unit 1 f321 chemistry ocr

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unit 1 f321 chemistry ocr
1 shells
1.1 1s shell -2 elec
1.1.1 2nd shell- 8 elec
1.1.1.1 3rd shell-18 elec
1.1.1.1.1 4th shell- 32 elec
1.2 sub-shells
1.2.1 S- 2 elec
1.2.1.1 P- 6 elec
1.2.1.1.1 D- 10 elec
1.2.1.1.1.1 F- 14 elec
1.2.2 ORBITAL - can hold up to 2 electrons, each sub-shell has different orbitals s,p,d,f
1.2.2.1 s orbital has a spherical shape
1.2.2.2 P orbital has a 3D dumbbell shape at right angles to one another
1.2.2.3 orbits are regions around the nucleus that have electron density
1.2.2.4 electron repel each other and have a property called a spin
1.2.2.4.1 in orbitals they have opposite spins to counteract the repulsion
1.3 carbon- 12= international standard
1.3.1 the mass of carbon-12 is exactly 12
1.3.2 the mass of 1/12 of an atom of carbon 12 is exactly 1
1.3.3 RELATIVE ISOTOPIC MASS- the mass of an atom of an isotope compared to 1/12 of the mass of an atom of carbon-12
1.3.3.1 RELATIVE ATOMIC MASS- is the weighted average mass of an atom of an element compared with 1/12 of the mass of am atom of carbon-12
1.3.3.2 RELATIVE MOLECULAR MASS- the weighted average mass of an molecule of a compound compared with 1/12 of the mass of an atom of carbon-12
1.3.3.2.1 ISOTOPES-atoms of the same element with different masses with the same number of protons and neutrons
1.3.3.2.2 RELATIVE FORMULA MASS- the average mass of the formula unit of a compound compared with 1/12 the mass of an atom of carbon-12
2 IONISATION ENERGY
2.1 the first ionisation energy is the energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of a gaseous 1+ ion
2.2 1st ionisation energy == Na(g)-->Na+(g) + e-
2.2.1 2nd i.e==Na+(g)-->Na2+(g) + e-
2.3 factors affecting ionisation energy
2.3.1 Atomic radius-the greater the distance between the nucleus and the outer electron, the less the attractive force
2.3.1.1 Nuclear charge - the greater the number of protons in the nucleus the greater the attractive force
2.3.1.1.1 Electron shielding- the outer electron shells are repelled by any inner shells between electrons and the nucleus
2.4 the second ionisation energy is the energy required to remove 1 electron fro each atom in 1 mole of a gaseous 1+ ion to form a mole of gaseous 2+ ions
2.5 trends in 1st I.E-- There is a increase in 1st I.E across a period because the there is a increase in nuclear charge as electrons are added ti the same shell across a period.
2.5.1 there is a sharp decrease in 1st I.E between the end of one period and the start of the next period, this is because there is a new outer shell which is due to distance and shielding
2.5.1.1 A decrease in 1st I.E down a group due to the presence of extra shells down the group
3 redox
3.1 oxidation- the loss of electrons
3.2 reduction is the gain of electrons
3.3 half equations-
3.3.1 Mg + Cl2 -->MgCl2
3.3.1.1 electron transfer
3.3.1.1.1 Mg-->Mg2+ + e-
3.3.1.1.1.1 Cl2--> 2e- + 2Cl-
3.3.1.1.2 Mg is the reducing agent, it has reduced Cl2
3.3.1.1.2.1 Cl2 is the oxidising agent , it has oxidised Mg to Mg 2+
3.4 oxidation agent accepts electrons from other reactants
3.4.1 non-metals are oxidising agents== F2,Cl2,O2
3.5 A reducing agent donates electrons to another reagent
3.5.1 metals are reducing agents==Na, Fe, Zn
4 Acids bases and Salts
4.1 All acids release H+ ions (aq)
4.2 All bases accept H+ ions (aq)
4.3 An alkali is a soluble base that dissolves in water, releasing OH- ions (aq)
4.3.1 strong alkalis- KOH (potassium hydroxide), LIOH (lithium hydroxide)
4.3.2 weak alkalis- NH3 (ammonia)
4.4 SALT- a compound formed from an acid, when H+ ion from the acid has been replaced by a metal ion or a another positive ion, such as the ammonium ion (NH3+)
4.5 water of crystallisation
4.5.1 hydrated- containing water molecules
4.5.2 anhydrous- without water
4.5.3 the molar proportion of water in hydrated crystals
5 Bonding structure and the periodic table
5.1 ionic
5.1.1 electrostatic attraction between oppositely charged ions
5.1.1.1 ionic lattices
5.1.1.1.1 structure compromising of hundreds of thousands of ions
5.1.1.1.1.1 an ion is surrounded by oppositely charged ions, forming a giant ionic lattice
5.1.2 properties of ionic compounds
5.1.2.1 high melting and boiling point due to lot of energy needed to break the strong electrostatic forces holding the ions rigidly in the solid lattice
5.1.2.2 electrical conductivity- (solid lattice)- ions are in a fixed position and there are no MOBILE charge carriers so ionic compounds are a non conductor of electricity in the solid state. When melted or dissolved in water the solid lattice breaks down and the ions are free to move as mobile charge carriers therefore ionic compounds are electricity conductors in a liquid or aqueous state
5.1.2.3 An ionic lattice often dissolves in polar solvents (e.g) water
5.2 covalent
5.2.1 a shared pair of electrons between atoms of non metals
5.2.2 a molecule is the smallest part of a covalent compound that can take part in a chemical equation
5.2.3 DATIVE COVALENT BOND FORMS WHEN THE SHARED PAIR OF ELECTRONS COMES FROM ONE BONDED ATOM
5.2.3.1 e.g- and ammonium ion NH4+ (containing3 covalent bonds and one dative bond
5.2.4 an electron that takes part in forming a chemical bonds is called a valence electron
5.2.5 simple covalent
5.2.5.1 small molecules with weak intermolecular forces
5.2.5.2 low melting and boiling points due to less energy required to break the weak intermolecular forces
5.2.5.3 no mobile charged particles so the structures are non-conductors of electricity
5.2.5.4 van der waals forces form between a simple molecular structure and a non-polar solvent which weakens the structure simple molecular structures are often soluble in non polar solvents
5.2.6 giant covalent
5.2.6.1 e.g- diamond, graphite and SiO2
5.2.6.2 high melting and boiling point due to lots of energy required to break the strong covalent bonds in the lattice
5.2.6.2.1 except fro graphite there are no mobile charged particles so they are non conductors
5.2.6.2.1.1 the strong covalent bonds in a lattice are too strong to be broken by polar or non-polar solvents
5.3 metallic
5.3.1 the electrostatic attraction between the positive metal ions and delocalised electrons
5.3.2 the delocalised electrons in metals are able to move throughout the structure can cannot assign to s positive metal ion
5.4 shapes of molecules
5.4.1 the shape of the molecule depends on the number of electron pairs surrounding the central atom
5.4.1.1 electron pairs repel each other and move as far apart as possible
5.4.2 molecules with bonded pairs
5.4.2.1 2- linear
5.4.2.1.1 3- trigonal planar
5.4.2.1.1.1 4- tetrahedral
5.4.2.1.1.1.1 6- ocrahedral
5.4.2.1.1.1.1.1 SF6
5.4.2.1.1.1.1.2 90
5.4.2.1.1.1.2 CH4
5.4.2.1.1.1.3 109.5
5.4.2.1.1.2 BF3
5.4.2.1.1.3 120
5.4.2.1.2 BeCl2
5.4.2.1.3 180
5.4.3 molecules with lone pairs
5.4.3.1 0- tetrahedral
5.4.3.1.1 1- pyramidal
5.4.3.1.1.1 2- non-linear
5.4.3.1.1.1.1 H20
5.4.3.1.1.1.2 104.5
5.4.3.1.1.2 NH3
5.4.3.1.1.3 107
5.4.3.1.2 CH4
5.4.3.1.3 109.5
6 electronegativity, polsrity and polarisation
6.1 ELECTRONEGATIVITY- the measure of the attraction of an atom in a molecule for the pair of electrons in a covalent bond
6.1.1 generally the smallest atoms are the most electronegative and the most electronegative atoms are those of highly reactive non-metallic elements- e.g.. O, Fe and Cl
6.2 Polar and non-polar molecules
6.2.1 polar bonds
6.2.1.1 the bonded electrons are shared equally between both atoms, and the bonded atoms have similar electonegativities
6.2.2 non-polar bonds
6.2.2.1 the bonded electrons are shared equally between both atoms and the bonded atoms have similar electronegativities
6.2.3 symmetrical and unsymmetrical molecules
6.2.3.1 each bond is polar but the dipoles act in different directions, the overall effect of the dipoles to cancel out
7
8 Intermolecular forces
8.1 Van der waals forces
8.1.1 weak intermolecular forces that act between molecules
8.1.2 caused by the uneven distribution of electron in molecules which inures a dipole in an neighbouring molecule which is induces more dipoles
8.1.3 the greater the number of electrons in each molecule the larger the oscillating and induced dipoles. the greater the attractive forces between molecules and the greater the van der waals forces
8.1.4 van der waals forces increase in strength with increasing number of electrons
8.2 Permanent dipole-dipole
8.2.1 the delta positive/negative charges on a polar molecule attract oppositely charged dipoles on another polar molecule
8.2.1.1 this gives a weak intermolecular force called a permanent dipole-dipole interaction- eg- HCL
8.2.1.1.1 stronger than van der waals forces although they are both weak
8.3 Hydrogen bonds
8.3.1 strong intermolecular bond between polar molecules
8.3.2 special properties of water due to hydrogen bonds are -that ice is less dense than liquid water this is because particles in solids are usually packed closely together than in liquids however hydrogen bonds hold water molecules apart in an open lattice structure -ice has a high melting point and water has a high billing point because of the strong hydrogen bonds between water molecules
9 Diamond
9.1
10 graphite
10.1

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