Chemistry revision F321

James Farley
Flashcards by James Farley, updated more than 1 year ago
James Farley
Created by James Farley about 6 years ago
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OCR AS level Chemistry F321

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Question Answer
Atomic number The number of protons in the nucleus of an atom.
Mass number The number of protons and neutrons in the nucleus of an atom.
Isotopes Atoms of the same element with different numbers of neutrons.
Why do isotopes have similar chemical properties? They have the same number of electrons in their outer shell.
Relative isotopic mass The mass of an isotope compared to 1/12th of the mass of one atom of carbon 12.
Relative atomic mass The weighted mean mass of an atom compared to 1/12th of one atom of carbon 12.
Protons Relative mass: 1 Relative charge: +1 Position in the atom: Nucleus
Neutrons Relative mass: 1 Relative charge: 0 Position in the atom: Nucleus
Electrons Relative mass: 1/2000 Relative charge: -1 Position in the atom: Shells
Relative Molecular mass The weighted mean mass of a molecule relative to 1/12 of one atom of carbon 12.
Orbital A volume of space that can hold up to two electrons with opposite spins.
Order of the orbitals 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 (max 36 electrons)
First ionisation energy The amount of energy required to remove one electron from each atom in one mole of gaseous atoms (to form one mole of positive gaseous ions.)
Successive ionisation energies Second ionisation energy: O+ (g) ----> 0 2+ (g) + e- Third ionisation energy: O 2+ (g) ----> O 3+ (g) + e-
Trend in successive ionisation energies Removal of additional electrons results in an increase in ionisation energy because: Same number of protons are now attracting fewer electrons in the ion One less electron means there is less electron-electron repulsion The O+ ion is smaller than the O atom, therefore the next electron to be removed will be closer to the nucleus Next electron will experience greater nuclear attraction.
CARS: across a period Increase in ionisation energy: C inc - less electrons added to same no. of protons A inc R dec - extra electron shell added S eq - electrons added to same shell
CARS - down a group Ionisation energy decreases: C inc - outweighed by radius and shielding A dec R inc - extra electron shell added S - more electron-electron repulsion
Acid + Metal ---> 2HCl (aq) + Mg (s) ---> Salt + hydrogen gas MgCl2 (aq) + H2 (g)
Acid + Metal Oxide ---> 2HCl (aq) + CaO (s) ---> Salt + Water CaCl2 (aq) + H20 (l)
Acid + Metal hydroxide ---> HNO3 (aq) + NaOH (s) ---> Salt + Water NaNO3 (aq) + H20 (l)
Acid + Metal carbonate ---> H2SO4 (aq) + CaCO3 (s) ---> Salt + Water + Carbon dioxide CaSO4 (aq) + H2O (l) + CO2 (g)
Ionic bond The electrostatic force of attraction between oppositely charged ions (occurs between a metal and a non metal.)
Covalent bond A shared pair of electrons (occurs between two non metals.)
Dative covalent bond A shared pair of electrons, both of which come from the same atom. E.g. NH4+ formed by reaction of NH3 + H+.
Metallic bond The electrostatic force of attraction between positive metal ions and delocalised electrons.
Exceptions of the octet rule When there are not enough electrons to complete the octet. E.g. Be & B in period 2 when they are the central atom.
Expansion of the octet rule When each atom has more than 8 electrons. E.g. Groups 5-7 from period 3 downwards when they are the central atom.
Linear shaped molecules E.g. Becl2 BeH2 CO2 2 bonded pairs 0 lone pairs 180° bond angles
Trigonal planar shaped molecules E.g. BF3 BCl3 BBr3 SO3 3 bonded pairs 0 lone pairs 120° bond angles
Tetrahedral shaped molecules E.g. CH4 SiH4 NH4+ CCl4 4 bonded pairs 0 lone pairs 109.5° bond angles
Octahedral shaped molecules E.g. SF6 SeF6 6 bonded pairs 0 lone pairs 90° bond angles
Pyramidal shaped molecules E.g. NH3 NCl3 PCl3 3 bonded pairs 1 lone pair 107° bond angles (-2 . 5° for each lone pair)
Non-linear shaped molecules E.g. H2O H2S 2 bonded pairs 2 lone pairs 104.5° bond angles
Electronegativity The ability of an atom to attract a pair of electrons in a covalent bond towards itself.
Pauling scale Electronegativity increases to the top right of the periodic table, with F being the most electronegative.
Permanent dipole A difference in electronegativity between 2 atoms results in a small charge difference. Eg. HCl H △+ Cl△- electrons always closer to Cl atom.
Polar bonds in a non polar molecule A molecule may have polar covalent bonds because the molecules are symmetrical. The dipoles act in opposite directions and cancel each other out.
Permanent dipole-dipole forces One molecule can attract the opposite permanent dipole in a neighbouring molecule.
Van der waals' forces Forces arise due to movement of electrons (uneven distribution.) This causes an instantaneous dipole (temporarily.) This temporary dipole induces a dipole in neighbouring molecule. Dipoles attract one another (weak VDW's)
Increasing van der Waals' forces VDW's forces increase with an increase in the number of electrons. More electrons result in larger temporary and induced diploles, which results in a greater force of attraction between molecules. Therefore boiling point increases.
Hydrogen bonding Special permanent dipole-dipole force between O-H or N-H bonds. The H△+ attracts the lone pair of electrons in the O△-/N△- in neighbouring molecule.
Special properties of water Ice is less dense than water: the hydrogen bonds hold H2O molecules further apart in an open lattice. H2O has a higher than expected boiling point: hydrogen bonds are the strongest intermolecular force therefore require more energy to break.
Simple molecular lattice E.g. H2 I2 H2O Low melting/boiling points: weak VDW's forces require little energy to be broken. Don't conduct electricity: no free charged particles.
Giant covalent lattice E.g. Si, SiO2, diamond&graphite High melting/boiling point: strong covalent bonds require lots of energy to break. Don't conduct electricity: no free moving charged particles
Diamond Giant covalent lattice. Doesn't conduct electricity: no free moving charged particles. Hard: strong covalent bonds throughout tetrahedral structure.
Graphite Giant covalent lattice with layers of carbon atoms. Conducts electricity: delocalised electrons between carbon layers. Soft: weak VDW's forces between the layers allowing them to slide.
Ionic compounds High melting/boiling point: strong electrostatic force of attraction between oppositely charged needs to be broken. Don't conduct in solid state: fixed position. Conduct when liquid/aqueous: ions free to move. Soluble in polar solvents, surround and break down lattice.
Metal bonding The strong electrostatic force of attraction between positive metal ions and delocalised electrons. High melting/boiling point: strong electrostatic force of attraction. Conduct electricity: delocalised electrons are free to move. Ductile: drawn out (wires.) Malleable: hammered into shape.
ionic charge affecting melting point of a metal E.g. Na+ Mg2+ Al3+ As ionic charge increases the number of delocalised electrons increases. Therefore attraction increases between delocalised electrons and metal ions (stronger metallic bonding.)
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