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» AS-Level Chemistry Quick Revision Notes: Atoms, Bonds and Groups
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Posted on: January 16th, 2013

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Isotopes – Atoms of the same element – same number of protons and electrons – same chemical reactions (speed varies) Physical properties such as melting point and density varies

(Relative atomic mass x Abundance) + (repeat) / 100

+ Protons relative mass 1

Neutrons relative mass 1

– Electrons 1/2000th the mass of a proton

Relative Atomic Mass is the weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of Carbon 12

Relative Isotopic Mass is the mass of an atom of an isotope compared with1/12 of the mass of an atom of Carbon 12

Equations

Mole the amount of any substance containing as many particles as there are carbon atoms in exactly 12g of the carbon-12 isotope

Number of moles = Mass / Atomic Molecular Mass

Empirical Formula simplest whole number ratio of atoms of each element present in a compound

Mass / Atomic Mass = Moles / smallest number = Ratio

Molar mass Number of atoms of each element in a molecule

Molecular Mass / Empirical = Number of molecules

Molecular Formula Number of atoms of each element in a molecule

Avogadro’s Law for any gas the number of particles is always the same

Volume of gas (cm3) / 24 (24000) = Number of moles

Concentration x Volume [cm3/1000] = Number of moles

1 Mole is equivalent to 6.023 x 10^23 (Avogadro’s number) units of any substance.

Acids

Acid + Metal > Salt + Hydrogen

Acid + Metal Oxide > Salt + Water

Acid + Carbonate > Salt + Water + Carbon dioxide

Neutralisation H+(aq) + OH-(aq) > H2O (l)

Acid + Alkali > Salt + Water

Salt – The H+ ion in an acid has been replaced by a metal ion

Acids – HCl, H2SO4, HNO3 Alkalis (soluble base) – metal oxides, metal hydroxides, ammonia Finding x in hydrated solutions

Mass of solution – mass after heated – Use empirical formula method

Redox / Disproportionation

Oxidation The number of electrons an element uses to bond to other atoms

Uncombined elements  – Ca, He, O2, Cl2 – Oxidation number of 0

Combined oxygen – H2O, CO2 – Oxidation number of -2

Combined hydrogen – HCl – Oxidation number = +1

Ions Li+ = +1, Ca2+ = +2, Cl- = -1 Oxidation number = depends on charge of ion

Metals generally lose electrons and have an increase in oxidation number to form positive ions – oxidised Non-metals generally gain electrons and have a decrease in oxidation number to form negative ions – reduced

Oxidation Involves Loss Reduction Involves Gain of electrons

Oxidation is an increase in oxidation number

Reduction is a decrease in oxidation number

Electron Structure

An atomic orbital is a space within an atom that can hold up to two electrons with an opposite spin S-Orbital 2e-  P-Orbital 6e-  D-Orbital 10 e-  F-Orbital 14 e-

Degenerate Orbitals (P, D, F)

The first ionisation energy of an element is the energy required to remove one mole of electrons from a mole of gaseous atoms to form a mole of 1+ ions

X(g) > X(g) + e-

Depends on –

Atomic Radius (smaller nuclear attraction)

Electron Shielding (inner electrons repel outer)

Nuclear Charge (attraction increases)

Lowest available energy needs to be filled up first (Chromium & Copper are exceptions) – In the case of degenerate orbitals, the electrons are filled in singly with a parallel spin, then pairing starts – Hund’s rule – Half full or completely full configuration is more stable
1st shell < 2nd shell < 3rd shell < 4th shell

s < p < d < f

1s2 < 2s2 < 2p6 < 3s2 < 3p6 < 4s2 < 3d10 < 4s6 <5s2 <4d10 < 5p2

Bonding and Structure

Ionic bonding A metal with a non-metal, held together by electrostatic forces Covalent bonding Two non-metals, bonded together by a shared pair of electrons Co-ordinate (dative) bond A covalent bond, (shared pair of electrons), both electrons come from the same atom

Ionic bonding

NaCl MgO Al2O3 Metal+Non-metal

Positive and negative ions have strong electrostatic forces of attraction

High melting and boiling point

Conducts electricity when melted or dissolved in water as particles are now free to move

Crystalline solids at room temperatureMost are soluble in water as ions can separate

Hard but brittle, not malleable

Bigger the charges the stronger the bonding

 

Covalent bonding

H2 CH4 CO2 Non-metal+ non- metal

Covalent bonds are formed by atoms sharing electrons to form molecules – strong electrical forces of attraction, weak intermolecular forces

Low melting and boiling point

Does not conduct electricity as there are no ions, does in water (electrolysis)

Soluble in non-polar solutions

Gas at room temperature

 

Giant covalent lattice

Diamond, silicon, graphite

Each carbon shares electrons with four other carbon atoms to form a 3D tetrahedral lattice – strong covalent bonds Graphite – 3D hexagonal layered structure

High melting and boiling point

Does not conduct electricity as elect rons are not free to move Does conduct electricity because delocalised electrons carry the charge of an electric current

Insoluble in all solvents

Solid at room temperature

Lattice can absorb external forces giving hardness Graphite – weak interlayer forces make overall structure soft and brittle

Metallic bonding

Metal atoms

Electrostatic attractive forces between the delocalized electrons gathered in an “electron sea”, and the positively charged metal ions.

High melting and boiling point

Most are soluble in water as ions can separate

Insoluble

Solid at room temperature

Malleability and ductility result from layers of atoms sliding over each other

Silvery surface that may be easily tarnished. Alloys are a mixture not compound.

Electronegativity

Electro-negativity is the ability of an atom to attract electrons
Linear 180°

Trigonal 120°

Tetrahedral 109.5°

Trigonal Pyramid 107° 1 lone pair

Non-linear 104.5° 2 lone pairs

Trigonal by-pyramid 120°

Octahedral 90°
Lone pairs have a bigger repulsion, which makes the bond angles smaller

Lone + Lone > Lone + Bond > Bond + Bond

Polar covalent

2.55 C — F 3.98

+ –|—-> –

e.g. – NH3  HCl

Good conductors in molten state – ions a free

High melting and boiling points

Soluble in Water

 

Non-polar

H2, Cl2 H2O CCl4

(Dipole moments cancel each other out)

Poor conductors – no mobile electrons or ions

Low melting and boiling points – weak molecular forces

Insoluble in Water

The smaller the atom, the larger the electronegativity – less shielding and small radius

Ionic bonding ——- Polar covalent ——- Non-polar covalent

Full charges ——- Partial charges ——- Electronically symmetrical
Polar covalent – Intermolecular forces of attraction need to be overcome

Non-polar covalent – Weak intermolecular forcers and temporary dipoles

Wan der Waals forces are attractive forces between induced dipoles in neighbouring molecules.

 

At any moment oscillations produce and instantaneous dipole. This induces further dipoles on neighbouring molecules. The small induced dipoles attract one another causing weak intermolecular forces.

Van der Waals are weak forces but as you go down the group the element will have a higher boiling point because the van der Waals forces get stronger due to size of atom. This explains why boiling points of noble gases increase down the group
Hydrogen Bonding Permanent dipole-dipole attraction between an electron deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom on a different molecule.
Ice is less dense than water – because the H bonds form a 3D crystalline structure, which is less closely packed.

Water is the only hydride with H-bonding – high melting and boiling point as strength of intermolecular forces increases.

Atomic Radius Half the distance between two atomic nuclei – in a covalent or metallic lattice

Across a period Increase in nuclear charge, reduction in radius, higher attractive force

Down a group Addition of electron shells, more shielding effect,  increase in radius, lesser attractive forces

Group 2

Alkali Earth Metals

Beryllium (Be)

Magnesium (Mg)

Calcium (Ca)

Strontium (Sr)

Barium (Ba)

Radium (Ra)

2M + O2 > 2MO
M + H2O > MO + H2
M + 2H2O > M(OH)2 + H2

Low densities
High melting / boiling points  Colourless states
Oxidises
Going down the group

More reactive

More alkaline

More soluble

More shells

Atomic radius increase

1st ionisation energy decreases
Group 7 Halogens

Fluorine (F)

Chlorine (Cl)

Bromine (Br)

Iodine (I)

Astatine (At)
Cl2 + 2Br- > 2Cl- + Br2      orange in water and cyclohexane

Cl2 + 2I- > 2Cl- + I2      brown in water and purple in cyclohexane

Br2 + 2I- > 2Br- + I2      brown in water and purple in cyclohexane

Ag+ + Cl- > AgCl    white precipitate   dissolves in dilute ammonia

Ag+ + Br- > AgBr    cream precipitate  dissolves in concentrated ammonia

Ag+ + I- > AgI           yellow precipitate    does not dissolve

Going down the group Less reactive

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