Covalent bonding

Metallic Bond

Attraction between cation lattice and delocalised electrons. The result is an element in a lattice.

Ionic Bond

Attraction between cations and anions. The result is a compound in a lattice.

Covalent Bond

The attraction between the nuclei of the atoms and the shared pair(s) of electrons. The result is either an element (H2) or compound (H2O) in a molecule.

Lone pair(s)

(Also known as non-bonding pair(s)) Any pair(s) of valence electrons that are not shared between the two atoms

Bonding pair(s)

Any pair(s) of valence electrons that are shared between the two atoms

Single covalent bond

Is formed when one bonding pair of electrons is shared between two atoms

Double/Triple covalent bond

is formed when two/three bonding pair of electrons is shared between two atoms

Lattice

a repeating structure of atoms arranging in a regular pattern.

Molecule

a group of two or more atoms covalently bonded together.

Matter

General term for anything

Materials

Well defined samples of stuff, retaining from matter

Mixtures

Several substances mixed together, retaining from materials

Substances

chemically pure samples containing one type of basic unit, retaining from materials

Element

A substance containing atoms/nuclei of the same number of protons, retaining from substances

Compound

A substance that contains more than one type of element which combined in a fixed ratio, retaining from substance

Metallic lattice

Cation lattice in a sea of delocalised electrons, retaining from element

Molecules

a group of two or more non-metal atoms covalently bonded together to form a discrete unit, retaining from element/compound

Ionic lattice

Containing atoms of metals and non-metals, retaining from compounds

Electron dot diagram

Represents the atoms using their element symbols and the valence electrons in pairs as dots around the element symbols.

• Show both lone pair(s) and bonding pair(s) of electrons.

Valence structure

Represents the atoms using their element symbols but clearly distinguishes between lone and bonding pairs:

• Lone pairs are dots.

• Bonding pairs lines (one line per bonding pair).

Structural formula

Represents the atoms using their element symbols but lone pairs of electrons are not shown.

Linear VSEPR

2 Bonds, 0 Lone pairs

Bent VSEPR

2 Bonds, 2 Lone pairs

Trigonal Planar VSEPR

3 Bonds, 0 Lone pairs

Trigonal Pyramidal VSEPR

3 Bonds, 1 Lone pair

Tetrahedral VSEPR

4 Bonds, 0 Lone pairs

Non-polar covalent bond

2 atoms with similar or identical electronegativity (>0.4) leads to the formation of a non-polar covalent bond (electrons shared equally)

Ionic bond vs Polar Covalent bond

If the difference in electronegativity is 1.7 and higher the electrons completely transfer to the more electronegative element creating an ionic bond.

If the difference is between 0.4 and 1.7 then a polar covalent bond will form and electrons are distributed unequally (partial sharing)

Polarity of F2

Since the two atoms are the same, there is no difference in electronegativity between the atoms.

The pair of electrons is shared evenly between the atoms.

F2 is non-polar.

Polarity of CH4

Since carbon and hydrogen have a very small difference in electronegativity, the C-H bonds are considered to be non-polar.

Since the bonds are non-polar, CH4 is non-polar.

Polarity of HF

There is one bond dipole that is not cancelled out.

The molecule will have a permanent net dipole and is polar.

The hydrogen atom is considered to have a partial positive charge (a positive dipole)

The fluorine atom is considered to have a partial negative charge (a negative dipole)

Polarity of H2O

The molecule has two bond dipoles (due to the polar bonds) that do not cancel each other out due to its V-Shape.

It would have a net dipole and is polar.

The oxygen atom is the negative dipole and has a partial negative charge

The hydrogen atoms are the positive dipoles and has a partial positive charge

Polarity of NH3

The molecule has three bond dipoles (due to its three polar bonds) that do not cancel each other out due to its trigonal pyramidal shape.

It would have a net dipole and is polar.

The nitrogen atom is the negative dipole and has a partial negative charge

The hydrogen atoms are the positive dipoles and have a partial positive charge

Polarity of CF4

Each C-F bond is polar.

The molecular geometry is tetrahedral so the bond dipoles are evenly spread out around the carbon atom.

This symmetry allows the polar bonds to cancel out, making the molecule non-polar.

The molecule has a net dipole of zero.

Polarity of CO2

Each C=O bond is polar.

The molecular geometry is linear so the two bond dipoles are pointing in opposite directions.

This symmetry of the molecule allows the polar bonds to cancel out, making the molecule non-polar.

The molecule has a net dipole of zero.

Ionic charge

• Due to the giving and taking of electrons.

• Overall, ions have charges.

• Much stronger than partial charge

  (Na+ and F-)

Partial charge/Dipole

• Due to the uneven distribution of electrons in an atom.

• Will always add up to zero. Overall, the molecule has no charge.

  (H :F)

Dipole-Dipole Interactions

• Is the attraction between the positive dipole of a polar covalent molecule and the negative dipole of another polar covalent molecule.

• While a solid line (—) is used to represent covalent bonds, the dashed line (----) is sued to represent dipole-dipole interactions since they are much weaker.

• Dipole-dipole interactions can be formed between any two polar molecules.

Hydrogen bonding

• When N, O, or F (greatly electronegative) are bonded to H (not really electronegative), the resulting bond is highly polar, creating a positive dipole that is highly positive and a negative dipole that is highly negative.

• They are significantly stronger than dipole-dipole interactions.

Temporary dipole

When electrons gather more closely together at one end of the molecule causing one end of the molecule to become negative and the other end to become positive.

Dispersion forces

Is the attraction between the temporary positive dipole of one molecule and the temporary negative dipole of another molecule.

All molecules have dispersion forces. However, they are masked by stronger intermolecular forces.

HBr has a higher boiling point than HCl

Because the more electrons a molecule has the more stronger/significant its dispersion forces. Greater electronegativity doesn’t compare to more electrons an element has (Cl - Br)

Intramolecular bonds

A force that holds the atoms within a molecule or a compound

E.g. covalent bonds, metallic bonds, ionic bonds

Intermolecular forces

The electrostatic forces of attraction between molecules.

They are generally weaker than intramolecular bonds and are eaier to break.

There are three types: dispersion forces, dipole-dipole and hydrogen bonds

Physical properties of substances

A property that can be observed and measured without changing the composition of the subject.

E.g. State at 25 degrees celsius, boiling temp, solubility in water or oil

States of matter: Solid, Liquid and Gas

The state of a substance is a reflection of the distance between the particles in the substance.

Variation in states of matter for covalent molecules

• Is a a reflection of the strength of the intermolecular forces between molecules.

• The stronger the intermolecular forces, the more energy it would require to break them, the harder it is to change state (i.e. higher melting temperature and higher boiling temperature)

Solid lattice vs Solid molecules

• Melting of an ionic solid required the breaking of strong ionic bonds (The melting temperature of solid NaCl is 801 degrees celsius)

• Melting a covalent solid requires the breaking of intermolecular forces (The melting temperature of solid H2O is 0 degrees celsius)