Bonding - Unit 1, Section 3 Flashcards
Ionic bonding - compounds
electrostatic attraction holds positive and negative ions together.
when atoms are held together in a lattice, it is called ionic bonding.
when oppositely charged ions form ionic bonds you get an ionic compoud.
e.g: NaCl, MgO, MgCl2
giant ionic lattices
ionic crystals are giant lattices of ions.
a lattice is a regular structure.
e.g: NaCl
Behaviours of ionic compounds x3
Electrical conductivity - conduct electricity when molten or dissolved but not when solid. Ions in a liquid are free to move and carry a charge, whereas in a solid they are in a fixed position by the strong ionic bonds.
Melting point - have high melting points as giant ionic lattices are held together by strong electrostatic forces. Takes a lot of energy to overcome the forces so melting points are very high.
Solubility - ionic compounds tend to dissolve in water. Water molecules are polar and pull the ions away from the lattice, causing it to dissolve.
Covalent bonding
occurs when 2 or more atoms are bonded together to form a molecule.
covalent bonds can be single, double or triple bonds.
Single covalent bond
2 atoms share electrons so they both have got a full outer shell.
Contains a shared pair of electrons.
Both the positive nuclei are attracted electrostatically to the shared electrons.
Double covalent bonds
contain 2 shared pairs of electrons
Triple covalent bonds
contains 3 shared pairs of electrons
Simple covalent compounds
compounds that are made up of lots of individual molecules.
the atoms in the molecules are held together by strong covalent bonds, but the molecules within the simple covalent compound are held together by weaker van der waals forces (intermolecular forces.
it is the intermolecular forces that determine the properties of simple covalent compounds.
have low melting and boiling points, are electrical insulators.
giant covalent structures
a type of crystal structure.
have a huge network of covalently bonded atoms.
sometimes called macromolecular structures.
carbon atoms form this type of structure because they can each form 4 strong covalent bonds.
graphite and diamond are 2 examples of giant covalent structures.
Graphite
- the carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each.
- The 4th outer electron of each carbon atom is delocalised. The sheets of hexagons are bonded together by weak van der Waals forces.
- the weak bonds between the layers in graphite are easily broken, so the sheets can slide over each other. Graphite feels slippery and is used as a lubricant and in pencils.
- the delocalised electrons in graphite are free to move along the sheets so an electric current can flow.
- the layers are quite far apart compared to the length of the covalent bonds, so graphite has a ow density and is used to make strong, lightweight sports equipment.
- because of the strong covalent bonds in the hexagon sheets, graphite has a very high melting point.
- it is insoluble in any solvent as the covalent bonds are too difficult to break.
Diamond
- each carbon atom is covalently bonded to 4 other carbon atoms.
- the atoms are arranged in a tetrahedral shape -its crystal lattice structure.
- diamond has a very high melting point due to strong covalent bonds.
- extremely hard due to strong covalent bonds - it is used in diamond-tipped drills.
- vibrations travel easily through the stiff lattice so it’s a good thermal conductor.
- it can’t conduct electricity as there are no delocoalised electrons.
- won’t dissolve in any solvent.
dative covalent bonds (co-ordinate bonds)
one of the atoms provides both of the shared electrons
e.g: NH4+ - nitrogen donates a pair of electrons
in diagrams it is shown by an arrow pointing away from the ‘donor’ atom.
co-ordinate bonds form when one of the atoms in the bond has a lone pair of electrons and the other doesn’t have any electrons available to share.
charge clouds
an area where you have a high chance of finding an electron
how do molecules get their shapes?
depends on the number of pairs of electrons in the outer shell of the central atom.
electrons can be shared (bonding pairs) or unshared (lone pairs)
Electron pair repulsion theory
electrons are all negatively charged, so charge clouds repel each other until they are as far apart as possible.
The shape of the cloud affects how much it repels other charge clouds.
Lone-pair charge clouds repel more than bonding pair charge clouds, so bond angles are often reduced because bonding pairs are pushed together by lone-pair repulsion.
lone-pair/lone-pair angles are biggest
lone-pair/bonding-pair angles are second biggest
bonding-pair/bonding-pair angles are smallest
drawing shapes of molecules
solid line - shows bonds that aren’t pointing away from or towards you
wedge - shows a bond pointing towards you
broken line - shows a bond pointing away from you
finding the number of electron pairs
- find the central atom
- work out how many electrons are in the outer shell of the central atom
- add 1 electron for every atom that the central atom is bonded to
- add up all the electrons
- divide by 2 to find the number of electron pairs
- compare the number of electron pairs to the number of bonds to find the number of lone pairs and the number of bonding pairs on the central atom
molecular shapes x11
LINEAR
- 2 bonding pairs
- bond angle = 180*
TRIGONAL PLANAR
- 3 bonding pairs
- bond angle = 120*
TETRAHEDRAL
- 4 bonding pairs - no lone pairs
- bond angle = 109.5*
TRIGONAL PYRAMIDAL
- 3 bonding pairs and 1 lone pair
- bond angle = 107*
NON-LINEAR
-2 bonding pairs and 2 lone pairs
-bond angle = 104.5*
TRIGONAL BIPYRAMIDAL
- 5 bonding pairs
- 3 bonds angles = 120*
- 2 bond angles = 90*
SEESAW
- 4 bonding pairs and 1 lone pair
- bond angles = 102* and 86.5*
T-SHAPE
-3 bonding pairs and 2 lone pairs
-bond angle = 87.5*
OCTAHEDRAL
- 6 bonding pairs
- bond angle = 90*
SQUARE PYRAMIDAL
- 5 bonding pairs and 1 lone pair
- bond angle = 90*
SQUARE PLANAR
- 4 bonding pairs and 2 lone pairs
- bond angle = 90*
Polarisation
occurs because of the nature of different atomic nuclei
electronegativity
the ability to attract the bonding electrons in a covalent bond
measured on the Pauling scale
a higher number means an element is more electronegative
fluorine is the most electronegative
Non-polar bonds
covalent bonds in diatomic gases are non-polar because the atoms have equal electronegativities so the electrons are equally attracted to the nuclei
Polar bonds
in a covalent bond between 2 atoms of different electronegativities, the bonding electrons are pulled towards the more electronegative atom. This makes he bond polar.
the greater the difference in electronegativity, the more polar the bond.
the difference n electronegativity between the 2 atoms causes a dipole.
what is a dipole?
a difference in charge between the 2 atoms in a covalent bond caused by a shift in electron density in the bond.
Polar molecules
molecules that have a permanent dipole.
a molecule has a permanent dipole if charge is distributed unevenly over the whole molecule.
whether or not a molecule is polar depends on whether it has any polar bonds and its overall shape.
Simple polar molecules
HCl has one polar bond that means charge is distributed unevenly across the whole molecule so it has a permanent dipole.
Complicated molecules - CO2
CO2 has 2 polar bonds that are arranged symmetrically so the dipoles cancel each other out. This means the molecule does not have a permanent dipole and is non-polar.
Complicated polar molecules - H20, CHCl3, CH2F2
the polar bonds are arranged so they all point in roughly the same direction but they do not cancel each other out. This means the charge is arranged unevenly across the whole molecule and this results in a permanent dipole and polar molecule.
what are intermolecular forces?
forces between molecules
much weaker than covalent, ionic or metallic bonds
3 types = induced dipole-dipole (van der Waals forces), permanent dipole-dipole forces and hydrogen bonding.
Van der Waals forces (Induced dipole-dipole)
- cause all atoms and molecules to be attracted to each other
-at any moment, the electrons in an atom are more likely to be at one side. This causes a temporary dipole.
- this dipole can cause another temporary dipole in the opposite direction on a neighbouring atom.
- the 2 dipoles are then attracted to each other.
- the second dipole can then cause another dipole in a third atom.
- it is a domino effects as electrons are constantly moving so dipoles are created and destroyed all the time.
- even though dipoles keep changing, the overall effect is for atoms to be attracted to each other.
- not all van der Waals forces have the same strength - larger molecules have larger electron clouds and therefore stronger forces.
-shape impacts forces - long, straight molecules can lie closer together than branched ones - the coser together the 2 molecules are, the stronger the forces.
permanent dipole-dipole forces
in a substance made up of molecules that have permanent dipoles, there will be weak electrostatic forces of attraction between the opposite charges on neighbouring molecules. These are called permanent dipole-dipole forces.
Hydrogen bonding
- the strongest intermolecular force.
- only happens when hydrogen is covalently bonded to fluorine, nitrogen or oxygen.
- fluorine, nitrogen and oxygen are very electronegative so they draw the bonding electrons away from the hydrogen atom.
- the bond is very polarised and hydrogen has a high charge density as it is so small, that the hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen atoms of any other molecule.
- molecule which have hydrogen bonding are usually organic, including OH and NH groups.
effects of hydrogen bonding
- substances with hydrogen bonding have higher melting and boiling points because extra energy is required to break the hydrogen bonds.
- as water turns into ice, more hydrogen bonds are formed and are arranged into a regular lattice structure.
- ice is less dense than water because the distance between H20 molecules is greater in ice than in water. This is unusual as most substance are more dense as solids than they are as liquids.
behaviour of simple covalent compounds
- have strong covalent bonds within molecules but weak forces between molecules
Electrical conductivity - don’t conduct electricity because there are no free ions or electrons to carry the charge.
Melting point - have low melting points because the weak forces between molecules are easily broken.
Solubility - some dissolve in water depending on how polarised the molecules are.
trends in melting and boiling points - Group 7
- the main factor affecting melting and boiling points is the strength of the van der waals forces, unless the molecule can form hydrogen bonds.
- as you go down group 7, the polarity of molecules decreases, so the strength of the permanent dipole-dipole interactions decreases.
- as you go down group 7, the number of electrons in the molecules increases so the strength of the van der waals forces increases.
- the boiling points increase as you go down because the increasing strength of the van der waals has a greater effect than the decreasing permanent dipole-dipole interactions.
metallic bonding
- the positive metal ions are attracted to the delocalised negative electrons. They form a giant lattice of closely packed positive ions in a sea of delocalised electrons.
properties of metals
Melting point - have high melting points because of the strong electrostatic attraction between positive metal ions and the delocalised electrons. The number of delocaised electrons per atom affects the melting point. E.g: the more there are, the stronger the bonding will be and the higher the melting point. The size of the metal ion and the lattice structure also effect the melting point.
Ability to be shaped - as there are no bonds holding specific ions together, metal ions can slide over each other when the structure is pulled, so metals are malleable and ductile.
Conductivity - delocalised electrons can pass kinetic energy to each other making metals good thermal conductors. The delocalised electrons can also move and carry a charge so metals are good electrical conductors.
Solubility - metals are insoluble, except in liquid metals, because of the strength of the metallic bonds.
melting and boiling covalent substances
- in simple covalent substances, the covalent bonds don’t break during melting and boiling. To melt or boil them, you only have to overcome the weak intermolecular forces that hold the molecules together.This is why they have relatively low melting/boiling points.
To melt or boil a giant covalent substance, you need to break the covalent bonds holding the atoms together. This is why they have very high melting/boiling points.
Physical properties of materials
Melting/Boiling points - determined by the strength of attraction between its particles.
Electrical conductivity - a substance can only conduct electricity if it contains charged particles that are free to move. e.g: delocalised electrons.
Solubility - how soluble a substance is depends on the type of particles that it contains.
summary of ionic bonding properties
- high melting/boiling point
- solid under standard conditions
- only conducts electricity when molten or dissolved
- soluble in water
summary of simple covalent bonding properties
-low melting/boiling points
- usually gas or liquid under standard conditions
- can’t conduct electricity
- an be soluble in water depending on how polar the molecules are.
summary of giant covalent bonding properties
-high melting/boiling points
- solid under standard conditions
-can’t conduct electricity (expect graphite)
- insoluble in water
summary of metallic bonding properties
- high melting/boiling points
-solid under standard conditions
-can conduct electricity - insoluble in water
