of tetrahedral shape as each carbon is bonded four times
diamond can conduct heat well due to its tightly packed rigid arrangement
diamond can be cut to make gemstones unlike graphite which is too fragile and would break into layers
diamond can’t conduct electricity as it does not have any delocalised electrons
diamond is hard and insoluble with a very high melting point due to many strong covalent bonds
2 bonding pairs, no lone pairs
name of shape: linear
bond angle: 180°
example: BeCl₂
3 bonding pairs, no lone pairs
name of shape: trigonal planar
bond angle: 120°
example: BF₃
4 bonding pairs, no lone pairs
name of shape: tetrahedral
bond angle: 109.5°
example: CH₄
5 bonding pairs, no lone pairs
name of shape: trigonal bipyramidal
bond angles: 120° and 90°
example: PCl₅
6 bonding pairs, no lone pairs
name of shape: octahedral
bond angle: 90°
example: SF₆
3 bonding pairs, 1 lone pair
name of shape: pyramidal
bond angle: 107°
example: NH₃
based on a tetrahedral shape, replacing one bonding pair with a lone pair therefore reducing bond angle by 2.5° due to lone pair repulsion
2 bonding pairs, 2 lone pairs
name of shape: non-linear
bond angle: 104.5°
example: H₂O
based on tetrahedral shape, replacing two bonding pairs with lone pairs therefore reducing bond angle by 5° due to lone pair repulsion
3 bonding pairs, 2 lone pairs
name of shape: trigonal planar
bond angle: 120°
example: ClF₃
based on trigonal bipyramidal shape, bond angle does not change due to lone pairs cancelling each other out
4 bonding pairs, 2 lone pairs
name of shape: square planar
bond angle: 90°
example: XeF₄
based on octahedral shape, bond angle does not change due to lone pairs cancelling each other out
relationship between electronegativity and polarity
the bigger the difference in electronegativity between atoms in a covalent bond, the more polar the bond
the more electronegative element is δ⁻
uneven distribution of charge leads to polar molecules (e.g. H₂O, dipole doesn’t cancel out due to bent shape)
if polar bonds are arranged symmetrically then there is no overall polarity (e.g. CO₂, oxygen is more electronegative however electrons are pulled towards oxygens symmetrically on either side)
explain types of intermolecular forces
induced dipole-dipole - a temporary dipole is formed when a molecule is nearby to another molecule, electrons will move from one end to another causing uneven distribution and one side to become partially negative (δ⁻), there will be a force of attraction between the partially positive (δ⁺) end of one molecule and the partially negative (δ⁻) of another, however when they move away this dipole interaction is lost
permanent dipole-dipole - these interactions exist in molecules with polarity, stronger than induced dipole-dipole, they are weak electrostatic forces between molecules, molecules that have permanent dipole interactions also have van der waals forces
we can test polar molecules by placing a charged rod near a steady stream of a polar liquid, and should see the liquid bend towards the rod as the molecules align to face the oppositely charged rod
hydrogen bonding - strongest intermolecular force, occurs with very electronegative elements, hydrogen bonded with the lone pair on nitrogen, oxygen, or fluorine, can show hydrogen bonding by using a dotted line between a lone pair and a hydrogen atom, molecules that have hydrogen bonding also have van der waals forces and dipole-dipole interactions
explain the structure of iodine
simple molecular structure
van der waals forces hold iodine in a crystal structure
strong covalent bonds hold the 2 iodine atoms together, but weak van der waals hold the I₂ molecules together
explain the trend in van der waals forces in hydrocarbons
hydrocarbons do not have polar bonds, therefore they are held together by van der waals forces
longer, straight chain hydrocarbons have more van der waals forces, therefore higher boiling points
branched hydrocarbons have weaker van der waals forces and lower boiling points as they cannot pack s closely together
the bigger the molecule or atom, the more van der waals forces as they have larger electron clouds
explain the structure of ice and compare it to the structure of water
ice forms a regular, simple molecular structure held together by hydrogen bonding
the hydrogen bonding pushes the molecules further apart than in water, making ice less dense than water
when we boil water, we are breaking the van der waals forces holding the molecules together, not the covalent bonds
explain the boiling points of hydrogen halides
HF has a higher boiling point than HCl as it has hydrogen bonding, so more energy is required to overcome the electrostatic forces
there is a slight increase in boiling point from HCl to HI due to increased mass of molecule, therefore a bigger electron cloud and more van der waals forces
explain the energy changes associated with changes of state
particles in a solid are tightly packed and in a regular arrangement, therefore they have a high density, particles vibrate on the spot and can’t be compressed
particles in a liquid have more energy than in a solid, as they can move freely, however it is still difficult for liquids to be compressed, they have a high density and are tightly packed but in a random arrangement
particles in a gas have more energy than in liquids and solids as they can move freely, and it is easy for them to be compressed, particles are spaced out and in a random arrangement, so gases have a low density
properties of metal
metals have giant metallic lattice structures
metals are good thermal conductors as the delocalised electrons can transfer kinetic energy
metals are good electrical conductors as the delocalised electrons are mobile and can carry a current
metals have high melting points due to the strong electrostatic attractions between positive metal ions and negative delocalised electrons (increases with charge and number of electrons donated to sea of electrons)
solid metals are insoluble as the metallic bond is too strong to break
compare the properties of giant covalent, simple molecular, and giant ionic bonding
giant covalent e.g. graphite, diamond, SiO₂ - usual state at room temperature and pressure is solid, can’t conduct electricity as a solid (graphite is an exception), cannot conduct electricity as a liquid (difficult to melt, and normally sublime), not soluble in water, high melting and boiling points as a lot of of energy is required to break the strong covalent bonds
simple molecular e.g. I₂, H₂O, NH₃ - usual state at room temperature and pressure is liquid or gas (iodine is solid), cannot conduct electricity as a solid or a liquid, solubility in water depends on the polarity of the molecule (polar molecules dissolve well in polar solvents), low melting and boiling points as only weak intermolecular forces need to be broken
giant ionic e.g. NaCl, CaO, MgBr₂ - usual state at room temperature and pressure is solid, cannot conduct electricity as a solid, can conduct electricity as a liquid due to free ions, soluble in water, high melting and boiling points as a lot of energy is required to break the strong electrostatic forces
