Physical Chemistry - Bonding

What is Ionic bonding?

1. Occurs between metals and non-metals.

2. Electrons are transferred from metal atoms to non-metal atoms.

3. Positive and negative ions are formed.

4. The atom that loses an electron becomes positively charged ion, the atom that gains the electron becomes negatively charged ion.

5. The two ions are attracted to each other and to other oppositely charged ions in the sodium chloride compound by electrostatic forces in all directions.

6. Ionic compounds always exist in a structure called a lattice (giant lattice).

What are the physical properties of ionic bonds?

1. Ionic compounds are always solids at room temperature. They have giant structures and therefore high melting temperatures. This is because in order to melt an ionic compound, energy must be supplied to break up the lattice of ions.

2. Ionic compounds conduct electricity when molten or dissolved in water, but not when solid. This is because the ions that carry the current are free to move in the liquid state but are not free in the solid state.

3. Ionic compounds are brittle and shatter easily when given a sharp blow. This is because they form a lattice of alternative positive and negative ions. A blow in the direction shown may move the ions and produce contact between ions with like charges.

4. Ionic compounds are polar, so are able to dissolve in water.

What is covalent bonding?

1. Between non-metals

2. Electrons are shared so that each atom has a stable noble gas arrangement.

3. A covalent bond is a shared pair of electrons.

How does sharing electrons hold atoms together in covalent bonds?

1. Atoms with covalent bonds are held together by the electrostatic attraction between the nuclei and the shared electrons.

What are double covalent bonds?

1. In a double bond, four electrons are shared. X=X

What are dative covalent bonds?

A single covalent bond consists of a pair of electrons shared between two atoms. In most covalent bonds, each atom provides one of the electrons. But, in some bonds, one atom provides both the electrons. This is called co-ordinate bonding or dative covalent bonding.

1. The atom that accepts the electron pair is an atom that does not have a filled outer main level of electrons - the atom is electron-deficient.

2. The atom that is donating the electrons has a pair of electrons that is not being used in a bond called a lone pair.

What can be used as a term for the measurement of a covalent bond strength.

Average bond enthalpy.

the higher the average bond enthalpy, the stronger the covalent bond.

What is a metallic bond?

1. A lattice of positive ions existing in a ‘sea’ of outer electrons. The electrons are delocalised.

2. The bond is held together by strong electrostatic forces of attraction between the delocalised electrons and the positive ions.

What are the properties of metals?

1. Good conductors of electricity as the delocalised electrons are free to move so can carry a charge throughout the structure.

2. Good conductors of heat - they have high thermal conductivities. The sea of electrons is partly responsible for this property, with energy also spread by increasingly vigorous vibrations of the closely packed ions.

3. Strength:

   1. The greater the charge on the ion, the greater the number of delocalised electrons and the stronger the electrostatic attraction between the positive ions and the electrons.

   2. The smaller the ion, the closer the electrons are to the positive nucleus and the stronger the bond.

4. Metals are malleable and ductile.

5. Metals generally have high melting and boiling points because they have giant structures. The strong attraction between metal ions and the delocalised sea of electrons. This makes the atoms difficult to seperate.

What is electron pair repulsion theory?

1. Each pair of electrons around an atom will repel all other electron pairs.

2. The pairs of electrons will therefore take up positions as far apart as possible to minimise repulsion.

What is the shape of a molecule if there are two pairs of electrons?

1. Linear - The furthest away from each other the two pairs can get is 180° apart.

What is the shape of a molecule of there are three pairs of electrons?

1. Trigonal Planar - 120°

What is the shape of a molecule if there are four pairs of electrons?

Tetrahedron - 109.5°

What is the shape of a molecule with 5 pairs of electrons?

Trigonal bipyramid

What is the shape of a molecule with 6 pairs of electrons?

Octahedral - 90°

What is the shape of a molecule with 3 bonding pairs and 1 lone pair?

Triangular pyramid - 107°

What is the shape of a molecule with 2 bonding pairs and 2 lone pairs?

Non - Linear - 104.5°

What is electronegativity?

1. Electronegativity is the power of an atom to attract the electron density in a covalent bond towards itself.

2. The Pauling scale is used as a measure of electronegativity.

What factors impact electronegativity?

1. Nuclear charge.

2. Distance between the nucleus and the outer shell electrons.

3. The shielding of the nuclear charge by electrons in inner shells.

4. The smaller the atom, the closer the nucleus is to the shared outer main level electrons and the greater its electronegativity.

5. The larger the nuclear charge (for a given shielding effect), the greater the electronegativity.

Describe the trends in electronegativity.

1. Going up a group in the periodic table, electronegativity increases and there is less shielding by electrons in inner shells.

2. Going across a period in the periodic table, the electronegativity increases. The nuclear charge increases, the number of inner main levels remain the same and the atoms become smaller.

What is polarity?

Polarity is about the unequal sharing of the electrons between atoms that are bonded together covalently.

Polarity in covalent bonds between two atoms of the same element.

If both atoms are the same, electronegativity is the same so the bond is non-polar.

Polarity in covalent bonds between two atoms that are different.

If electronegativity is different, then the electrons will not be shared equally so the bond will be polar.

What are the three different types of intermolecular forces?

1. van der Waals forces - act between all atoms and molecules

2. Dipole-dipole forces - act only between certain types of molecules

3. Hydrogen bonding - acts only between certain types of molecules

What are Dipole-dipole forces?

Dipole moments:

1. Molecules with polar bonds may have a dipole moment. This sums up the effect of the polarity of all the bonds in the molecule.

2. In molecules with more than one polar bond, the effects of each bond may cancel, leaving a molecule with no dipole moment. The effects may also add up and so reinforce each other. It depends on the shape of the molecule.

Dipole-dipole forces act between molecules that have permanent dipoles.

What are van der Waals forces? How do they work?

1. All atoms and molecules are made up of positive and negative charges. These charges produce very weak electrostatic attractions between all atoms and molecules.

2. Distribution of charge is changing every instant as electrons are moving.

3. The dipole is caused by the changing position of the electron cloud, so the more electrons there are, the larger the instantaneous dipole will be.

4. The size of van der Waals forces increases with number of electrons present.

What is a hydrogen bond?

1. A hydrogen atom ‘sandwiched’ between two very electronegative atoms.

2. You need a very electronegative atom with a lone pair of electrons covalently bonded to a hydrogen atom.

3. Hydrogen bonds are considerably stronger than dipole-dipole attractions, though much weaker than a covalent bond. They are usually represented by dashes - - -.

When do hydrogen bonds form?

In order to form a hydrogen bond there must be the following:

1. A hydrogen atom that is bonded to a very electronegative atom. This will produce a strong partial positive charge on the hydrogen atom.

2. A very electronegative atom with a lone pair of electrons. These will be attracted to the partially charge hydrogen atom in another molecule and form the bond.

With which atoms or elements do hydrogen bonds form with?

1. Oxygen

2. Nitrogen

3. Flourine

The nitrogen-hydrogen-oxygen system is linear. This is because the pair of electrons in the N—H bond repels those in the hydrogen bond between nitrogen and hydrogen. This linearity is always the case with hydrogen bonds.

Explain the structure and density of ice.

1. In water in its liquid state, the hydrogen bonds break and reform easily as the molecules are moving about. When water freezes, the water molecules are no longer free to move about and the hydrogen bonds hold the molecules in fixed positions.

2. In order to if into this structure, the molecules are slightly less closely packed than in liquid water. This means that ice is less dense than water and forms on top of ponds rather than at the bottom. This insulates the ponds and enables fish to survive through the winter. This must have helped life to continue, in the relative warmth of the water under the ice, during the Ice Ages.

What are the anomalous properties of hydrogen bonding?

1. The density of ice compared to water

2. They have a relatively high melting and boiling points because they are strong intermolecular forces (IMFs) that require significant energy (heat) to overcome

What is a simple molecular lattice and how are the particles bonded and attracted?

• A simple molecular lattice is a solid made of small covalently bonded molecules.

• Strong covalent bonds hold atoms together within each molecule.

• Weak intermolecular forces (IMFs) — mainly London dispersion forces, sometimes permanent dipole–dipole or hydrogen bonding — hold molecules together in the solid.

• The lattice structure depends on the shape and polarity of the molecules; stronger IMFs give a slightly higher melting/boiling point.

How do structure and bonding affect the physical properties (mp/bp, solubility, electrical conductivity) of simple molecular lattices?

• Melting/boiling points:

  • Low because IMFs between molecules are weak and easy to overcome

  • Covalent bonds within molecules stay intact when melting/boiling

• Solubility:

  • Often soluble in non-polar solvents (like dissolves like)

  • Often insoluble in water unless they form hydrogen bonds or are polar

• Electrical conductivity:

  • Do not conduct in solid or liquid states because they have no mobile charged particles (no ions, no delocalised electrons)