Bonding

Octet Rule Octet Rule states that when bonding occurs atoms tend to reach an arrangement with 8 electrons in the outer shell Exceptions: - d-block elements, including the transition metals (variable valency) - small elements like He, Li (want full outer shell)

Valency The valency of an element is the number of bonds an atom of the element forms when it reacts Valency is equal to the number of electrons an atom needs to lose or gain when forming a bond

Group number: Valency: I 1 II 2 III 3 IV 4 V VI VII VIII

Transition metals have a variable valency They can lose different amounts of electrons depending on what they bond with

Ionic Bonding An ion is a charged atom or group of atoms When an atom loses an electron it forms a positive ion (more protons than electrons) When an atom gains an electron it forms a negative ion (more electrons than protons) An ionic bond is the force of attraction between oppositely charged ions

Examples: Sodium chloride (NaCl) Na + Cl → Na+ Cl- 2,8,1 2,8,7 2,8 2,8,8 Na + Cl → Na + Cl -

Magnesium fluoride (MgF2) Mg + F + F → Mg2+ F- F- 2,8,2 2,7 2,7 2,8 2,8 2,8 Mg + F + F → Mg 2+ F - F -

Group ions: Name: Formula: Hydroxide ion Nitrate ion Hydrogencarbonate ion Permanganate ion OH- NO3- HCO3- MnO4- Carbonate ion Chromate ion Dichromate ion Sulfate ion Sulfite ion Thiosulfate ion CO32- CrO42- Cr2O72- SO42- SO32- S2O32- Phosphate ion PO43- Ammonium ion NH4+

Ionic structure Ions are arranged in compounds in a crystal lattice structure Crystal lattices are made up of unit cells linked together (repeating unit) E.g., NaCl molecule = unit cell NaCl compound = crystal lattice

Ionic compound formulas Chemical formula represents the compound using chemical symbols and numbers Have to be able to work out ionic formulas for 1st 36 elements Ionic bonds usually form between Group I & II and Groups VI & VII

Group I & II (metals) lose electrons and form positive ions, e.g., Li → Li+ , Mg → Mg2+ Group VI & VII (non-metals) gain electrons and form negative ions, e.g., O → O2-, Cl → Cl-

These oppositely charged ions are attracted to each other forming an ionic bond, e.g., Li+ + Cl- → LiCl Li ion Cl ion Neutral LiCl compound Mg2+ + O2- → MgO Mg ion O ion Neutral MgO Ionic componds are neutral because the charges of the ions that form them should balance

Covalent bonding A covalent bond consists of one or more shared pairs of electrons A molecule is a group of atoms joined together. It is the smallest part of an element that can exist independently. Need to know 5 examples:

(a) Hydrogen molecule - H has 1 electron in its outer shell so needs to gain another electron to have a full outer shell - The two H atoms share their electrons to make them stable H H

(b) Chlorine molecule - Cl has 7 electrons in its outer shell so needs to gain another electron to have a full outer shell - The two Cl atoms share 1 of their electrons with the other to make them stable Cl Cl

(c) Water molecule - H has 1 electron its outer shell so needs to gain another electron to have a full outer shell - O has 6 electrons in its outer shell so needs to gain another 2 to have a full outer shell - Each H atom shares its electron with the O atom to make them stable O H H

(d) Ammonia molecule - H has 1 electron its outer shell so needs to gain another electron to have a full outer shell - N has 5 electrons in its outer shell so needs to gain another 3 to have a full outer shell - Each H atom shares its electron with the N atom to make them stable N H H H

(e) Methane molecule - H has 1 electron its outer shell so needs to gain another electron to have a full outer shell - C has 4 electrons in its outer shell so needs to gain another 4 to have a full outer shell - Each H atom shares its electron with the C atom to make them stable H H C H

Types of covalent bonds Single bonds share one pair of electrons e.g., H Cl Double bonds share two pairs of electrons e.g., O O Triple bonds share three pairs of electrons e.g., N N

Sigma bond is the head on overlap of two atomic orbitals Pi bond is the side on overlap of two atomic orbitals Single bond = sigma bond e.g., H Cl

Double bond = 1 sigma & 1 pi bond e.g., O O Triple bond = 1 sigma & 2 pi bonds N N

Electronegativity In a covalent bond between two atoms of the same element the pair of electrons in the bond is shared equally But in a covalent bond between atoms from different elements one of the atoms has a greater attraction for the electrons than the other This can cause a slight charge to form on the atoms in the bond

Example - HCl The Cl atom is more attracted to the electrons than the H atom, δ+H Cl δ- Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond

Very reactive elements, like K and Na, have low electronegativity values – can be called electropositive Very unreactive elements, like F and Cl, have high electronegativity values

Increase across a period because: Increase in nuclear charge – Increase in number of protons but electrons all added to same shell (b) Decrease in atomic radius – Electrons are closer to the nucleus (c) No screening effect – Electrons are added to the same shell so there are no inner electrons to shield nuclear charge

Decreases down a group because: Increase in atomic radius – Electrons are further away from the nucleus (b) Screening effect – Inner electron shells shield the outer ones from the nuclear charge (c) Decrease in nuclear charge –