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73 Cards in this Set
- Front
- Back
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Isotopes
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Atoms of an element with different numbers of neutrons and different masses
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Relative atomic mass
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The weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of the C-12 isotope which has a mass of exactly 12
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Relative isotopic mass
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The weighted mean mass of an atom of an isotope compared with 1/12 the mass of the C-12 isotope which has a mass of exactly 12
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The relative molecular mass
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The weighted mean mass of the molecule relative to 1/12 the mass of an atom of the C-12 isotope which has a mass of exactly 12
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A mass spectrometer... Does what?
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Provides a trace which shows the mass of each isotope and its relative abundance
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How to calculate relative atomic mass...
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(% of relative mass 1) + (% of relative mass 2)
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A mole
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Is the amount of any substance containing as many particles as there are carbon atoms in exactly 12g of the C-12 isotope
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Avogadro constant...
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6.02 x 10^23
(Is the number of particles per mole) |
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Molar mass...
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Add together the relative atomic masses for each atom that make up 1 mole
n=m/ M M= molar mass |
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Empirical formula
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The simplest whole number ratio of atoms of each element present in a compound
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Molecular formula
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The actual number of atoms of each element in a compound
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Molecular formula 2
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= relative molecular mass/ empirical formula mass
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1 mole of gas at RTP and pressure occupies...
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24dm^3 = 24000cm^3
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Moles =
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=Mass/ molar mass
=Volume in cm/ 24000cm^3 = cv/ 1000 |
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The concentration of a solution
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Is the amount of solute of mol dissolved per 1dm^3
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A lower PH...
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= higher concentration of hydrogen + ions
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A diprotic acid
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Has two replaceable H+ ions
And can form normal and acidic salts Eg: sulphuric, H2SO4 |
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An acid
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Releases H+ ions in aqueous solution.
Is a Proton donor |
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An alkali
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Is a soluble base that releases OH- ions in aqueous solution
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A base
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Metal oxides, metal hydroxides, ammonia
A proton acceptor They neutralise acids and carbonates |
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A strong acid
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One that completely dissociated in solution
HCl NaOH H2SO4 |
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A weak acid
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Is only partially ionised in solution
CH3COOH--> ethanoic acid Has a reversible sign CH3COOH<-> CH3COO- & H+ H+ = hydroxonium ion |
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Acid reactions
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Acid + metal= salt + hydrogen
Acid+carbonate=salt+CO2+ H2O Acid+ base= salt + H2O Acid+alkali= salt + H2O |
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Hydrated...
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A crystalline compound containing water molecules
[anhydrous=crystalline compound containing no water molecules] |
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Water of crystallisation
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Water molecules that form an essential part of the crystalline structure of a compound
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A salt
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Is produced when the H+ ion of an acid is replaced by a metal ion or NH4+
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Methyl orange
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Acid= red
Base= yellow |
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Phenolphthalein
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Acid= colourless
Base = pink |
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Oxidation
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Loss of elections
An increase in oxidation number= reducing agent |
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Reduction
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Gain of electrons
Decrease in oxidation number= oxidising agent |
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An unreacted element
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Has an oxidation number of zero
The sum of oxidation numbers in a compound is zero The most electronegative element in a compound has a negative oxidation number |
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Oxidation numbers
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Fluorine= -1
Oxygen= -2 Hydrogen= +1 [ in peroxides, oxygen= -1] Oxidation numbers are constant in groups 1&2 |
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Redox reactions
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Metals form ions by losing elections with an increase of oxidation Numbers to form positive ions
Nonmetals generally react by gaining electrons with a decrease in oxidation number to form negative ions |
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Ionisation energies(g)
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Decrease down the group as the atomic radius increases, the shielding increases, The nuclear charge increases but is outweighed. It takes less energy to overcome the nuclear attraction
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Ionisation energies(p)
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The ionisation energy increases across a period as shielding is the same, radius is smaller, as increase in nuclear charge. Harder to overcome nuclear attraction
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In electron shielding
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See electrons are shielded from the electrostatic attraction of the positive nucleus by the inner shells of electrons
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Anomally at aluminium
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Electron entering 3p orbital which has a higher energy level and is shielded by spherical 3S orbital
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The first ionisation energy
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Is the energy needed to remove one mole of electrons from one mole of atoms in the gaseous state
Na(g) --> Na+ (g) & e- |
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The second ionisation energy
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Is the energy needed to remove one mole of electrons from each ion in a mole of gaseous 1+ ions to form one mole of gaseous 2+ ions
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The second ionisation energy level
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Is greater as shielding decreases, distance decreases, charge increases. Harder to overcome nuclear attraction
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An orbital
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A region that can hold up to 2 electrons with opposite spins
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Number of orbitals in subshells
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S=1
P=3 D=5 |
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S,p,d sub shells
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Each orbital filled with one electron first before pairing starts to reduce repelling and so less energy is needed to separate them
S= group 1&2 P= group 345678 D= transition metals F= actinides & lanthanides |
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Energy order of sub shells
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1s
2s 2p 3s 3p 4s 3d 4p 4d |
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Chromium and copper
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Cr-1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu-1s2 2s2 2p6 3s2 3p6 4s1 3d10 |
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Ionic bonding
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Strong Electrostatic attraction between oppositely charged ions
Held in a giant crystalline lattice Metal and non metal Can only conduct when molten Greater charge= greater force of attraction |
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Covalent bond
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Shared pair of electrons
Strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond Between non metals Can be simple or giant |
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Dative covalent bond
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A shared pair of electrons in which the pair are provided by only one of the bonding atoms
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Shape and angle of molecules
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Shape determined by repulsion between electron pairs surrounding a central atom
Lone pairs of electrons repel more than bonding pairs |
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Shapes and angles molecules 2
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Electronegativity
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The ability of an atom to attract the bonding electrons in a covalent bond
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Permanent dipole-dipole force
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Is a weak attractive force between permanent dipoles in neighbouring polar molecules
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Permanent dipoles
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Occur when covalently bonded atoms have different Electronegativities resulting in a polar bond
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Type of bond and Electronegativity difference
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Van Der Waal's forces
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Weak attractive forces between induced dipoles in neighbouring molecules
- the more electrons= larger charge cloud= more induced dipoles= more VdW forces= higher BP - movement of electrons causes instantaneous dipoles. Creates induced dipoles in neighbouring molecules. Electrostatic attraction between neighbouring induced dipoles= VdW forces |
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Metallic bonding
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The attraction of positive ions to delocalised electrons
In a metal= giant structure of positive ions surrounded by a sea of delocalised electrons Higher charge= more delocalised electrons= greater electrostatic attraction Are malleable(can roll over each other into new positions) Can conduct |
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Ionic structures
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Giant 3d crystalline lattice structure held together by the attraction between oppositely charged ions. Can be split along certain angles, have high MP, often soluble, conduct when molten or dissolved
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Highest BP
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Group 3= giant covalent
Group 2= metallic- higher charge of ions |
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Min BP
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Group 8- simple molecular, weak VDW forces
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Hydrogen bonding
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A strong 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
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In jce
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The extensive hydrogen bonding produces a very open tetrahedral lattice structure in which the water molecules are held further apart than in water
This makes the density of ice less than that of water so it floats. Hydrogen bonding account for anomalously high BP- more energy needed to overcome additional intermolecular forces |
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Graphite
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Diamond
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Giant covalent=Covalent bonding extended in 3 dimensions in a repeating network of bonds
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Draw hydrogen bonding
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Elements in a group
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Have similar outershell electron configurations, resulting in similar properties
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Periodicity
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Regular or a periodic variation in the properties of an element with atomic number and position in the periodic table
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PH of group 2 elements with water
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10
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Thermal decomposition of the carbonates of elements in group 2
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Higher temperature needed down the group
Charge stays the same, size increases with more shells, the smaller the eye on the higher the charge density. The smaller Ion distorts The anion charge cloud and weakens the covalent bond of the carbonate ion. The carbonate breaks more easily |
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Use of Ca(OH)2 and Mg(OH)2
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Neutralise acid soils and indigestion tablets as an antacid
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Trend in reactivity of group 7 elements (VdW, colour, state)
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Disproportionation
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A reaction in which an element is simultaneously oxidised and reduced
Occurs in the reaction of chlorine with water( water purification) : Cl2 + H2O = HClO + HCl Chloric (1) acid is unstable and breaks down easily on standing 2HClO= 2HCl + O2 (this is the bleaching action of chlorine water). Kills bacteria but hazards of toxic chlorine chemicals. Chlorinated hydrocarbons are carcinogenic The reaction of chlorine with cold delude aqueous sodium hydroxide (to form bleach): Cl2 + NaOH= NaClO +NaCl+H2O |
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Halide tests
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Adding ammonia helps distinguish between precipitates. Dilute ammonia for AgCl. Conc ammonia for AgBr. AgI does not dissolve with any ammonia
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Displacement reactions with halides
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