Chapter+Two+Chemical+Bonding

//__ 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) 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



__**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

a) Crystalline solids b) high melting points c) conduct electricity in aqueous solution (when dissolved in water) d) do not conduct **electrisity in so**lid state e) have definite (fixed) values for MP's
 * __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:

__**Covalent Bonds**__

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


 * __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.

[Lewis Acids] __Electron Deficient:__** they don't have the right number of electrons to satisfy the orbits (empty orbits)
 * [[image:NH4.JPG]]
 * __Acceptors:__** elements that receive electrons
 * __[__Lewis Base] __Electron Rich:__** when a molecule has a lone pair
 * __Donor:__** when an element gives electrons


 * __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.

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

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
 * Tertiary Compounds:**

__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 ||

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


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

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


 * 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 ||

//Binary acids: Ending Aqueoushydrogen--// __Oxy Acids of Chlorine__
 * Acids:**
 * 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 ||


 * 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

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

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

**// __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 || [[image:http://www.ausetute.com.au/images/linea2at.gif width="50" height="50"]] || linear || 180o || H2, HCl ||
 * 2 || linear || 2 || 0 || [[image:http://www.ausetute.com.au/images/linea3at.gif width="100" height="100"]] || linear || 180o || CO2, HCN ||
 * 3 || trigonal planar || 3 || 0 || [[image:http://www.ausetute.com.au/images/trigplan.gif width="100" height="100"]] || triganol planar || 120o || BCl3, AlCl3 ||
 * 4 || tetrahedral || 4 || 0 || [[image:http://www.ausetute.com.au/images/tetrahed.gif width="100" height="100"]] || tetrahedral || 109.5o || CH4, SiF4 ||
 * ^  || 3 || 1 || [[image:http://www.ausetute.com.au/images/trigpyra.gif width="100" height="100"]] || trigonal pyramidal || <109.5o (bond angles in ammonia, NH3, are 107o) || NH3, PCl3 ||^   ||
 * ^  || 2 || 2 || [[image:http://www.ausetute.com.au/images/bent.gif width="50" height="50"]] || bent || <109.5o (bond angles in water, H2O, are 105o) || H2O, SCl2 ||^   ||
 * 5 || trigonal bipyramidal || 5 || 0 || [[image:http://www.ausetute.com.au/images/trigbipy.gif width="100" height="100"]] || trigonal bipyramidal || 120o in the trigonal planar part of the molecule, 90o for the others || PCl5 ||
 * 6 || octahedral || 6 || 0 || [[image:http://www.ausetute.com.au/images/octahedr.gif width="100" height="100"]] || octahedral || 90o || SF6 ||