1.3 Bonding Flashcards
Ions
lons are formed when electrons are transferred from one atom to another.
The simplest ions are single atoms which have either lost or gained 1, 2 or 3 electrons so that they’ve got a full outer shell.
compound ions
ions that are made up of groups of atoms with an overall
charge.
lonic compounds
When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms. The metal atoms lose electrons to become positively charged ions (cations) with a full outer shell of electrons. The non-metal atoms gain electrons and become negatively charged ions (anions) with a full outer shell of electrons.
The oppositely charged ions are strongly attracted to each other, and this strong electrostatic attraction holds the ions together in the ionic compound in a lattice. This is known as ionic bonding.
Giant ionic lattices
lonic crystals are giant lattices of ions. A lattice is just a regular structure. The structure’s called ‘giant’ because it’s made up of the same basic unit
repeated over and over again. In sodium chloride, the Na+ and Cl- ions are packed together. Sodium chloride is an example of a compound with an ionic
crystal structure.
Behaviour of ionic compounds
The structure of ionic compounds decides their physical properties
things like their electrical conductivity, melting point and solubility.
Electrical conductivity
lonic compounds conduct electricity when they’re molten or dissolved but not when they’re solid. The ions in a liquid are free to move (and they carry a charge). In a solid they’re fixed in position by the strong ionic bonds.
Melting point
lonic compounds have high melting points. The giant ionic lattices are held together by strong electrostatic forces. It takes loads of energy to overcome these forces, so melting points are very high (801 ℃ for sodium chloride).
Solubility
lonic compounds tend to dissolve in water. Water molecules are polar - part of the molecule has a small negative charge, and the other bits have small positive charges (see pages 92-93). The water molecules pull the ions away from the lattice and cause it to dissolve.
covalent bonds
A covalent bond is a shared pair of electrons between 2 atoms, so they’ve both got full outer shells of electrons. A single covalent bond contains a shared pair of electrons.
Both the positive nuclei are attracted electrostatically to the shared electrons.
Double bonds
One carbon atom (C) can bond to two oxygen
atoms (O). Each oxygen atom shares two pairs of
electrons with the carbon atom. So, each molecule
of carbon dioxide (CO2) contains two double bonds.
Triple bonds
When a molecule of nitrogen (N2) forms, the nitrogen
atoms share three pairs of electrons. So, each
molecule of nitrogen contains one triple bond.
Simple covalent compounds
Compounds that are made up of lots of individual molecules are called simple covalent compounds. The atoms in the molecules are held together by strong covalent bonds, but the molecules within the simple covalent compound are held together by much weaker intermolecular forces of attraction.
It’s the intermolecular forces, rather than the covalent bonds within the molecules, that determine the properties of simple covalent compounds. In general, they have low melting and boiling points and are electrical insulators
Giant covalent structures
Giant covalent structures are type of crystal structure. They have a huge network of covalently bonded atoms. (They’re sometimes called macromolecular structures). Carbon atoms can form this type of structure because they can each
form four strong, covalent bonds. There are two types of giant covalent carbon structure you need to know about - graphite and diamond.
Graphite
The carbon atoms in graphite are arranged in sheets of flat hexagons covalently bonded with three bonds each. The fourth outer electron of each carbon atom is delocalised. The sheets of hexagons are bonded together by weak van der Waals forces
Graphite properties
Graphite’s structure means it has certain properties:
. The weak bonds between the layers in graphite are easily broken, so the sheets can slide over each other making graphite feels slippery and is used as a dry 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 low 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 sublimes at over 3900 K).
Graphite is insoluble in any solvent. The covalent bonds in the sheets are too difficult to break.
Diamond
Diamond is also made up of carbon atoms. Each carbon atom is covalently bonded to four other carbon atoms . The atoms arrange
themselves in a tetrahedral shape - its crystal lattice structure.
Diamond properties
Because of its strong covalent bonds:
- Diamond has a very high melting point - it actually sublimes at over 3800 K.
- Diamond is extremely hard - it’s used in diamond-tipped drills and saws.
- Vibrations travel easily through the stiff lattice, so it’s a good thermal conductor.
- It can’t conduct electricity - all the outer electrons are held in localised bonds.
- Like graphite, diamond won’t dissolve in any solvent.
Co-ordinate (dative covalent) bonds
In a normal single covalent bond, atoms share a pair of electrons - with one electron coming from each atom. In a co-ordinate bond, also known as a
dative covalent bond, one of the atoms provides both of the shared electrons.
Co-ordinate (dative covalent) bonds eg 2
2 things that affect the strength of an ionic lattice (this affects melting point and solubility)
Size of each ion (as you go down group ionic size increases, as you go across period ionic size decreases, as ionic size decreases bond strength increases)
Strength of the charge on each ion (e.g. 1+, 2+, 1-, 2-), as ionic charge increases bond strength increases
Charge clouds - bonding and lone pairs
The shape depends on the number of pairs of electrons in the outer shell of the central
atom. Pairs of electrons can be shared in a covalent bond or can be unshared.
Shared electrons are called bonding pairs, unshared electrons are called lone pairs or non-bonding pairs.
Bonding pairs and lone pairs of electrons exist as charge clouds.
Charge clouds
A charge cloud is an area where you have a big chance of finding an electron.
The electrons don’t stay still - they whizz around inside the charge cloud.
Valence Shell Electron Pair Repulsion Theory
Electrons are all negatively charged, so charge clouds repel each other until
they’re as far apart as possible. The shape of a charge 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.
Drawing shapes of molecules
It can be tricky to draw molecules showing their shapes, because you’re trying
to show a 3D shape on a 2D page. Usually you do it is by using different
types of lines to show which way the bonds are pointing. In a molecule
diagram, use wedges to show a bond pointing towards you, and a broken
(or dotted) line to show a bond pointing away from you
Finding the number of electron pairs (lone and bonding) e.g.PH3
Molecular shapes - Central atoms with two electron pairs
Molecules with two electron pairs have a bond angle of 180° and have a
linear shape. This is because the pairs of bonding electrons want to be as
far away from each other as possible.
Molecular shapes - Central atoms with three electron pairs
Molecules that have three electron pairs around the central atom don’t always
have the same shape - the shape depends on the combination of bonding
pairs and lone pairs of electrons.
If there are three bonding pairs of electrons the repulsion of the charge clouds is the same between each pair and so the bond angles are all 120°. The shape of the molecule is called trigonal planar.
Central atoms with four electron pairs (4bp)
If there are four pairs of bonding electrons and no lone pairs on a central atom, all the bond angles are 109.5° - the charge clouds all repel each other equally. The shape of the molecule is tetrahedral.
Central atoms with four electron pairs (3BP + 1LP)
If there are three bonding pairs of electrons and one lone pair, the
lone-pair/bonding-pair repulsion will be greater than the bonding-pair/
bonding-pair repulsion and so the angles between the atoms will change.
There’ll be smaller bond angles between the bonding pairs of electrons and
larger bond angles between the lone pair and the bonding pairs. The bond
angle is 107° and the shape of the molecules is trigonal pyramidal.
Central atoms with four electron pairs (2BP + 2LP)
If there are two bonding pairs of electrons and two lone pairs
of electrons the lone-pair/lone-pair repulsion will squish the bond angle
even further. The bond angle will be around 104.5° and the shape of the
molecules is bent (or non-linear).
Central atoms with five electron pairs
Some central atoms can ‘expand the octet’ - which just means that they can have more than eight bonding electrons in their outer shells.
A molecule with five bonding pairs will be trigonal bipyramidal. Repulsion between the bonding pairs means that three of the atoms will form a trigonal planar shape with bond angles of 120° and the other two atoms will be at 90° to them.
Central atoms with five electron pairs (4BP + 1LP)
If there are four bonding pairs and one lone pair of electrons, the
molecule forms a disphenoidal / seesaw shape. The lone pair is always positioned where
one of the trigonal planar atoms would be in a trigonal bipyramidal molecule.
Central atoms with five electron pairs (3BP + 2LP)
If there are three bonding pairs and two lone pairs of electrons,
the molecule will be T-shaped.
Central atoms with six electron pairs
A molecule with six bonding pairs will be octahedral. All of the bond angles
in the molecule will be 90°.
Central atoms with six electron pairs (5BP + 1LP)
If there are five bonding pairs and one lone pair, the molecule forms a
square pyramidal structure. (Molecules with this shape are very rare.)
Central atoms with six electron pairs (4BP + 2LP)
If there are four bonding pairs and two lone pairs of electrons, the
molecule will be square planar.
Finding the shape of an unfamiliar molecule - Example - Predict the shape of the BF - ion.
Awkward molecules - Carbon dioxide
Carbon has four bonding pairs of electrons (found in two carbon-oxygen double bonds) and no lone pairs. Double bonds can be treated as one bond, so you can say that there are two bonds and no lone pairs - CO2 will be linear.
Awkward molecules - Sulfur dioxide
Sulfur has four bonding pairs of electrons (found in two sulfur-oxygen double
bonds) and one lone pair. Double bonds can be treated as one bond so you can
say there that are two bonds and one lone pair - SO2 will be bent (non-linear).
The extra electron density in the double bonds cancels out the extra repulsion from the lone pair, so you get 120° angles.
non-polar bonds
The covalent bonds in diatomic gases (e.g. H2, Cl2) are non-polar because
the atoms have equal electronegativities and so the electrons are equally
attracted to both nuclei (see Figure 2). Some elements, like carbon and
hydrogen, have pretty similar electronegativities, so bonds between them
are essentially non-polar.
Polar bonds
In a covalent bond between two atoms of different electronegativities, the bonding electrons are pulled towards the more electronegative atom. This makes the bond polar. The greater the difference in electronegativity, the more polar the bond.
Polar molecules
If charge is distributed unevenly over a whole molecule, then the molecule will have a permanent dipole. Molecules that have a permanent dipole are called polar molecules. Whether or not a molecule is polar depends on whether it has any polar bonds, and its overall shape.
In simple molecules, such as hydrogen chloride, the one polar bond means charge is distributed unevenly across the whole molecule, so it has a permanent dipole
More complicated non Polar molecules with polare bonds
More complicated molecules might have several polar bonds. The shape of the molecule will decide whether or not it has an overall permanent dipole.
If the polar bonds are arranged symmetrically so that the dipoles cancel each
other out, such as in carbon dioxide, then the molecule has no permanent
dipole and is non-polar
More complicated Polar molecules
If the polar bonds are arranged so that they all point in roughly the same
direction, and they don’t cancel each other out, then charge will be arranged
unevenly across the whole molecule. This results in a polar molecule - the
molecule has a permanent dipole
Metallic bonding
Metal elements exist as giant metallic lattice structures. The outermost shell of electrons of a metal atom is delocalised - the electrons are free to move about the metal. This leaves a positive metal ion, e.g. Na+, Mg2+, Al3+. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons - this is metallic bonding.
Solids
A typical solid has its particles very close together. This gives it a high density and makes it incompressible. The particles vibrate about a fixed point and can’t move about freely.
liquids
A typical liquid has a similar density to a solid and is virtually incompressible. The particles move about freely and randomly within the liquid, allowing it to flow.
gases
In gases, the particles have loads more energy and are much further apart. So the density is generally pretty low and it’s very compressible. The particles move about freely, with not a lot of attraction between them, so they’ll quickly diffuse to fill a container
summary of different properties of different types of bonding
Van der Waals forces
Van der Waals forces cause all atoms and molecules to be attracted to each other. Electrons in charge clouds are always moving really quickly. At any particular moment, the electrons in an atom are likely to be more to one side than the other. At this moment, the atom would have a temporary dipole. This dipole can cause another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to each other. The second dipole can cause yet another dipole in a third atom. It’s kind of like the domino effect. Because the electrons are constantly moving, the dipoles are being created and destroyed all the time. Even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each other.
Factors affecting the relative strength of van der Waals forces (which further affects melting point / viscosity)
Not all van der Waals forces are the same strength - larger molecules have
larger electron clouds, meaning stronger van der Waals forces.
The shape of molecules also affects the strength of van der Waals
forces. Long, straight molecules can lie closer together than branched ones -
the closer together two molecules are, the stronger the forces between them.
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 delta+ and delta- charges on neighbouring molecules. These are called permanent dipole-dipole forces.
polar water attraction to charged rods
If you put an electrostatically charged rod next to a jet of a polar liquid, like water, the liquid will move towards the rod. It’s because polar liquids contain molecules with permanent dipoles. It doesn’t matter if the rod is positively or negatively charged. The polar molecules in the liquid can turn around so the oppositely charged end is attracted towards the rod.
The more polar the liquid, the stronger the electrostatic attraction between the rod and the jet, so the greater the deflection will be. By contrast, liquids made up of non-polar molecules, such as hexane, will not be affected at all when placed near a charged rod.
Hydrogen bonding
Hydrogen bonding is the strongest intermolecular force. It only happens when hydrogen is covalently bonded to nitrogen, oxygen or fluorine (that’s eNOF). Fluorine, nitrogen and oxygen are very electronegative, so they draw the bonding electrons away from the hydrogen atom.
The bond is so polarised, and hydrogen has such a high charge density because it’s so small, that the hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen atoms of other molecules. Molecules which have hydrogen bonding are usually organic, containing -OH or -NH groups.
Hydrogen bonding affect on boiling points
Hydrogen bonding has a huge effect on the properties of substances. Substances with hydrogen bonds have higher boiling and melting points than other similar molecules because of the extra energy needed to break the hydrogen bonds. This is the case with water which has a much higher boiling point than the other group 6 hydrides
Hydrogen bonding affect on water density compared to ice
As liquid water cools to form ice, the molecules make more hydrogen bonds and arrange themselves into a regular lattice structure. Since hydrogen bonds are relatively long, the average distance between H2O molecules is greater in ice than in liquid water - so ice is less dense than liquid water (see Figures 12 and 13). This is unusual - most substances are more dense as solids than they are as liquids.
Trends in melting and boiling points of simple covalent compounds
In general, the main factor that determines the melting or boiling point of a substance will be the strength of the induced dipole-dipole forces (unless the molecule can form hydrogen bonds).
