Bonding Flashcards
What is ionic bonding?
Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions formed by electron transfer.
Formation of cation
A positive ion (cation) is formed when an atom loses an electron
Formation of anion
A negative ion (anion) is formed when an atom gains an electron
Migration of ions as evidence for existence of charged particles (ions):
Electrolysis of copper (III) chromate (VI) on a wet filter paper shows evidence for charged particles
When a drop of green copper (III) chromate (VI) is placed on wet filter paper and electricity passed through it, the ions start to separate
The positive Cu2- ions move towards the negative cathode (opposite charges attract). You can see the blue solution move
The negative CrO42- ions move towards the positive anode (opposite charges attract). You can see the yellow solution move
Physical properties of ionic compounds:
Ionic compounds, like sodium chloride, have a giant ionic structure
They conduct electricity when molten or dissolved in solution as the ions are free to move around
Most ionic compounds dissolve in water as water molecules are polar and so they can attract the positive and negative ions and break up the structure
Have high melting points as there are many strong electrostatic forces between oppositely charged ions.
Lots of energy is needed to overcome these forces
They are brittle. When struck with a hammer the layers slide and you get positive ions aligned with positives and negative ions aligned with negatives. They repel, and the structure breaks apart
What are covalent bonds?
Covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them:
Forms between elements that have high electronegativity values (non metals)
There are single, double and triple covalent bonds. More electrons are being shared
How can covalent bonds be represented?
Covalent bonds can be represented by lines too
An arrow represents a coordinate bond
The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient
Reasons for the shapes offend bond angles in, simple molecules
Electron pairs repel and try to get as far apart as possible (or to a position of minimum repulsion.)
If there are no lone pairs the electron pairs repel equally
If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs.
Lone pairs repel more than bonding pairs and so reduce bond angles (by about 2.5o per lone pair
2 bonding pairs
0 lone pairs
Linear shape
Bond angle - 180
3 bonding pairs
0 lone pairs
Trigonal planar shape
Bond angle - 120
4 bonding pairs
0 lone pairs
Tetrahedral shape
Bond angle - 109.5
5 bonding pairs
0 lone pairs
Trigonal bipyramidal shape
Bond angle - 120 and 90
6 bonding pairs
0 lone pairs
Octahedral shape
Bond angle - 90
3 bonding pairs
1 lone pairs
Trigonal pyramidal shape
Bond angle - 107
2 bonding pairs
2 lone pairs
Brent shape
Bond angle 104.5
Factors affecting electronegativity:
Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more.
It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
Ionic and covalent bonding
Ionic and covalent bonding are the extremes of a continuum of bonding type.
Differences in electronegativity between elements can determine where a compound lies on this scale
A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent
A compound containing elements of very different electronegativity and hence a very large electronegativity difference(> 1.7) will be ionic
Why may molecules with polar bonds may not be polar molecules?
If the polar bonds in a molecule are arranged symmetrically, the partial charges cancel out and the molecule is non-polar.
Due to directly opposing pulls by the electronegative element.
London forces
London forces occur between all molecular substances and noble gases. They do not occur in ionic substances
In any molecule the electrons are moving constantly and randomly.
As this happens the electron density can fluctuate, and parts of the molecule become more or less negative. This leads to instantaneous dipoles
Instantaneous dipole induces a dipole in nearby molecules. Induced dipoles attract one another
Main factor affecting size of London Forces
The more electrons there are in the molecule the higher the chance that temporary dipoles will form.
This makes the London forces stronger between the molecules and more energy is needed to break them so boiling points will be greater.
Permanent dipole-dipole forces
- Permanent dipole-dipole forces occurs between polar molecules
- It is stronger than London forces and so the compounds have higher boiling points
- Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
- Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
Hydrogen bonding:
It occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a–O-H -N-H F-H bond.
There is a large electronegativity difference between the H and the O,N,F
Understand the interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bonding:
Hydrogen bonding is stronger than the other two types of intermolecular bonding.
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
How does the shape of the molecule affect the size of London forces?
The shape of the molecule can also have an effect on the size of the London forces.
Long straight chain alkanes have a larger surface area of contact between molecules for London forces to form than compared to spherical shaped branched alkanes and so have stronger London forces
Explain how hydrogen bending results in ice being more dense than water
For most substances, as the temperature drops below its freezing point the density increases as a solid forms. This is not the case for water.
The density of ice changes at around 4*C.
There are 4 groups around each oxygen atom, arranged tetrahedrally in three dimensions by hydrogen bonding.
This leads to an ‘open’ structure with large spaces in it.”
Because of its open structure ice is less dense than water.
Explain how hydrogen bending results in the relatively high melting/ boiling points of water
Each water molecule can potentially form four hydrogen bonds with surrounding water molecules.
There are exactly the right numbers of δ+ hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding.
This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride
Trends in boiling temperatures of alkanes with increasing chain length
The only intermolecular force in alkanes are London forces
These forces are very weak. With more electrons, the forces get stronger, as the instantaneous dipoles are likely to be stronger.
This is why longer chains of alkanes have higher boiling and melting points.
Effect of branching in the carbon chain on the boiling temperatures of alkanes
The only intermolecular force that exist in alkanes are London forces
Straight chain molecules have more places along its length where they can be attracted to other molecules, so there are more chances of London forces to be developed.
Hence, they have stronger intermolecular forces (and higher boiling points) as compared to the branched chain molecules which have a compact shape, and therefore fewer spaces where they can be attracted to other molecules.
The relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
Alcohols are not as volatile as alkanes with similar numbers of electrons, due to hydrogen bonding
Alkanes have weaker London forces
The trends in boiling temperatures of the hydrogen halides, HF to HI
Among hydrogen halides, only hydrogen fluoride exhibits hydrogen bonding between molecules, and therefore has the highest melting and boiling points of the Hydrogen halides.
From HCl to HI the boiling point rises. This trend is attributed to the increasing strength of intermolecular van der Waals forces , which increases as the numbers of electrons in the molecules increases.
For a substance to dissolve:
Solvent bonds must break
Substance bonds must break
New bonds must be formed between the solvent and substance
Ionic substances dissolving in water
When an ionic lattice dissolves in water it involves breaking up the bonds in the lattice and forming new bonds between the metal ions and water molecules.
The negative ions are attracted to the δ+ hydrogens on the polar water molecules and the positive ions are attracted to the δ- oxygen on the polar water molecules.
The water molecules surround the ions in a process called hydrogen
Solubility of simple alcohols
The smaller alcohols are soluble in water because they can form hydrogen bonds with water.
The longer the hydrocarbon chain the less soluble the alcohol.
Insolubility of compounds in water
Compounds that cannot form hydrogen bonds with water molecules ,e.g. polar molecules such as halogenoalkanes or non polar substances like hexane will be insoluble in water.
Solubility in non-aqueous solvents
Compounds which have similar intermolecular forces to those in the solvent will generally dissolve
Non-polar solutes will dissolve in non-polar solvents e.g. Iodine which has only London forces between its molecules will dissolve in a non polar solvent such as hexane which also only has London forces.
Propanone is a useful solvent because it has both polar and nonpolar characteristics. It can form London forces with some non polar substances such as octane with its CH3 groups. Its polar C=O bond can also hydrogen bond with water
What is metallic bonding?
Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
The three main factors that affect the strength of metallic bonding are:
- Number of protons/ Strength of nuclear attraction. The more protons the stronger the bond
- Number of delocalised electrons per atom (the outer shell electrons are delocalised)
The more delocalised electrons the stronger the bond - Size of ion.
The smaller the ion, the stronger the bond.
Properties of metals
Metals have high melting points because the strong electrostatic forces between positive ions and sea of delocalised electrons require a lot of energy to break
Metals can conduct electricity well because the delocalised electrons can move through the structure
Metals are malleable because the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another. The attractive forces in the lattice are the same whichever ions are adjacent
