Topic 4 Bonding

4.1 Ionic Bonding (SL/HL)

Metals and non metals

Metals donate electrons, non-metals accept.

Metals form positive ions, non-metals negative.

To get to a full outer electron shell.

Ions held together by electrostatic forces.

Common examples NaCl, MgO, Na2O, AlF3, Ca3P2

e.g. deduce the formula of (a) magnesium fluoride (b) iron II chloride (c) copper I sulfide.

Ions with more than 1 element – polyatomic ions, e.g. CaCO3, Al(OH)3, Mg(NO3)2, (NH4)2SO4, Ba3(PO4)2, LiHCO3

  • E.g. deduce the formula of (a) sodium nitrate (b) iron III sulfate (c) ammonium carbonate
  • Transition metals are able to form more than one type of ion. E.g. FeCl2 and FeCl3.
  • On a larger scale, ions are arranged in a regular giant structure known as a lattice. Each cation is directly connected to an anion with strong electrostatic forces. This results in the crystalline properties of ionic compounds.

4.2 Covalent Bonding (SL/HL)

  • Covalent bonding occurs between a pair of electrons and positively-charged nuclei

Sharing of electron pairs occurs between the nuclei.

Lewis structures (Dot and cross structures) e.g. Cl2, NH3, CCl4.

Multiple bonds. O2, CO2, C2H4, C2H2, SO3.

  • Dative covalent bonds are formed when one atom supplies both the shared electrons to form the covalent bond – e.g. NH4+, H3O+, CO

Bond length and bond strength. Single covalent bonds tend to be longer and weaker than double bonds. Triple bonds are the shortest and strongest. E.g. (CH3COOH)

Ionic or Covalent? (SL/HL)

It can be predicted whether a bond will be ionic or covalent from electronegativity values.

Electronegativity is the ability of an atom to attract electron density towards itself within a covalent bond.

Atoms with large electronegativity differences will have a large separation of charge and are classed as ionic.

Those with a smaller electronegativity difference will have less separation of charge and are covalent.

  • Metals in the bottom left of the periodic table tend to be the least electronegative.
  • Non-metals in the top right of the table tend to the most electronegative.

E.g. K (0.8) and F(4) have a large difference in electronegativity (3.2) and so will be an ionic bond.

Whereas Al (1.5) and Cl (3) have a smaller difference (1.5) and so will be intermediate between an ionic and covalent bond.

When Al bonds with Br(2.8), the electronegativity difference is small enough (1.3) to predict that this bond will now be covalent.

The larger the difference, the more the charge separation, the more ionic the bond is. (a difference greater than 1.8 is mainly ionic, less then 1.8 and it is mainly covalent)

Bond Polarity (SL/HL)

Atoms which are bonded together and have the same or very similar electronegativity values are perfectly covalent, as there is no separation of charge. E.g. N2 or C(2.5) and H(2.2) in CH4.

Other covalent bonds with a larger separation of charge due to different electronegativities are said to be polar (delta + or -).

E.g. H(2.2) and O(3.5) in water is a very polar bond due to the large difference in electronegativity between the atoms (1.2)

In ammonia the bond is less polar between H(2.2) and N(3) with a smaller difference (0.8)

  • In a carbon(2.5) and chlorine(3) bond there is less polarity (0.5)

Some covalently bonded substances may contain polar bonds, but because these are arranged symmetrically to cancel each other out the molecule as a whole is non-polar. E.g. CCl4 and CO2.

Shapes and Bond Angles of Covalent Molecules (SL/HL)

VSEPR Theory

The valence shell electron pair repulsion theory states that electrons will arrange themselves around a central atom as far apart as possible. This is because electrons repel each other.

Electrons may be in a bonded pair, a non-bonded pair, or as part of a multiple bond (double or triple).

These are all called negative charge centres

This theory gives us 5 basic shapes based on the number of negative charge centres.

Negative chargeNameShapeBond Angles

Centres

2Linear1800

3 Triganol Planar1200

4 Tetrahedral109.50

14.2 HL ONLY

5 Triganol bipyramid90, 120

6 Octahedral90

Non-Bonding Electrons and Bond Angles.

 Non-bonding electrons (lone pairs) repel non-bonding electrons more than they repel bonding electrons.

 Bonding electrons repel other bonding electrons by the least.

 This leads to a distortion of the molecules when there are non-bonding electron pairs.

 E.g. In H2O the bond angle is 104.50

 E.g. 2. In NH3 the bond angle is 1070

Multiple Bonds

  • Double and triple bonds count as one negative charge centre.
  • E.g. CO2, SO3, HCN, SO42-

Molecular polarity and Shape

  • The shape of a molecule can also effect if it will be polar.
  • E.g. In CO2 the C-O bonds are polar but since it only has 2 negative charge centres the molecular shape is linear. This means that the molecule overall will be non-polar as the polar bonds will cancel out their effect.
  • The same is true in the tetrahedral CCl4 molecule.
  • However in H2O, the molecule is based on a tetrahedral structure so that the polar bonds cannot cancel, leading to water being a very polar substance.

Structures of the Allotropes of Carbon

An allotrope is when an element can exist with more than one structure.

The following examples of allotropes of carbon need to be studied…

Diamond

Graphite

‘Bucky Balls’ (Buckminsterfullerene)

Diamond

All the 4 bonds formed between carbons are equally strong.

The structure is based upon a tetrahedral to form a macromolecule which is extremely hard to break down.

All of the carbons electrons are involved in bonding. It has no free ‘delocalised’ electrons and so it doesn’t conduct electricity.

Graphite

All the carbon atoms involved form 3 very strong sigma covalent bonds.

The other electron exists in a p-orbital. This can overlap with other p-orbitals to give a much weaker 4th bond.

The strong bonds exist in layers of hexagons, whilst the weak bond exists between these layers.

This explains why graphite can be used as a lubricant. The layers of hexagons can easily be scraped off.

It also explains why graphite is such a good conductor of electricity. These p electrons are able to move between the layers of hexagons. They are delocalised.

Bucky Balls

This is the common name given to a family of spherical carbon molecules arranged in pentagons and hexagons like a football.

The first member of the family to be found was bucksminsterfullerene containing 60 carbon atoms.

They have 3 strong sigma bonds arranged as pentagons and hexagons, which perfectly join up with each other to form the ball shape.

This makes them extremely strong yet lightweight.

An unbonded electron exists in a p orbital. This allows them to conduct electricity but not as well as the layered graphite structure.

Structure and bonding in silicon and silicon dioxide

  • Silicon’s atoms are arranged in the same way as the diamond structure, but does not show allotropic behaviour (ie no graphite or bucky-ball equivalent). Covalent bonding predominates throughout.
  • Silicon dioxide contains silicon atoms bonded tetrahedrally to four oxygen atoms. The empirical formula is given as SiO2, because each O atom is shared equally by 2 Si atoms. It exists in a variety of forms (e.g. glass, sand, quartz) but all are covalently bonded giant structures.

4.3 Intermolecular Forces (SL/HL)

  • These are forces betweencovalent molecules, rather than the covalent bond within them.
Hydrogen Bonding
  • This is the strongest type of intermolecular force. (about 1/10th of a covalent bond.)
  • H-bonds occur when a highly electronegative element (F, N, or O) are bonded to a hydrogen atom.
  • The electronegative element must also contain a non-bonded electron pair.
  • The large polarisation of the bond attracts much of the electron density away from the H-atom.
  • This H-atom experiences a large force of attraction towards the non-bonded electrons on another nearby electronegative atom.
  • This is a Hydrogen bond.
  • HF, H2O, and NH3 experience this type of bonding between the molecules. (Also some organic compounds in topic 11, CH3OH, CH3COOH, CH3NH2)

The formation of hydrogen bonds has no effect on the chemical properties of the molecules, but can have a massive impact on the physical properties such as solubility and melting/boiling points.

  • This is particularly true for H2O since it is able to form 2 hydrogen bonds per molecule in a tetrahedral structure (ice).

Dipole-Dipole Forces

  • Molecules that contain atoms of differing electronegativity will have polar bonds.
  • This results in a one atom being slightly positive and the other being slightly negative.
  • These molecules will have an attractive force between them called a dipole-dipole bond.
  • These bonds are much weaker than Hydrogen bonds.
  • Dipole-dipole bonds lead to an increase in melting and boiling points that is much less than with molecules that are hydrogen bonded.
  • E.g. HCl, CH3Br, CH2=O, PH3, CH3CHO.

Van Der Waals Forces

  • Non-polar molecules experience an even weaker force of attraction between their molecules.
  • When the particles are close enough, the electrons spinning in their orbitals induce a temporary dipole into a neighbouring atom.
  • This causes a weak force of attraction between the two particles.
  • These are the weakest type of intermolecular force.
  • Molecules that only contain this type of force have very low melting and boiling points.
  • E.g. C4H10 = Mr= 58 = Van Der Waals = b.p. at –10C.

C3H6O = Mr = 58 = Dipole-dipole = b.p. at 560C.

  • Since Van Der Waals forces depend on the interaction between the electrons in 2 particles, the greater the molar mass, the greater the interaction between the electrons, the greater the strength of these forces.
  • This explains why the melting points of the halogens increases as the mass increases. (F, Cl, Br, I)
  • The same trend happens with organic molecules such as the alkanes. (CH4, C2H6, C3H8)

4.4 Metallic Bonding

Metallic bonds consist of a giant structure (lattice) of positive ions.

This gives the metal its strength.

In between the ions there is a sea of delocalised valence electrons.

This acts like an electron ‘glue’, allowing the ions to slide over each other.

This is why metals are bendy and ductile.

The fact that these electrons are able to move freely within the lattice structure explains why metals are such good conductors of electricity and heat.

  • The metallic bond exists due to the electrostatic attraction between the positive ions and the delocalized valence electrons.
  • The strength of a metallic bond therefore increases as the number of valence electrons increases and the charge on the metal ion increases.

(This is why metals Na, Mg Al increase in melting point)

4.5 Physical Properties

Melting Points/Boiling Points

  • Covalent macromolecules have the highest melting points. (Diamond, graphite and silicon)
  • Giant structures involved in metallic and ionic bonding have high melting points.
  • Simple covalent molecules have the lowest melting points. (Those with Hydrogen bonding are greatest, then dipole-dipole, then lowest are Van Der Waals)
  • E.g. Propane, ethanal and ethanol all have similar mass, but propane has a much lower b.p. since it only has Van Der Waals forces, ethanal is higher as it has dipole-dipole forces, and ethanol has the highest due to its H-bonding. (Topic 11)
  • Impurities will lower the melting point of all of these types of compounds, since they will interfere with the attractive forces between the particles.

Volatility

  • A volatile liquid will evaporate into a vapour at a low temperature.
  • As the strength of the forces between the particles in the liquid increases, its volatility decreases.
  • Hydrogen bonded liquids such as water or alcohol tend not to be very volatile, but alkanes are very volatile (and hence smell).

Solubilities

  • Water is a polar solvent. Other polar compounds will therefore dissolve in water. (NaCl)
  • Alkanes are non-polar and so other non-polar compounds will dissolve well in them. (iodine in hexane)
  • Alcohols are very good solvents since they have a polar OH group and a non-polar alkane group. They can hence be used to dissolve many polar and non polar compounds.
  • As the length of the alkane group increases, the alcohol becomes more and more non-polar and hence less and less soluble in water. (See TOPIC 11)

Conductivity

  • To conduct electricity a compound must contain charged particles that are able to move.
  • Metals (and graphite) contain delocalised electrons and hence are excellent conductors.
  • Ionic compounds contain charged ions that are only able to move when the compound is molten or dissolved in water.

14.2 Hybridisation(HL ONLY)

Atoms that covalently bond do so by overlapping their orbitals to share electrons.

The problem is that some of the electrons to be shared are in

a) different types of orbitals(s, p, d, f)

b) they often have different shapes(spherical or dumb bell)

c) the electrons it wants to share are already paired up.

Atoms overcome this problem by putting any of its electrons involved in bonding into a bonding orbital.

This bonding orbital is a combination of the types, shapes, and energies of the s, p, d, f orbitals involved and hence is called a hybrid orbital (the process is called hybridisation)

Sp3 Hybridisation

E.g CH4 carbon = 1s2 2s2 2p12p12p0

In order to overlap with each of the s orbitals of hydrogen, the carbon must combine its 2s and its three 2p orbitals into a new set of four bonding orbitals each containing one electron and each in the correct orientation to overlap with the hydrogen orbitals and each of the same energy. Since it involves 1 s and 3 p orbitals this is called sp3 hybridisation

E.g. In CH4 the carbon atom is sp3 hybridised.

In diamond, the carbon atoms are sp3 hybridised.

Sp2 Hybridisation

When an atom needs to bond only 3 times it only combines its outer s orbital with two p orbitals. This is called sp2 hybridisation.

E.g. BF3

In graphite and bucky-balls, sp2 hybridisation is found.

Sp Hybridisation

An atom needing to share twice only needs to create 2 bonding orbitals by combining one s and one p orbital. This is sp hybridisation.

E.g. BeCl2

TOK: Does hybridization really exist? Or is it a mathematical model/prediction to aid understanding? What are the benefits/problems of using such models to help us understand complicated concepts?

Lone Pairs

Any lone pairs of electrons which are in the outer orbital of an atom but are not being used in bonding are still put into bonding orbitals to keep them as far away as possible from the bonded electron pairs.

E.g. NH3

H2O

Sigma and Pi Bonds (HL ONLY)

  • The formation of a bonding orbital by the hybridization of atomic orbitals and then the subsequent ‘end on’ overlapping of these orbitals with those of another atom leads to the formation of a single covalent bond which is called a sigma bond.
  • The electrons are distributed between the 2 atoms.

(Note: Sigma bonds are normally formed from overlap of hybridized bonding orbitals, but in some cases they can also be formed by the overlap of atomic s or p orbitals.)

  • When a double covalent bond forms, it is impossible for 2 sigma bonds to form (VSEPR theory)
  • One sigma bond forms and then there is a’ parallel’ overlapping of non-bonded electrons in a non-hybridized atomicp-orbital in each of the atoms involved.
  • The overlap occurs above and below the sigma bond.
  • This accounts for the shorter length of a double bond, and also a restriction in its ability to rotate.
  • This type of bond is called a pi bond.
  • In a triple covalent bond there are 1 sigma and 2 pi bonds formed by 2 p orbitals on each of the atoms that overlap.

Multiple Bonds and Hybridisation (HL ONLY)

When an atom forms a double or a triple bond with another atom, it only uses one bonding orbital to form a sigma bond.

  • It’s other electrons are kept in its atomic p-orbitals where they can be used to form one or more pi bonds.

E.g. C2H4

E.g. C2H2

E.g. N2H2

Hybridisation and Shape (HL ONLY)

  • It is the type of bonding orbitals formed in a compound which dictates its shape. (VSEPR)
  • Sphybridised molecules are linear
  • E.g. C2H2
  • Sp2 hybridised molecules are triganol.
  • E.g. C2H4
  • Sp3 hybridised molecules are tetrahedral.
  • E.g. CH4
  • The type of hybrid orbitals for the triganol biplanar and octahedral shapes involves s, p and d hybrids.
  • Give the shape and bond angles of the following.
  • NH3, CH2=O, N2, H2O, N2H4, N2H2.

14.3 Delocalisation of Electrons (HL ONLY)

  • When there are a number of atomic p-orbitals next to each other in a covalent molecule.
  • And there are unpaired electrons/charges involved.
  • It will be possible that there will be overlapping of all of these orbitals above and below the sigma bonds in order to stabalise the molecule and spread the charge out.
  • The electrons involved are said to be delocalised (spread) over the atoms concerned.
  • Rather than single and double bonds formed, the bonds are somewhere between.
  • This means that their length and strength is somewhere between that of a double and a single bond.
  • The formation of a delocalised structure in this way leads to a very stable structure.
  • C6H6
  • CH3CO2-
  • NO3-
  • CO32-
  • NO2-
  • O3

TOK: Kekule claimed the inspiration for the model of benzenes structure was from a dream. Is this a rational way of discovery? What is the difference between a scientific rational hypothesis and a non scientific one?