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CHAPTER 4 : CHEMICAL BONDING - Coggle Diagram
CHAPTER 4 : CHEMICAL BONDING
Type of Bonds and Lewis Structure
Why form chemical bonds
:
Acquire stable noble gas configuration of minimum energy
Intra-molecular bonds bind atoms within a molecule or compound
with each other
Inter-molecular forces bind compounds or molecules together
The attractions between molecules (intermolecular) are not nearly as strong as the intramolecular attractions that hold compounds together.
Type of Bonds
Three principle classes of bonding :
Bonding involves the transfer
ionic bonding
:
Bonding involves the transfer of electrons between atoms
covalent bonding
:
Bonding which involves the sharing of electrons
metallic bonding
:
Metal atoms bonded to other metal atoms
Octet rule
:
Atoms ONLY come together for a single reason:
to produce a more stable electron configuration.
Atoms bond together by either transferring or sharing electrons.
Atoms like to have 8 electrons in their outer shell.
Octet rule (gain, lose or share electrons)
Valence Electron and Lewis Dot Symbol
:
Uses symbol of element to represent nucleus and inner electrons
Uses dots around the symbol to represent valence electrons
Puts one electron on each side first, then pair
elements in the same group have the same number of valence electrons; therefore, their Lewis dot symbols will look alike.
Ionic Bond (Electrocovalent Bond)
Ionic bonds are formed when electrons are transferred from one atom to another
The atom which donates electrons becomes positively charged (cation) and the atom which receives
electrons becomes negatively charged (anion)
The electrostatic attraction between the oppositely charged ions is called the ionic bond
The compounds with ionic bond is called ionic compound.
Metal to nonmetal
Metal loses electrons to form cation (IE )
Nonmetal gains electrons to form anion (EA )
Ionic bond results from + to − attraction.
Larger charge = stronger attraction, LE increases.
Smaller ion = stronger attraction, LE increases.
Lattice energy (LE) released must be high to formed stable ionic structures. LE is the energy released by a process in which isolated ions come together to form a crystal or ionic compound
Lattice energy is the energy required to separate 1 mole of an ionic solid into a gaseous ions
Lattice energy is the measure of the strength of the ionic bond
Factor affecting Lattice energy
Ionic size
as ionic size increases, lattice energy decreases
ionic charge
as ionic charge increases, lattice energy increases
Properties of Ionic Compound
High melting and boiling point
The stronger the attraction force, the higher the melting point and boiling point
Hard and brittle crystalline solid
If the ions are displaced from their position in the crystal lattice, repulsive force will occur
The crystal will become unstable and break apart
ionic solids are brittle. When struct they shatter. Like the charges repel each other
Good electrical conductor in molten and aqueous state
In the ionic solid, the ions are locked in position and cannot move around. Thus, does not conduct electricity.
In molten state or when dissolved in water, there are free moving ions. Thus, conduction electricity
Writing Lewis Structure for Ionic Bonds
:
Draw the Lewis dot symbols of the elements.
Transfer all the valance electrons from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost from the metal atoms and all nonmetal atoms have 8 electrons.
Covalent bonds
Formation of covalent bonds
:
Often found between two nonmetals
Typical of molecular species
Atoms bonded together to form molecules
Strong attraction
Atoms share pairs of electrons to attain octets
Molecules generally weakly attracted to each other
Observed physical properties of molecular substance due to these attractions.
A covalent bond is a chemical bond formed by sharing valence electron between two atoms and held by electrostatic attractions between bonding electrons and nuclei
A shared electron would then be count towards each atom's octet.
Single Covalent Bonds
:
Two atoms share one pair of electrons
One atom may use more than one single bond to fulfill its octet
Double Covalent Bond
:
Two atoms sharing two pairs of electrons
Shorter and stronger than single bond
Triple Covalent Bond
:
Two atoms sharing 3 pairs of electrons
Shorter and stronger than single or double bond
Properties of Covalent Bond
Low melting and boiling point
The covalent bond are strong, but the attraction between the molecules are generally weak intermolecular forces.
Except for giant covalent structure : higher melting point
Do not conduct electricity
:
There are no charged particles around to allow the material to conduct
Solubility
:
In general, covalent compound are insoluble in water
Bonding and Lone Pair Electrons
:
Electrons that are shared by atoms are called bonding pairs
Electrons that are not shared by atoms but belong to a particular atom are called lone pairs
Also known as nonbonding pairs
Drawing Lewis Structure for Covalent Bond
:
Step 1: Count the total number of valence electrons
Step 2: Write the skeletal structure.
Step 3: Use two valence electrons to form each bond.
Step 4: Complete the octets of the outer atoms by distributing the remaining valence electrons in pairs.
Step 5: Place remaining electrons on the central atom
Step 6: If central atom does not form an octet, form double or triple bonds
Attach the atoms together in a skeletal structure.
Halogens and hydrogen are generally terminal
Many molecules tend to be symmetrical
In oxyacids, the acid hydrogens are attached to an oxygen.
Multiple bonds
:
If there are not enough electron for the central atom to attain an octet, a multiple bond is present
Step 5 : If the central atom does not have a full octet, change a lone pair on a surrounding atom into another bonding pair to the central atom, thus forming a multiple bond
Polyatomic Ions
:
The polyatomic ions are attracted to opposite ions by ionic bonds
Form crystal lattices
Atoms in the polyatomic ion are held together by covalent bonds
Formal Charge and Lewis Structure
:
It is the apparent charge on an atom in a molecule or poly-atomic ion
When several Lewis structures are possible, those with smallest formal charges are the most stable and preferred
Formal charge on any atom can be calculated as follows:
Formal charge = (no of valence electrons) – (no of bonds + no. of unshared electrons)
Resonance Structure
:
When there are more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structure
-to draw it, look for multiple bonds or lone pairs
Rules of resonance structure
:
Resonance structures must have the;
same connectivity - only electron positions can change.
same total number of electrons.
Period 2 elements have a maximum of eight electrons.
bonding and nonbonding
third row can have expanded octet.
Total formal charges for neutral molecules must be zero.
Total formal charges for polyatomic ions is the same as ionic charge.
Lewis structures often do not accurately represent the electron distribution in a molecule
Real molecule is a hybrid of all possible Lewis structures.
Resonance stabilizes the molecule.
Drawing Resonance Structures
:
Draw Lewis structure that maximizes octets.
Move electron pairs from outside atoms to share with central atoms
In all resonance structure the arrangement of the nuclei is the same
Most Plausible (stable) state
:
All atoms obey the octet rule.
Zero formal charges or smaller formal charges.
Negative formal charge on the more electronegative atom
Sum of formal charges equal zero or equal the ionic charge for polyatomic ions.
When possible, choose a structure that have different charges on adjacent atoms.
Exceptions to the Octet Rule
:
Incomplete octet
Be forms two bonds with no lone pairs ins compounds
B and Al forms three bonds with no lone pairs in its compounds
Expanded octet
Many elements may end up with more than eight valence electrons in their structure if they can use their empty d orbitals for bonding.
Odd number electron species
Have one unpaired electron, free radical and very reactive
Dative / Coordinate Bond
:
A coordinate covalent bond is formed when both bonding electron are donated by one atom
Electron donor
Must have at least one lone pair electron to donate the electron pair
Electron accepter
Must have at least one vacant orbital in its outer shell to receive the electron pair
A dative/coordinate bond is a covalent bond in which both shared/bonding electrons originate from one of the joined atoms
The bond may be shown using the symbol '->' to designate the source of shared electrons
Examples: The pair of electrons that bind the N atom to another atom in the NH4+ ion and in the NH3BCl3 molecule originate from N only.
In metal complexes such as [Cu(NH3)4]2+ the bond exist between Cu-NH3 occurred through coordinate bond. Thus the compound is known as coordination compound.
Intermolecular forces
Type of forces
Intramolecular forces (Bonding forces)
:
Attractive forces that hold atoms together in a molecule.
These forces exist within each molecule.
They influence the chemical properties of the substance
Covalent bond
Result of large changes interacting at a very close distance
Intermolecular forces (nonbonding forces)
:
Attractive forces between molecules
These forces exist between molecules
They influence the physical properties of the substance
Van der Waals forces and Hydrogen bond
Result of smaller change interacting at a greater distance
Types of Intermolecular Forces
London Dispersion Forces
:
the weakest intermolecular forces.
temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles.
In nonpolar molecules, there is no dipole-dipole force, yet there is still a force of attraction.
Thus London dispersion forces exists between all substances/compounds.
The instantaneous dipole of the neon atom induces an instantaneous dipole in adjacent atoms, resulting in an attractive force between them.
As example, In a neon atom, though the electrons are spherically distributed over time, at any instant, one side of the nucleus may posses a higher amount of electrons. The atom has a small, instantaneous dipole.
The motion of electrons in one atom affects the motion of electrons in another atom. This causes the instantaneous dipoles of the atoms change together, maintaining a net attractive force.
Dipole-dipole Forces
:
Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule.
Only occurs between polar molecules
When polar molecules come closer to one another, their partial charges act as tiny electric fields that orient and give rise to dipole- dipole attraction. The positive end will attracts the negative end of another
In polar molecules, the molecules tends to align with the opposite polarities for maximum attraction intersection
Hydrogen Bond
:
Hydrogen bond is the strongest intermolecular forces
All polar compounds has dipole-dipole forces but not all polar compounds has hydrogen bond.
The hydrogen bond is a special dipole-dipole forces between the hydrogen atom and either fluorine, oxygen or nitrogen atom (F, O, N)
Make sure that either fluorine, oxygen or nitrogen atom (F. O. N) directly bonded to H within the molecules
Strength of Hydrogen Bond
:
The strength of the hydrogen bond decrease with decrease in the electronegativity value of the atom bonded to the H atom
Since the electronegativity value of the 3 atoms decrease in the order F>O>N, therefore, the strength of the hydrogen bond also decreases in the order:
H-F H-F > H-O H-O >H-N H-N
Metallic Bond
The Electron Sea Model
:
Metallic bonding results from attraction of the metal cations for the mobile, highly delocalized valence electrons.
In metal, bonding electron are delocalized over the entire crystal. Metal atoms in the crystal can be imagined as an array of positive ions immersed in the sea of delocalized valence electron
Physical properties of Metal
high melting points and boiling points
Due to the strong attraction between the metal ions and the delocalised electrons must be overcome to melt or to boil a metal.
Metals are malleable and ductile
:
Due to the ability of the atoms to roll over each other into new positions without each other into new positions without breaking the metallic bond.
Metals conduct heat and electricity
:
Due to the delocalised electrons that can move freely within the metal structure when an electrical current is applied
The Strength of Metallic Bond
:
ionic size
ionic charge
number of valence electron
Generally, the strengt. increases across the period due to the decreasing in atomic size and the increasing number of valence electrons. →resulting in stronger attractions among the metal ions
Trend going down the group: Generally, the strength decreases down the group because the atomic radii