Chemistry Revision Notes (Chapter 3 – Bonding)

Hopefully these notes make sense to you. I was attempting to summarize the entire bonding chapter for Chemistry HL/SL.

Chemical bond – interaction between atoms/ions that results in a reduction in the potential energy of the system, making it more stable

Bond type (ionic, covalent, metallic) depends on electronegativities

Very different EN – ionic (EN difference > 1.8)

Metal loses electrons to become cation; non-metal becomes anion; they do not share electrons

Presence of electrostatic attraction between ions which holds them together

Isoelectronic with noble gases (same electronic structure)

3D arrangement (crystal lattice) of anions/cations results in very strong electrostatic attractions, resulting in high melting point and stability of ionic solids

Transition metals: +2 charge most common

-Hard brittle crystalline solids

-Relatively high melting and boiling points

-Non volatile

-If one layer is shifted, then that causes the following atoms to repel each other, causing the substance to break

-firmly held in place so they cannot conduct electricity

• Can conduct electricity when they are not fixed (ie. Molten or in solution)
  

-Do NOT conduct electricity when SOLID, but do when molten/aqueous

-More soluble in water (highly polar) than other solvents (ion-dipole interactions) (the ions react with the dipoles in the water molecule), but insoluble in most solvents

If ion-ion forces too strong, it will not dissolve in water

When naming, don’t forget to include transition state if necessary (III)

High EN – covalent

Acids covalently bond

Usually between non-metals, though metals in high oxidation states can form covalent bonds

Electrons shared because neither element loses electrons easily

When one atom donates both electrons to form a pair of shared electrons, this is called a dative covalent bond

Indicate dative covalent bond by drawing an arrow pointing in the direction that the electron pair is donated (eg from N to H in ammonia)

Dative bonds occur in triple bond in carbon monoxide, in all acid-base interactions, and in all ligand metal ion interactions

Covalent bonds in molecular are very strong, but IMFs are much weaker

- Soft in solid state

- Don’t conduct electricity

- More soluble in non-polar solvents than in water

- Low melting/boiling points

Low EN – metallic

Lewis structures must have all electrons listed whereas structural formulas just have bonds

1. Add up number of valence electrons.
2. Decide which one is the central atom
3. Octet rule
4. Calculate lone pairs
5. If molecule is ion, then draw square brackets around it and put charge in upper right hand corner

Molecular Shapes

Repulsion between a lone pair and a bonding pair is however greater than that between two bonding pairs

Lone pair distorts geometry, causing angle reductions

0 lone pairs, 4 bonds – tetrahedral, 109.5 degrees (eg. CH4)

1 lone pair, 3 bonds – trigonal pyramidal, 107 degrees (eg. NH3 ammonia)

2 lone pairs, 2 bonds – bent, 104 degrees (eg. Water)

Expanded octet:

Elements with d orbitals can promote one or more electrons from a doubly filled s/p orbital into an unfilled low energy d-orbital, which increases # of unpaired electrons available

Usually occurs when the element forms strong polar covalent bonds to small atoms

Non-bonding pairs always occupy the equatorial positions (minimizes electron repulsion)

Page 109, textbook

Polarity

Covalent bonding sharing of electrons is not identical unless the two molecules are identical because of different EN

Stronger EN atom = dipole negative, weaker EN atom = dipole positive

A covalent bond in which the atoms have dipole/fractional charges = polar bond

Illustrating a dipole bond

Can also be illustrated by drawing arrows pointing to the more EN element parallel to the bond line

The EN increases across a period, increases when one goes up a group (with N,O,F being most EN)

The greater the difference in EN, the greater the polarity of the bond

In the extreme case of very large EN differences, like Na and Cl, it is ionic since the electron can be considered to have been transferred

Polar bonds may result in the molecule having a resultant dipole (positive and negative end to the molecule)

However this may be negated by the symmetry of the molecule; for example the centres of positive and negative charge in CO2 neutralize each other since they are in the same place, so there is no overall dipole

Water is not symmetrical, so dipoles do not cancel

THEREFORE, for a molecule to be polar:

Must have polar bonds

Shape must be such that the centres of +/- charges are not in the same place (tetrahedral = nonpolar??)

This can be tested by bringing an electrostatically charged rod close to the jet of liquid in question; polar liquids will be attracted to the stream of liquid, but non-polar will be unaffected

Molecules orientate themselves so that the end closest to the rod has the opposite charge to the rod, so that the electrostatic force of attraction > force of repulsion

Dipole moment – measure of the polarity of a molecule (non-polar = 0 dipole moment, for polar = greater polarity greater dipole moment)

Hybridization

Sigma bond - “end on” interaction of electrons in a s-orbital; electron density greatest on the inter-nuclear axis (an imaginary line joining the two nuclei)

Pi bond – “side on” interaction of orbitals in p-orbitals at right-angles to the aforementioned axis; low electron density on the axis; regions of high electron density on opposite sides, each p-orbital controls one electron or more?

Eg) methane – one of the pair of electrons in the s-orbital is promoted to a vacant p-orbital to produce carbon with 4 Ve; these then form hybrid orbitals

Hybridization – when an atom bonds, the atomic orbitals involved in forming the sigma bonds (or accommodating the lone pairs) interact with each other to form an equal # of highly directional hybrid orbitals of EQUAL energy

Total energy is unchanged, it has just been redistributed equally amongst the orbitals

Ie) outer atomic orbitals interact with each other to produce hybrid orbitals

Involves same amount of orbitals, just rearranged

2 electron pairs (bonds/lone pairs) – sp hybridization (linear, one s and one p)

3 electron pairs – sp2 hybridization (trigonal planar, one s and two p)

4 electron pairs – sp³ hybridization (tetrahedral, one s and three p) (even NH₃ and water but 1 or 2 hybrid orbital has lone pair)

Hybrid orbitals can form sigma bonds, pi bonds are only produced by the sideways interaction of unhybridized p-orbitals

Single, double or triple bonds all count as one bond in hybridization as each only involves one sigma bond

Important for organic

Single bonds – 1 sigma bond

Double bonds – one sigma bond and one pi bond

Triple bonds – one sigma bond and two pi bonds (bond lengths very short, bond energies very high)

Resonance – use double headed arrows

Delocalized bonds

Sigma bonds affect shape along with lone pairs

Pi bonds do not affect molecular shape

Species with unpaired valence electrons are called free radicals

p-orbitals in pi bonds may have differing numbers of electrons and this may cause delocalized pi bonds to form

Delocalization allows pi electrons to spread over more than two nuclei (are mobile and shared by many atoms)

Gives the species a lower potential and makes it more stable than it would if it had only double/single bonds

Eg) benzene – theory predicted alternating double and single bonds, which would not lead to hexagonal shape

Each carbon is sp2 hybridized; one of these orbitals forms a sigma to the hydrogen it is attached to, the other two form sigma bonds to the carbons on either side; remaining 4th electron is in a p-orbital perpendicular to plane of sigma bonds

These p-orbitals interact to produce delocalized pi bond, giving rings of high electron density above and below the ring of carbon atoms

In Lewis structure terms, this is known as resonance structures; the actual species is the resonance hybrid of these extreme structures; same sigma bonds, differ in the arrangement of the pi bonds

Bond order – average number of bonds in different resonance structures (eg benzene due to resonance has bond order of 1.5 because of alternating single and double)

Difference in stability known as delocalization energy/stabilization energy

Hydrogenation of benzene yields less heat energy released then expected because of resonance energy

Graphite (sp2) is similar to benzene in terms of this madness. Can also have delocalized pi bond. Gives C-C bond order of 1.33333.

READ PAGE 120 AND 121.

Intermolecular forces

Covalent bonding can cause giant structure like diamond/silicon dioxide or simple molecular structure like methane

Giant covalent substances are very hard and have very high melting and boiling points because of the strong forces holding the substance together; it follows that the substance is solid at room temp, and insoluble in all solvents

Does not conduct electricity since electrons are held together (with the exception of graphite)

Molecular covalent structures have strong covalent bonds but weak IMFs

Therefore liquids/gases at room temp and pressure

Usually quite soft, will dissolve in non-polar solvents like hexane

Insoluble in very polar solvents like water

Does not conduct electricity either since electrons are firmly held in covalent bonds

Van der Waals’ forces

Exist in all species because temporary dipoles on one molecule from random movement of electrons (especially valence) can have effect on neighbouring molecules

Net result is that the attractive forces between molecules are on average stronger than the repulsive forces, causing temporary attraction

Random accumulation of electrons on one side can cause that side to have a temporary partial negative, repelling the electrons in a neighbouring atom; hence causing temporary dipoles on both sides

Does not last long

Strength of this IMF is proportional to molar mass (increase in electrons and size of instantaneous dipoles)

Therefore, as one goes down group 7 (halogens), molar mass increases, so IMFs increase. Therefore the state of the element changes from gas (fluorine/chlorine) to liquid (bromine) to solid (iodine)

Effective over short range, dependent on surface area of molecules

Eg) elongated molecules have strong van der Waals’ forces, like pentane (straight chain) which is higher than dimethylpropane (almost spherical)

Can become quite strong in plastics/polymer chains

Dipole-dipole forces

Electrostatic attraction between molecules with permanent dipoles

Significantly stronger than van der Waals’ forces

Molecules with this still have van der Waals’ forces

Hydrogen bonding

Hydrogen bonded to N, O or F

Interaction between non bonding electron pair on one of these atoms with hydrogen atom that is carrying a high partial positive charge

Part way between dative covalent and dipole-dipole bond

For maximum strength, the two atoms and hydrogen should be in straight line

Considerably stronger than other IMFs

Basis of base pairing in DNA, alpha/beta sheets in proteins

- Higher boiling point, lower volatility, greater solubility in water, higher viscosity, crystals are harder, more brittle than solids with other IMFs (eg. Sugar)

Hydrides are tetrahedral, non-polar, therefore van der Waals’

Structure of ice:

Tetrahedral because each water molecule can form 4 H-bonds

Similar to diamond structure, H-bonds instead of covalent bonds

Very open structure with large empty spaces, hence low density; thus ponds and lakes freeze from surface down

As water cools, becomes less dense, stays on top until it freezes

Density of liquid water increases when heated from 0 to 4 because structure preserved

Solubility:

If can’t H-bond, then insoluble in water

Propanone can’t bond to propanone, but it can bond to water because of the O-H

Metallic Bonding

Occurs in atoms of low EN

Atoms packed as closely together in 3D (close-packed lattice)

Metals have low ionization energies, many low energy unfilled orbitals, so valence electrons are delocalized (share) amongst the atoms, and the Ve are free to move throughout the metal

Metal atoms become cations in the process

Metallic structure: lattice of cations filled by mobile sea of valence electrons

Attractions between ions and electrons, not ion-ion, so layers of ions can slide past each other because of no need to break bonds in the metal; hence metals are malleable and ductile

If impurities are introduced, the cations slide less easily, so alloys are stronger than pure metals

Metals conduct electricity because delocalized electrons are free to move from one side of the lattice to the other when a potential difference is applied; good conductors of electricity

Mobile electrons also make good conductors of heat and interaction with light causes lustre of metals

Strength of the bond between metal atoms depends on how many electrons each atom shares with the others

Eg) K (one Ve) is 337K, calcium (2 Ve) is 1123K for melting points

Also depends on how far the electrons are from the positive nuclei (ie. Ionic radius)

Eg) sodium compared to lithium has lower melting and boiling points, metallic bonding is weak, substance is soft

Mercury is an exception.

Physical Properties

Depends on IMF. Melting point is also very dependent on existence of regular lattice structure. Impurities disrupt the lattice and lower the melting point.

Hence melting point can be used to check for impurities.

It follows that alloys have lower melting points than metals.

Volatility – how easily the substance is converted to a gas – is also dependent on IMFs

Electrical conductivity depends on whether the substance contains electrically charged particles that are free to move through it when a potential difference is applied.

Dissolving: in order for one substance to dissolve in another the forces between the 2 types of particles must be strong, or stronger than between the particles in the two pure substances

“like dissolves like”

In metals, the hardness, volatility, melting point, boiling point depend on the delocalized electrons

Metals can conduct electricity in all states

Metal atoms cannot form bonds of comparable strength to substances that are ionic or covalent. Metals don’t dissolve in other substances unless they react with them chemically.

Metals can dissolve in other metals to form alloys or even non-metals in some circumstances

Alloys generally less malleable and ductile than pure metal though

Mercury can form a wide range of alloys (amalgams)

Solubility Rules

All nitrates are soluble

All sodium, potassium, ammonium compounds are soluble

All sulfates are soluble except barium sulfate, lead sulfate; calcium sulfate is barely soluble

All chlorides, bromides and iodides are soluble except that of silver

Lead halides soluble in boiling water

All other compounds are insoluble though barium hydroxide and calcium hydroxide are barely soluble

Hydration enthalpy -enthalpy change when one mole of gaseous ions is converted to one mole of hydrated ions is measure of strength of interaction of water molecule with the ion; depends on charge to size ratio

Greatest for small highly charged cations like aluminum

Additional covalent bonding

Diamond is sp3, giant 3D covalent structure; tetrahedral

Exceptionally hard, very high melting and boiling point

Graphite – another allotrope

Allotropes – different forms of an element that exist in same physical state (ozone/oxygen molecules)

Only in 2D, weak van der Waals’ forces between sheets of carbon atoms

Delocalized pi-bond between all of the sp2 hybridized carbon atoms in a given sheet, so bond-order of C-C bonds is 1.33

Carbon-carbon bond length is slightly less than that found in diamond

Distance between sheets is quite large, very few forces between them, hence they can easily slide over each other

Therefore lower density than diamond, can be used as lubricant, “lead” pencils

Delocalized electrons can move between layers so graphite can conduct electricity in 2D

Fullerenes – allotrope of carbon also

Approximately spherical molecules of 5-6 membered carbon rings

C60 – like soccer ball

Sp2 hybridized carbons are bonded by sigma bonds to three other carbons, surface of sphere is not planar so there is little delocalization of unpaired bonding electrons, so electrons cannot easily flow from one C60 molecule to the next

Although higher electrical conductivity than diamond, still much less than graphite/fullerene derivs.

Chemically behaves as electron deficient (accepts electrons from reducing agnets) to form anions

Addition reactions can also occur

Molecular unlike diamond and graphite, so they will dissolve in non-polar solvents and have low melting points

Spherical, have compressibility properties

Related to nanotubes

Silicon has an almost identical crystal structure; can also form 4 covalent bonds

Sideways overlap between p-orbitals of larger atoms is less so other allotropes that involve pi bonding do not occur

Silicon dioxide (quartz) is extremely like diamond, but each carbon is replaced by a silicon and C-C bonds are replaced by an oxygen bridging the silicon atoms

DONE. SUMMARY TABLE P133 TEXTBOOK.