Chemical Bonding 

A compound is a substance that is made up of two or more different element combined together chemically. Eg H2O, CO2

Chemical bonds are the attractive force that hold the atom together.

The Octet Rule

Noble gases are very unreactive; they form practically no compounds (also called inert gases)

Octet Rule- when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost energy level.

When elements react together to form compounds, their atoms tend  to change their electron arrangement to try to end up with 8 electrons in their outermost energy level (stable configuration)

Exceptions: 

• Transition metals- in many of their compounds they can have more or fewer than 8 electrons in their outermost energy level.

• Element near helium (hydrogen, lithium, beryllium) tend to achieve the electron arrangement of helium with 2 electrons in the outermost energy level.

Ionic Bonding

An Ion is a charged atom or group of atoms

In a positive ions, since the atom has lost electrons, there are more protons in the nucleus than there are electrons in orbit around it.(also called cations)

In a negative ion the atom has gained electrons (also called anion) 

An Ionic Bond is the force of attraction between oppositely charged ions in a compound. Ionic bonds are always formed by the complete transfer of electrons from on atoms to another.

Writing Ionic Compound

Potassium Bromide

Calcium Chloride

Writing Formulas of Compounds with Group Ions

Sodium Sulfate

Calcium Hydrogencarbonate

Transition Metals

Iron (II) Carbonate

Chromium (ll) Sulfate

Dots and Cross Diagram

Crystal Lattices 

Crystal Lattices is the 3D arrangement of ions

Variable Valency

Iron combines with chlorine to form either FeCl2 or FeCl3
-FeCl2 is called iron (ll) chloride
-FeCl3 is called iron (lll) chloride

Copper combines with oxygen to form either Cu2O or CuO 

-Cu₂O called copper (l) oxide

-CuO is called copper (ll) oxide

Transition metals exhibit variable valency because there is such a small energy difference between the 4s and 3d sublevels. They can lose different number of electrons from these sublevels to give metal ions with different positive charges.

-ide => two elements
-ate => two elements and oxygen

D-elements and Transition Elements

Scandium and Zinc are quite different to those of the other 8 elements in the row hence are not transition metals as they do not have typical transition metal properties.

A Transition metal is one that forms at least one ion with a partially filled 'd' sublevel.

Transition Metal Characteristics:

-Transition metals have variable valency (scandium only forms Sc3+ ion and zinc only forms Zn3+ ions)
-Transition metals usually form coloured compounds (scandium and zinc only form white compounds)
-Transition metals are widely used as catalysts (scandium and zinc show little catalyst activity)
-Neither Sc3+ nor Zn2+ has a partially filled d-sublevel.

Covalent Bonding

Covalent Bonding is the sharing of the pair/s of electrons.

A molecule is a group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.

Valency- of an element is defined as the number of atoms of hydrogens or any other monovalent element with whihc each atom of the elements combines.

Covalent Bonding

Hydrogen Molecule, H ₂

Chlorine Molecule,Cl₂

Water Molecule,H₂O

Ammonia Molecule,NH₃

Methane Molecule,CH₄

Double and Triple Bonds

A double bond is formed when two pairs off electrons are shared between two atoms.

A triple bond is formed when three pairs of electrons are shared.

Sigma and Pi Bonds 

A sigma bonds is formed by the head on overlap of two orbitals 

σ -> is the symbol for sigma

A pi bond is formed by the sideways overlap of p orbitals.

In a double bond there is one sigma and one pi bond.

In a triple bond there is one sigma and two pi bonds.

Sigma bonds are stronger than pi bonds as there is more overlapping of orbitals in sigma bonds.

Ionic	Covalent
Contain a network of ions in crystals	Contain individual molecules
Usually hard and brittle	Usually soft
High melting and boiling points	Low melting and boling points
Uusally solid a room temp	Usually liquid or gases at room temp
Conducts electricity	Doesnt conduct electricity

Shape of Covalent Molecules

The molecules of covalent compounds have particular shapes. VSEPR Theory Valence Shell Electron Pair Repulsion. The shape of a molcules depends on the number of pairs of electrons around the central atoms. Since electrons are negatively charged, the electron pairs repel each other and arrange themselves in space so that they are as far apart as possible.

Lone pair/Lone pair> lone pair/bond pair> bond pair/bond pair because lone pairs are closer to the nucleus, they exert a greater force of repulsion on each other than bond pairs do.

Linear Molecules, Beryllium Chloride BeCl₂

Triangular Planar, Boron Trichloride BCl₃

Tetrahedal, Methane CH₄

Pyramidal, Ammonia NH₃

V-Shaped, Water H₂O

Electronegativity

Electronegativity is the relative atrraction that an atom in a molecule has for the shared pair of electron in a covalent bond.

Since the 'electron pulling power' of chlorine is greater that than of hydrogen in HCL, we say that chlorine is more electronegative than hydrogen.

Pauling set up a scale of relative value of electronegativity.

Elements with low electroonegativity values are siad to be electropositive.

A polar covalent bond is a bond in which there is unequal sharing of the pair(s) of elctrons. This causes one end of the bond to be slightly positive (δ+) and the other end slightly negative (δ–)

A pure covalent is often  used to refer to a covalent bond where there is equal sharing of the two electrons in the bond.

Uses of Electronegativity Values

1. Predict the polarity of Covalent Bonds

- The greater the electronegativity difference, the more polar the bond.

- There are some molcules, which,even though they have polar covalent bonds, are not polar molecules.These are usually symmetrical molecules. The centres of partial charges and centres of partial negative charges coincide, the polairites of the bonds cancel each other out due to symmetry of the molecules.

2. To Predict which compounds are ionic and which are covalent

-An electronegativity difference greater  than 1.7 indicates ionic bonding in a compound.
-An electronegativty difference less than or equal too 1.7 indicates covalent bonding in a compound.
-An electronegativty difference greater than 0.4 and less than 1.7 indicates that the covalent bond is polar.
-An electronegativity difference less than/equal to 0.4 is a non polar covalent bond

Dissolving of Ionic Compounds in Water

Water is an excellent solvent. The property depends on the fact that water is a polar molecule. Most ionic and most polar covalent substances dissolve in water. The ionic bonding ( in NaCl) is overcome by the strong attraction between the ions and the polar water molecules.

Intramoleuclars+Intermolecular Bonding

Intramolcular Bonding is the bonding that takes place within a  moleucle.

Intermolecular forces are the forces of attraction that exist between molecules.
Intramolecular are stronger than intermolecular.

Types Of Intermolecular forces

Van Der Waals Forces

V.D.W Forces are weak attractive forces between molecules resulting from the formation of temporary dipoles. They are the only forces of attractions between non-polar molecules.

Imagine, two electrons moving inside a moleucles of hydrogen. As the 2 electrons move inside the molecule, it may happen that, at any one instant, both electrons may be closer to one end than the other.

A temporary dipole is set up in the molecule ( one δ+ and the other δ– ) only exist for a short time, hence are weak.

Only force of attraction that exist between non polar molecules  eg. H₂, Cl₂, O₂ and also exist in the case of noble gases

The strength of V.D.W forces increase as the molecules get bigger because the increase in the number of electrons means bigger electron cloud, allowing the temporary dipoles to form more easily.

The progression of boiling point is due to the increasing strength of the van der waals forces as the relative molecular masses increase.

E.g Butane has a higher boiling point because it is a bigger molecule than propane and has more electrons and a bigger electron cloud than propane.

Dipole-Dipole Forces

Dipole-Dipole forces are forces of attrction between the negative pole of one polar molecule and the positive pole of another polar molecule.

Dipole-Dipole forces exist permanently

Since Dipole-Dipole forces are between partial electric charges ( δ+ and δ–) these forces are much weaker than the ionic bonds holding ionic compounds together but are  much stronger than V.D.W forces.

The Dipole-Dipole forces between the molecules give rise to higher boilng points than those of similar non-polar molecules. E.G boiling point of HCL (-85℃) is much higher than that of H ₂ (-253℃) (as V.D.W weaker) as more heat must be supplied to  overcome the Dipole-Dipole forces in HCL.

Hydrogen Bonding

Hydrogen Bonds are particular types of dipole-dipole attraction between molecules in which a small highly electromagnetive element such as hydrogen atoms are bonded to Nitrogen, Oxygen or Fluorine.

The Hydrogen atom carries a partial positive charge and is attracted to the electronegative atom in another molecule.

The high boilng points in H₂O,NH₃ and HF are caused by the fact that considerable energy must be supplied to break the hydrgen bonds.

Hydrogen bonds are much stronger than both V.D.Ws and Dipole-Dipole forces however much weaker than ordinary covalent bonds.