Chemistry - Chapter 4 Bonding

What are the 3 main types of bonding and how are they formed?

Why does this bonding happen in some of the types of bonding?

State the strength of these bonds.

What are the 3 main types of bonding and how are they formed?

Why does this bonding happen in some of the types of bonding?

State the strength of these bonds.

1. Ionic (the electrostatic force of attraction between anions and cations in a crystal latttice)

2. Covalent (and dative/coordinate)

   When the outer electrons of 2 atoms are shard.

3. Metallic (the electrostatic force of attraction between positive metal ions and the delocalised electrons)

Ionic and covalent bonding results in the bonding atoms to gain an octet of electrons (the electronic configuration of the nearest noble gas).

All 3 types of bonding is very strons.

Define lattice.

The arrangement of atoms/ions/molecules in a regular repeating pattern in the 3rd dimension throughout the whole crystal structure.

Define Van der Waals’ forces.

What are the different types of Van der Waals’ forces?

Weak forces of attraction between molecules, this could be id-id, pd-pd (including hydrogen bonding). It is the term that covers all types of intermolecular forces.

• instantaneous dipole-induced dipole (also called London dispersion)

• permanent dipole-permanent dipole (including hydrogen bonding, which is stronger)

Define ionic bonding.

Explain how ionic bonding happens (also state another term for ionic bonding).

The electrostatic force of attraction between oppositely charged ions.

• Atoms form ions (by the loss or gain of electrons).

• Metals usually form +vely charged ions

• Non-metals form usually -evely charged ions.

• Metals lose their outer electrons to the outer shell of non-metals, for both to obtain an octet of electrons.

• This causes both to become charged - they become oppositely charged, and since opposite charges attract, they have a strong electrostatic force of attraction between them.

Another term for ionic bonding is electrovalent bonding.

How are ions arranged in an ionic structure?

Ions are arranged in a regular repeating pattern, where an ion is held in place by strong electrostatic forces of attraction between itself and the oppositely charged ions surrounding it.

It is a very strong bond, because the force of electrostatic attraction is very great.

Define covalent bonding.

The electrostatic force of attraction between the 2 nuclei of atoms and the shared pair of electrons.

How does covalent bonding happen. And what electrons are involved in bonding? (3 mark question)

When 2 non-metal atoms combine, they share 1/more pairs of electrons. Not all electrons are involved in bonding:

• Bonding pair

• Lone pair (pairs of electrons in the outer shell of atoms not involved in covalent bonding).

Are there always an octet of electrons when bonding takes place? Explain.

No.

In the case of covalent bonding sometimes an atom that is covalently bonded may not form an octet of electrons:

• they could be ‘electron deficient’ (less than the octet of electrons - less than usual share of electrons)

• they could have an ‘expanded octet’, where there are more than the octet of electrons.

Expanded octets can happen because the extra electrons could fill unfilled p/d orbitals (in principal quantum shell 3).

Define coordinate/dative bonding. Expand on it by stating the conditions required and using at least 2 examples.

Dative bonding is the sharing of a pair of electrons between 2 atoms, where both the electrons come from the same atom.

The conditions required are:

• At least 1 lone pair of electrons in 1 atom.

• The other atom must be electron deficient.

Examples:

1. The formation of the ammonium ion NH₄⁺

   • This happens when an ammonia molecule combines with a H⁺ ion

   • The H⁺ ion is e^- deficient (can have 2 e^- in shell).

   • The N atom in ammonia has a lone pair of e^-s.

2. Aluminium Chloride

   • At high temperatures, it exists as AlCl₃ - electron deficient

   • At lower temperatures AlCl₃ combines to form Al₂Cl₆

   • They can combine due to the lone pair of electrons present in the Cl atoms.

   • 2 Cl atoms form dative bonds with the 2 Al atoms.

The difference in bond length between double and single bonds. Explain.

• Double bonds are shorter.

• They have a greater quantity of negative charge between the 2 nuclei.

• There is therefore a greater electrostatic force of attraction between the nuclei and the 2 pairs of electrons. This causes the nuclei to be pulled closer together.

• Double bonds are therefore more stable.

Define bond energy and bond length.

• Bond energy

  The energy required to break 1 mole of a particular covalent bond in it’s gaseous state. The units are kJmol^-

• Bond length

  The distance between the nuclei of 2 covalently bonded atoms (this depends on the 2 atoms that form the bond).

How does bond strength affect the reactivity of a molecule.

For a reaction to take place, molecules must collide (as in fluids where particles can freely move) with the appropriate amount of energy to break the bonds of one/both the molecules.

It takes a lot of energy to break triple bonds (explaining why N is so unreactive) compared to double bonds (explaining why O₂ is far more reactive).

NOTE: bond strength isn’t the only factor that influences reactivity (polarity, sigma and pi bonds).

State the full form of VSEPR.

Explain what it is about in 5 points.

Valence shell electron pair repulsion theory.

1. Valence shell electrons are the electrons present in the outer shell.

2. Pair of electrons have the same (-ve) charge, and so, repel one another.

3. Repulsion between multiple and single bonds are treated the same as the repulsion between single bonds. But, repulsions between pairs of double bonds are greater.

4. The shape of a molecule can be deduced using this theory. As the most stable shape is the one which minimises the repulsion the most.

What does the shape of the molecule depend on.

And compare the strength of repulsion between bonds. Explain.

• The shape depends on the number of electrons present in the valence shell.

• And which of them are bonded/lone pairs.

Lone pairs of electrons have a more concentrated electron charge cloud. The clouds are more wider and closer to the nucleus of the central atom:

• lone pair-lone pair = most repulsion (lp-lp)

• lone pair-bonded pair = medium (lp-bp)

• bonded pair-bonded pair = least repulsion. (bp-bp)

Give 7 examples of molecules with different shapes and bond angles.

1. Methane

   • Tetrahedral (bond angle = 109.5°)

   • Because all electrons are involved in single bonds. This means equal repulsion.

2. Ammonia

   • Pyramidal (bond angle = 107°)

   • Because there is 1 lone pair present.

     • Since lp-bp causes a larger repulsion than bp-bp;

     • bonding pairs are pushed closer together

3. Water

   • non-linear (bond angle = 104.5°)

   • Because there are 2 lone p[air of electrons.

     • lp-lp causes stronger repulsion than lp-bp and bp-bp;

     • Bonding pairs are pushed even more closer together

4. Bromine trifluoride

   • trigonal planar (120°)

   • No lone pair or electrons, only 3 bonding pairs - equal repulsion).

5. Carbon dioxide

   • Linear (180°)

   • 2 double bonds and no lone pairs (4 electrons in each double bond repels each other in same way as in single bond)

6. Phosphorous pentafluoride

   • trigonal bipyramidal

   • 3 F atoms are in same plane as P, no lone pair of electrons - equal repulsion; 120°

   • The other 2 f atoms are above and below plane of P, and forms a right angle (90°)

7. Sulfur hexafluoride

   • octahedral

   • All single bonds, no lone pairs = equal repulsion - all 90° bonds.

How do covalent bonds form (the 5 mark type)?

• When 2 non-metal atoms combine.

• Each atom that combines has an atomic orbital containing an unpaired electron.

• In covalent bonding these atomic orbitals combine to form a molecular orbital containing 2 electrons.

  • The amount of overlap decides the strength - more overlap means more strong.

• In a shell there are different types of orbitals - usually s and p.

• p orbitals must modify itself to obtain some s character for bonding to take place.

  • Modification causes 1 lobe of p orbital to become slightly bigger.

• s and p orbitals combine to form hybrid orbitals involved in bonding.

  • s can mix with 1/2/3 p orbitals to form sp/sp²/sp³ hybrids.

  • It’s these hybridised orbitals than are involved in bonding (when more than 1 bond is involved)

Explain how sigma and pi bonds are used and where they are found.

• Sigma bonds are formed due to the end-on overlap of orbitals.

  • They are found in single bonds and 1 bond in the double/triple bonds will be a sigma bond.

  • electron density is symmetrical about line joining 2 nuclei

• Pi bonds are formed due to the sideways overlap of p orbitals.

  • One of the bonds in a double bond and 2 of the bonds in a triple bond will be a pi bond.

  • electron density asymmetrical about line joining 2 nuclei

When a molecule uses all 4 orbitals in the 3rd shell, there are 3 possibilities:

• sp³ hybridisation when all are single bonds

• sp² hybridisation when all there are 2 single bonds and 1 double bond. Where, 1 bond is a pi bond.

• sp hybridisation when there is 1 single bond and a triple bond (where 2 are pi bonds)

How are atoms arranged in a metal?

Explain why metallic bonding is strong.

• Atoms are packed together closely and are arranged in a regular repeating pattern in a lattice structure.

• The metal atoms tend to lose their valence electrons to become +ve ions.

• The valence e- occupies new energy levels and are free to move throughout the metal structure (delocalised - not associated with any 1 atom)

• They are held in place by strong electrostatic forces of attraction between the positive metal ions and the negative delocalised sea of electrons.

Metallic bonding is strong due to the strong electrostatic force of attraction between the +ve metal ions and the -vely charge delocalised e-.

What factors increase the strength of metallic bonding?

• increasing +ve charge of the nucleus

• decreasing the size of the metal ions in lattice

• increasing the no. of mobile e- per atom.

Define electronegativity.

Explain the trend of electronegativity across a period and down a group.

State the most common scale of electronegativity values.

The power of a particular atom involved in covalent bonding with another atom to attract the shared (bonding) pair of electrons towards itself.

The greater the electronegativity of an atom, the greater the ability to attract the bonding pair toward itself.

Electronegativity increases across a period and decreases down a group. Br<Cl<N<O<F

Paulings electronegativity values

State the factors which influence electronegativity.

• nuclear charge

  Greater nuclear charge is more likely to attract the bonding pair towards itself

• atomic radius

  Larger atomic radius means outer electrons are further from the nucleus and it is less likely to attract bonding pair (pull on electron pair is lower).

• shielding

  The greater number of full electron shells between the nucleus and valence electrons, the less effective the nuclear charge is on the bonding pair.

Explain polarity and non-polarity.

• When the electronegativity values of 2 atoms are the same/similar, it means the electron pair is equally shared; non-polar.

• When a covalent bond is formed between atoms with diff. electronegativity values, the more electronegative atom pulls the electron pair towards itself.

  • The centre of +ve charge doesn’t coincide with centre of -ve charge.

  • e- distribution asymmetrical.

  • The 2 atoms are partially charged.

• less e-neg = delta +ve

• more e neg = delta -ve

• The is when the bond is polar

What is the degree of polarity and how is it measured?

The degree of polarity is the how polar the bond is (depending on the diff. in e-neg of atoms).

It is measured by the dipole moment.

Define polar bonds.

• Polar bonds:

  Electron pair in bond is drawn towards more electronegative atom, making one end of molecule slightly more +ve than the other.

Define bond polarity

The partial separation of charge when 2 diff. atoms are covalently bonded together. Resulting in the unequal attraction of the bonding pair of electrons.

How does bond polarity influence reactivity, give examples to help explain?

Many reactions start with a reagent attacking a charged end of a polar molecule.

• N₂ and CO both have triple bonds, and so requires approx. the same amount of energy to break them. But, N₂ is non-polar and is therefore very reactive.

  CO is polar, which explains it’s reactivity with O₂ and its use as a reducing agent.

• CH₃Cl is a polar molecule and is prone to attack by reagents like OH^-, as the C in CH₃Cl has a δ⁺ charge due to the polar C-Cl bond.

  • This attack is not possible with ethane, as C-H bonds are virtually non-polar, which also explains why alkanes are unreactive.

Define instantaneous-induced dipole forces.

The weakest intermolecular force. It results from the temporary instantaneous dipole induced in both polar/non-polar molecules (also known as London dispersion forces). These forces are not permanent. The ‘induced’ in this context means one molecule brings out an attractive/repulsive effect on another.

How do id-id forces arise?

Explain why they are temporary?

• Electron charge clouds in molecules (especially non-polar molecules) are constantly coving.

• Often, more of the charge cloud is on one side of the atom than the other.

  • This means for a short time, one side of the atom is more -vely charged than the other side.

• A temporary dipole is therefore set up.

• This dipole induces dipoles in neighbouring molecules.

• As a result there will be attractive forces between the δ⁺ end of the dipole in 1 molecule and the δ^- end of the dipole in the neighbouring molecule.

They are temporary because electron clouds are always moving.

How to increase id-id forces?

1. Increasing the number of protons and electrons in the molecule

   (example: the boiling point and enthalpy change of vaporisation increases down the group of noble gases for this reason)

2. Increasing the number of contact points between molecules

   • Contact points are places where molecules can come closer together.

   (example: pentane and 2,2-dimethyl propane (same no. e-). Molecules of pentane line up alongside each other and has many contact points - higher bp. 2,2-dimethyl propane is more compact and has less surface area for contact - lower bp.

id-id forces between individual atoms is very small, but the total id-id forces between long chains (like polythene) is much larger.

Define permanent dipole- permanent dipole forces.

The forces of attraction between molecules with permanent dipoles.

How do pd-pd forces work?

• Molecules with pds are polar molecules and are always charged with both +ve and -ve ends.

• The δ⁺ on one molecule and the δ^- on a neighbouring molecule causes a weak attractive force between molecules, as there is an attractive force between the opposite charges.

Compare the strength of pd-pd and id-id.

• pd-pd is stronger than id-id between molecules with the same no. of e-.

• more energy required to break pd-pd

• those substances whose molecules are held together by pd-pd have higher boiling points.

Define hydrogen bonding.

The strongest type of intermolecular force, but weaker than covalent bonding. It is a strong pd-pd force.

Conditions required for hydrogen bonding.

Explain how H-bonding works.

• A molecule to have hydrogen atom covalently bonded to F/O/N

• A second molecule to have F/N/O (highly electronegative atom) to have a lone pair of electrons.

• When H is bonded to a highly e-neg atom, the molecule is highly polarised.

• The δ⁺ on the H atom has a high enough charge to form a bond with a lone pair of e- on O/N/F of a neighbouring molecule.