Definitions
Binary compound:
a compound composed of two kinds of atoms or two kinds of monatomic ions
Bonding capacity: the number of electrons lost, gained, or shared by an atom when it bonds chemically
Chemical bond: the forces of attraction holding atoms or ions together
Chemical nomenclature: the system, such as the one approved by IUPAC, of names used in chemistry
Coordinate covalent bond: a covalent bond in which both of the shared electrons come from the same atom
Covalent bond: the attractive force, between two atoms of nonmetallic elements, that results when electrons are shared by the atoms
Crystal lattice: a regular, ordered arrangement of atoms, ions, or molecules
Diatomic molecule: composed of two atoms of the same or different elements
Dipole-dipole attraction: an attractive force acting between polar molecules
Electrical conductivity: the ability of a material to allow electricity to flow through it
Electron dot diagram: a representation of an atom or ion, made up of the chemical symbol and dots indicating the number of electrons in the valence energy level; also called Lewis symbol
Hydrate: a compound that decomposes to an ionic compound and water vapour when heated (empirical definition); a compound that contains water as part of its crystal structure (theoretical definition)
Hydrogen bond: a relatively strong dipole-dipole force between a positive hydrogen atom of one molecule and a highly electronegative atom (F,O or N) in another molecule
Intermolecular force: the attractive force between molecules
Intramolecular force: the attractive force between atoms and ions within a compound
Ionic bond: the electrostatic attraction between positive and negative ions in a compound; a type of chemical bond
Ionic compound: a pure substance formed from a mental and a nonmetal
Lewis structure: a representation of covalent bonding based on Lewis symbols; shared electron pairs are shown as lines and lone pairs as dots
Lewis symbol: a representation of an atom or ion, made up of the chemical symbol and dots indicating the number of electrons in the valence energy level; an electron dot diagram
London dispersion force: an attractive force acting between all molecules, including nonpolar molecules
Lone pair: a pair of valence electrons not involved in bonding
Molecular compound: a pure substance formed from two or more nonmetals
Multivalent: the property of having more than one possible valence
Octet rule: a generlization stating that when atoms combine, the convalent bonds between them are formed in such a way that each atom achieves eight valence electrons (two in the case of hydrogen)
Polar covalent bond: a covalent bond formed between atoms with significantly different electronegativities; a bond with some ionic characteristics
Polar molecule: a molecule that is slightly positively charged at one end and slightly negatively charged at the other because of electonegativity differences
Stabel octet: a full shell of eight electrons in the outer energy level of an atom
Tertiary compound: a compound composed of three different elements
Valence: the charge of an ion
Van der Waals force: weak intermolecular attractions, including London dispersion forces and dipole-dipole forces

Forces of Attraction

INTRAmolecular forces of attraction (inside the molecule)
1.) Covalent Bonds
2.) Ionic Bonds
3.) Co-ordinate Covalent
4.) Metallic Bond

H2 H-H
Cl2 Cl- Cl
H20 H-O-H

INTERmolecular forces of attraction (between the molecule)

1.) Hydrogen Bonds (Strongest Force)
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2.) Dipole-Dipole Attraction ( Vanderwaals Force)
3.) London's Dispersion ( Weakest Force)

Valence Electrons- The s and p electrons in the outer energy level. It is the highest occupied eneergy level.

Core electrons- Those in the energy levels below

Electron Configuration for Cations

Metals lose elctrons to attain noble gas configuration
They make positive ions ( cations)
If we look at the electron configuration, it makes sense to lose electrons:
NA :1s2, 2s2, 2p6, 3s1- 1 valence electron
NA+ :1s2, 2s2, 2p6- noble gas configuration

Electron Dots for Cations

Metals will have few valence electrons, these will come off forming positive ions

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Electron configuration for Anions

Non-metals gain electrons to attain noble gas configuration
They make negative ions(anions)
Halide- Ions- ions from chlorine or other halogens that gain electrons

S- 1s2, 2s2, 2p6, 3s2, 3p4- 6 valence electrons
S- 1s2, 2s2, 2p6, 3s2, 3p6- Noble gas configuration

Electron Dots for Anions

Non-metals wil have many valence electrons ( usually five or more)
They will gain electrons to fill outer shell

P- 1s2, 2s2, 2p6, 3s2, 3p3

Stable Electron Configuration

All atoms react to achieve noble gas configuration
Noble gases have 2s and 6p electrons
Eight valence electrons
Also called the octet rule


Ionic Bonds
  • Formed between metals and non-metals
  • Metal ions are called cations
  • Non-metal ions are called anions
  • Force of attration between ions are electrostatic in nature (very strong)
  • These Compounds are:
a) Crystalline solids
b) high melting points
c) conduct electricity in aqueous solution (when dissolved in water)
d) do not conduct electrisity in solid state
e) have definite (fixed) values for MP's
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Covalent Bonds

Formed between two or more non-metals by sharing of electrons
eg: CH4, H2O, NH3, CO2, N2

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Coordinate Covalent bonds:

Formed between an electron rich atom or ion and an electron defiecient atom or ion. Only method of formation is different but once it is formed it is similar to a covalent bond. The bond is indicated by an arrow from the donor to the acceptor.

NH4.JPG
[Lewis Acids] Electron Deficient:
they don't have the right number of electrons to satisfy the orbits (empty orbits)
Acceptors: elements that receive electrons
[Lewis Base] Electron Rich: when a molecule has a lone pair
Donor: when an element gives electrons
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Naming of Compounds:

Chemical Nomenclature: A system such as the one approved by IUPAC

International Union of Pure and Applied Chemistry or IUPAC: An organisation in Paris France which is the regulating body for naming of compounds and providing standards for measurements of other units like mass, length, time etc.

Oxidation number or valency of an ion: The number indicated on a cation or anion showing the number of electrons gained or lost is also known as valency.

Oxynations: A polyatomic ion containing oxygen.

Multivalent ions: A cation (netal ion) capable of showing more than one valency e.g. Pb2+ and Pb4+

Classical Approach for multivalent ions:
-ous lower oxidation state Fe2+ Ferrous
-ic higher oxidation state Fe3+ Ferric

Hydrates: Tertiary ionic compounds whihc contains water molecules within the crystal structure are called hydrates e.g. CuSO4.5H20

Acids: Common definition would be compounds that are capable of furnishing H+ ions in aqueous solutions.

Bases: Classical definition would be compounds capable of furnishing OH- ions in aqueous solution.

Binary Componds:
Metals first followed by non-metals
For multivalent ions indicate oxidation number or ionic charge in brackets using roman numerals

Tertiary Compounds:
Polyatomic ions are a group of covalently bonded atoms which has a net charge may be positive or negative
Metals first followed by poly atomic ions that are negatvely charged
Poly atomic cat ion first followed by poly atomic negative ion

Polyatomic Oxyanions of Halogens:
General Formula
Name
Cl
Br
I
XO-
Hypohalite
Hypochlorite
Hypobromite
Hypoiodite
XO-²
Halite
Chlorite
Bromite
Iodite
XO-³
Halate
Chlorate
Bromate
Iodate
XO-⁴
Perhalate
Perchlorate
Perbromate
Periodate

Ending
Number of Oxygen Atoms
Formula of Ion
Hyposulphite
2
SO²- 2
Sulphite
3
SO²-3
Sulphate
4
SO²- 4
Persulphate
5
SO²- 5
Hypo is low-less oxygen
Hyper is high-more oxygen

Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
2
3
4
5
6
7
8
9
10


Hydrates:
Name of the compound followed by the number of molecules of water of crystallization using the above terminology

Compound
IUPAC Name
CuSO⁴. 5H²O
Copper(II) Sulphate Penta Hydrate
Na²SO⁴.10H²O
Sodium Sulphate Deca Hydrate
MgSO⁴ .7H²O
Magnesium Sulphate Hepta Hydrate
CaSO⁴ .2H²O
Calcium Sulphate Di Hydrate
Al²O³ .2H²O
Aluminum Oxide Di Hydrate


Acids:
Binary acids: Ending Aqueoushydrogen----------
Formula
Classical or Trivial Name
IUPAC
HF(aq)
Hydrofluoric acid
aqueous hydrogen fluorid
HCl(aq)
Hydrochloric acid
aqueous hydrogen chloride
HBr(aq)
Hydrobromic acid
aqueous hydrogen bromide
HI(aq)
Hydroiodic acid
aqueous hydrogen iodide
H²S(aq)
Hydrosulphuric acid
aqueous hydrogen sulphide
Oxy Acids of Chlorine

Formula
IUPAC
Classical Name
HClO(aq)
Aqueous Hydrogen hypochlorite
Hypochlorous acid
HClO²(aq)
Aqueous Hydrogen chlorite
Chlorous acid
HClO³(aq)
Aqueous Hydrogen chlorate
Chloric acid
HClO⁴(aq)
Aqueous Hydrogen perchlorate
Perchloric acid

Aqueous hydrogen-----
e.g. HCl Aqueous hydrogen chloride
e.g. HOCl Aqueous hydrogen hypochlorite

Bases:
Hydroxides
Aqueous metal hydroxide....
e.g. NaOH Aqueous sodium hydroxide

VSEPR (Valence Shell Electron Pair Repulsion)Theory:
Lone pair-lone pair repulsion is greater than
Lone pair-bond pair repulsion is greater than
Bond pair-bond pair repulsion

Shapes Of Molecules



Total Number of electron pairs
Arrangement of electron pairs
Number of bonding pairs of electrons
Number of lone pairs of electrons
Shape of Molecule
Name of Shape
Bond Angle
Examples
not applicable
linear
1
not applicable
external image linea2at.gif
linear
180o
H2, HCl
2
linear
2
0
external image linea3at.gif
linear
180o
CO2, HCN
3
trigonal planar
3
0
external image trigplan.gif
triganol planar
120o
BCl3, AlCl3
4
tetrahedral
4
0
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tetrahedral
109.5o
CH4, SiF4
3
1
external image trigpyra.gif
trigonal pyramidal
<109.5o (bond angles in ammonia, NH3, are 107o)
NH3, PCl3
2
2
external image bent.gif
bent
<109.5o (bond angles in water, H2O, are 105o)
H2O, SCl2
5
trigonal bipyramidal
5
0
external image trigbipy.gif
trigonal bipyramidal
120o in the trigonal planar part of the molecule, 90o for the others
PCl5
6
octahedral
6
0
external image octahedr.gif
octahedral
90o
SF6