A level Chemistry - Bonding

What is ionic bonding?

Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

• occurs between a metal and a non metal 

• Metal atoms lose electrons to form +ve ions. Non-metal atoms gain electrons to form -ve ions.

• structure: giant ionic structure   

when is ionic bonding stronger?

Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges. E.g. MgO has a higher melting point than NaCl as the ions involved (Mg2+ & O2- are smaller and have higher charges than those in NaCl , Na+ & Cl-

Explain what affects the size of ionic radii

• Positive ions are smaller compared to their atoms because they have one less shell of electrons and the ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely.

• The negative ions formed from groups five to seven are larger than the corresponding atoms.

  The negative ion has more electrons than the corresponding atom but the same number of protons. So the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger.

Within a group the size of the ionic radii increases going down the group. This is because as one goes down the group the ions have more shells of electrons.

What is a covalent bond?

A covalent bond is a shared pair of electrons

What is dative covalent/ coordinate bonding?

A dative covalent bond forms when the
shared pair of electrons in the covalent bond
come from only one of the bonding atoms

What is metallic bonding?

Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons

What factors affect the strength of a metallic bond?

1. Number of protons/ Strength of nuclear attraction.

The more protons the stronger the bond

2. Number of delocalised electrons per atom (the outer shell electrons are delocalised)

The more delocalised electrons the stronger the bond

3. Size of ion.

The smaller the ion, the stronger the bond.

Explain the properties and structure of simple covalent substances

examples: Iodine ,Ice ,Carbon dioxide, Water, Methane

• boiling/melting point : low- because of weak intermolecular forces between molecules (specify type e.g van der waals/hydrogen bond)

• solubility: poor

• conductivity: poor: no ions to conduct and electrons are localised (fixed in place)

• generally: mostly gases and liquids

explain the properties and structure of ionic substances

• Melting/ Boiling point: high- because of giant lattice of ions with strong electrostatic forces between oppositely charged ions.

• soluble in water

• conductivity: only when molten or dissolved, ions are free to move and carry charge through the structure

• generally: crystalline solids

• giant ionic lattice structure , regular and repeating

What are the key macromolecular / giant covalent structures?

Diamond, Graphite , Silicon, Silicon dioxide 

explain the structure and properties of diamond 

• Giant covalent lattice

• Each C bonded to 4 other C atoms in a tetrahedral shape

• Strong covalent bonds throughout the lattice

Properties:

Property	Reason
Very hard	Each atom strongly bonded to 4 others
Very high melting/boiling point	Many strong covalent bonds must be broken
Does not conduct electricity	No free electrons or ions
Insoluble in water	Covalent bonds too strong to break with solvents

explain the structure and properties of graphite 

Structure:

• Giant covalent lattice

• Forms hexagonal layers of graphene 

• Weak van der Waals forces between layers

• Each carbon has one delocalized electron per atom

Properties:

Property	Reason
Soft/slippery	Layers slide over each other (weak forces)
High melting point	Covalent bonds within layers are strong
Conducts electricity	Delocalized electrons can move along layers
Insoluble in water	Covalent bonds too strong

Explain the structure and properties of silicon dioxide

Structure:

• Giant covalent lattice

• Each silicon atom bonded to 4 oxygen atoms (tetrahedral)

• Each oxygen bonded to 2 silicon atoms

• Repeating 3D structure

Properties:

Property	Reason
Very hard	Strong Si–O covalent bonds
Very high melting/boiling point	Many bonds to break
Does not conduct electricity	No free electrons or ions
Insoluble in water	Strong covalent lattice

Explain the structure and properties of silicon

Structure:

• Giant covalent lattice (like diamond)

• Each Si bonded to 4 other Si atoms tetrahedrally

Properties:

Property	Reason
Hard	Covalent bonds throughout lattice
High melting point	Strong covalent bonds
Semi-conductor	Can conduct electricity slightly under certain conditions
Insoluble in water	Covalent bonds too strong

SHAPES OF MOLECULES LEARN ON PAPER :)

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What is electronegativity?

Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.

what are the most electronegative atoms?

F, O, N and Cl are the most electronegative atoms

The most electronegative element is fluorine and it is given a value of 4.0

what is electronegativity measured with?

Electronegativity is measured on the Pauling scale (ranges from 0 to 4)

What factors affect 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

How can differences in electronegativity determine whether compounds are covalent/ ionic bonds?

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

What is a permanent dipole/ polar covalent bond?

• A polar covalent bond forms when the elements in the bond have different electronegativities . (Of around 0.3 to 1.7)

• When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends.

• The element with the larger electronegativity in a polar compound will be the δ- end.

What is a symmetric molecule?

when all bonds identical and no lone pairs

will a symmetric molecule be polar?

No, it will not be polar even if individual bonds within the molecular are polar.

this is because the individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule. There is no net dipole moment: the molecule is non-polar

Describe induced dipole-dipole forces ( Van Der Waals/ dispersion/ London forces )

• They occur between all simple covalent molecules and the separate atoms in 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 i.e. small temporary or transient dipoles form.

  These instantaneous dipoles can cause dipoles to form in neighbouring molecules. These are called induced dipoles. The induced dipole is always the opposite sign to the original one. A temporary dipole forms between the ∂+ in one molecule and the ∂- in a neighbouring molecule.

What factor affects the size of Vaan Der Waals forces? 

The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the Van der Waals stronger between the molecules and so boiling points will be greater.

-The shape of the molecule can also have an effect on the size of the Van der Waals forces. Long chain alkanes have a larger surface area of contact between molecules for Van der Waals to form than compared to spherical shaped branched alkanes and so have stronger Van der Waals.

• this explains increasing boiling points down group 7 and of the alkane homologous series

Describe permeant dipole-dipole forces

•Permanent dipole-dipole forces occurs between polar molecules
•They are stronger than Van der Waals 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.

Permanent dipole-dipole forces occurs in addition to Van der Waals forces

Descrive hydrogen bonding 

• 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

• There is a large electronegativity difference between the H and the O,N,F

• strongest type of dipole-dipole interaction  

• Hydrogen bonding occurs in addition to van der waals forces

can you explain the importance of hydrogen bonding in the low density of ice?

• Ice is solid water (H₂O).

• Each water molecule forms four hydrogen bonds with neighboring molecules.

• This forms a rigid, open, tetrahedral lattice.

Key Points:

1. The lattice holds water molecules further apart than in liquid water.

2. Liquid water has fewer hydrogen bonds and molecules are closer together.

3. Therefore, ice is less dense than water → it floats.

can you explain the importance of hydrogen bonding in the anomalous boiling points of compounds?

• Normally, boiling point increases with molecular mass due to stronger van der Waals forces.

• However, H₂O, NH₃, HF have higher boiling points than expected.

Why:

1. They have hydrogen bonding in addition to van der Waals forces.

2. Extra energy is required to break hydrogen bonds when boiling.

can you explain how melting/boiling points are influenced by intermolecular forces? and how to tell which forces are present in unfamiliar molecules?

Force	Present in	Strength	Effect on BP/MP
Van der Waals	All molecules	Weak	Increases with size
Permanent dipole–dipole	Polar molecules	Medium	Higher BP than non-polar of similar size
Hydrogen bonding	H bonded to N/O/F	Strong	Much higher BP/MP, affects solubility, ice density