Chemistry (Mrs Foord)

The changing atom

Dalton's atomic theory (early 1800s): atoms are tiny particles making up elements, cannot be divided, atoms of a given element are the same. he used symbols for elements & developed the table of atomic masses.

Joseph John (J.J.) Thomson discovers electrons (1897-1906): realised cathode rays were streams of particles with a negative charge, can be deflected by a magnet & electric field, have a very small mass. Proposed atoms made from negative electrons around a 'sea' of positive charge (plum-pudding atom).

Ernest Rutherford's gold-leaf experiment (1909-11): did experiments by directing a-particles towards a sheet of thin gold foil, they measured any deflection (disproved plum-pudding model). He proposed the new model, that most of the positive charge & mass is in the nucleus at the centre & negative electrons orbit the nucleus, but charges balance.

Niels Bohr's planetary model (1913): he altered Rutherford's model, to allow electrons to follow certain paths, orbit the nucleus in 'shells' and explains properties with the distance of electrons from the nucleus.

Henry Moseley's work on atomic numbers (1913): he discovered a link between X-ray frequencoes andatomic number.

Rutherford discovers the proton (1918): which explained Moseley's findings and tell us the number of protons is an atoms atomic number.

Wave and particle behaviour - Luois de Broglie & Erwin Schrodinger (1923 & 1926): suggested atoms have wave & particle properties, an introduced the idea of atomic orbitals.

James Chadwick discovers the neutron (1932): observed a new type of radiation of uncharged particles.

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Atomic structure & masses

The current model of an atom: protons & neutrons found in the nucleus, electron's orbit in shelld, the nucleus is tiny compared with the total volume of an atom (consists of empty space). Proton: 1 relative mass, 1+ relative charge. Neutron: 1 relative mass, 0 relative charge. Electron: 1/2000 relative mass, 1- charge.

Isotopes: atoms of the same element with a different number of neutrons. Different isotopes of the same element react in the same way, due to the fact chemical reactions involve electrons (they all have the same) & neutrons make no different to chemical reactivity.

Ions: a positively or negatively charged stom or (covalently bonded) group of atoms (a molecular ion). (Be,B,C & Si - dont form ions require too much energy to transfer outer shell electrons).

Ammonium (NH4 +), Hydroxide (OH -), Nitrate (NO3 -), Carbonate (CO3 2-), Sulfate (SO4 2-)

  • Relative isotopic mass: is the mass of an atom of an isotope compared with one-twelfth of the mass of an atom of carbon-12.
  • Relative atomic mass,Ar: is the weighted mean mass of an atom of an element compared with one-twelfth of the mass of an atom of carbon-12. (isotope 1 mass x %) + (isotope 2 x %) .... /100 - Mass spectrometry m/z
  • Relative molecular mass, Mr: the ratio of the average mass of one molecule of an element or compound to one twelfth of the mass of an atom of carbon-12.
  • Relative formula mass: is worked out be adding together the relative atomic masses of each atom making up a formula unit.


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  • A mole: is the amount of any substance containing as man particles as there are carbon atoms in exactly 12g of the carbon-12 isotope. 
  • The Avogadro constant, Na: is the number of atoms per mole of carbon-12 isotope. 
  • Molar mass, M: is the mass per mole of a substance (g mol -1).   moles = m (mass) /M (molar mass)
  • Molar gas volume: is the volume per mole of a gas (dm3 mol -1). At room temperature and pressure, the molar volume is approximately 24.   moles= V (volume in dm3) / 24

Ideal gas equation: gases behave in an ideal way; they are in continuous motion & do not experience any intermolecular forces, exert pressure when they collide & all collisions are elastic & kinetic energy of gases increases with temp.

pV=nRT (p-pressure, V-volume, n-moles, R-gas constant 8.314J mol -1K, T-temperature)

  • The concentration of a solution: is the amount of solute,in mol, dissolved per 1dm3 of solution.  moles= c (concentration) x V (volume in dm3). Mass concentration- amount of mass dissolved.
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Chemical equations

  • Empiral formula: the simplest whole-number ratio of atoms of each element present in a compound, always used for compounds with giant structures.
  • To calculate empirical formulae: divide the amount of each element present by its molar mass, divide the answer for each element by the smallest number. Make sure the ratio is in whole numbers.
  • Molecular formula: shows the numbers and type of the atoms of each element in a compound, does not tell you the order in which atoms are bonded to each other.
  • To calculate molecular formulae: use experimental results to find the molecular mass of a compound.

A chemical reaction: starts with the reactants, which change chemically to form the products. A species is a type of particle that takes part in a reaction, could be an atom, ion molecule, empiral formula or electron.

Percentage yield- with a fully balanced chemical equation it is assumed all of the reactants will converted to products, 100% yield. = actual amount, in mol, of product / theoretical amount, in mol, of product x100. There is not 100% yield because: reaction may be at equilibrium, side reactions may occur, reactants not pure, loss of product due to purification & some products left behind in apparatus.

Atom economy- considers not only the desired product, but also all the by-products of a chemical reaction, describing efficiency of a reaction in terms of all the atoms involved. = molecular mass of the desired product / sum of molecular masses of all products x 100. By using high atom economy reactions amount of waste is reduced (addition reactions).

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Acids, Bases & Salts

Acids: solution less than pH 7. It is a proton donor that releases H+ ions when added to water e.g. HCl (g) -> H+(aq) + Cl-(aq). Strong acids are good at giving up H+ ions, therefore they are fully dissociate / weak are partially dissociate. Examples: sulfuric acid( H2SO4) & nitric acid (HNO3).

Bases: is a proton acceptor, which neutralises acids. Examples: metal oxides (MgO), metal hydroxides (NaOH), ammonia (NH3) & amines (CH3NH2). Ammonia is a weak base / gas that dissolves in water to form a weak alkaline solution (only a small portion reacts with water).

Alkalis: (very corrosive) any substance that gives a solution greater than pH 7 when dissolved in water. They releas OH- (aq) ions e.g. NaOH (s) + aq -> Na+(aq) + OH-(aq). The hydroxide ions can neutralise the protons from acids e.g. H+(aq) + OH-(aq) -> H2O(l). Examples: sodium hydroxide (NaOH), potassium hydroxide (KOH) & ammonia (NH3).

Amphoteric substances: behave as acids & bases, for example an amino acid molecule - carboxyl acid group & amino base group.

Salts: is an ionic compound with a - positive ion (cation) a metal ion or an ammonium ion (NH4+), negative ion (anion) derived from an acid. Same as parent acid but H+ ion replaced by postive ion. 

  • Salts from carbonates- form a salt, carbon dioxide & water 
  • Salts from metal oxides- form a salt & water
  • Salts from alkalis- form a salt & water
  • Salts from metals- form a salt & hydrogen gas
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Crystallisation & Titrations

Water of crystallisation: refers to water molecules that form an essential part of the cystalline structure of a compound.

  • The dot formula gives the ratio between the number of compound molecules & the number of water molecules within the crystalline structure. Can be determined using experimental results & the empiral formula, gained from percentage or mass compositions.
  • You have the mass of hydrated salt e.g. MgSO4 : xH2O , mass of anhydrous salt e.g. MgSO4 & can be used to find the mass of water. First calculate the mol of anyhdrous compound and then calculate the mol of water. Use molar ratio and divide by the smallest number.

Titrations: the slow addition of one solution of a known concentration (titrant) to a known volume of another solution of unknown concentration, until the reaction reaches an end point (colour change). Use: a pipette, burette, conical flask & indicator.

  • methyl orange - in acid (red), in base (yellow) , end point (orange)
  • bromothymol blue - in acid (yellow), in base (blue) , end point (green)
  • phenolphthalein - in acid (colourless), in base (pink) , end point (pale pink/ colourless)
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Redox reactions

Oxidation numbers: are worked out by a set of rules:

  • uncombined element 0 / simple ion (charge on ion)
  • combine oxygen -2 / combined oxygen in peroxides -1 (H2O2)
  • combine hydrogen +1 / combined hydrogen in metal hydrides -1 (LiH)
  • combined fluoride -1
  • Compounds- the sum of the oxidation numbers must equal the overall charge of 0
  • Molecular ions- the sum of the oxidation numbers must equall its overall charge
  • Compounds of transition elements e.g. iron (II) chloride - Fe: oxidation number +2
  • Oxyanions- negative ions that contain an element as well as oxygen e.g. nitrate (V) NO3 - N: oxidation number +5

Redox reaction: a reaction in which oxidation and reduction both take place. Oxidation is the loss of electrons, increase in oxidation number. Reduction is the gain of electrons, reduction of oxidation number.

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Shells & orbitals

Energy levels or shells: the principle quantum number, n indicates the shell electrons occupy. 1st shell-2 , 2nd shell-8, 3rd shell-18, 4th shell-32.

Atomic orbitals: can hold a maximum of 2 electrons with opposite spins. An orbital is a region of space where electrons may be found. The two electrons in orbital must have opposite spins so the negative charges dont repel.

  • S-orbitals: has a spherical shape. From n=1, each shell contains one s-orbital 1s^2
  • P-orbitals: have a three-dimensionl dumb-bell shape. From n=2, each shell contains three p-orbitals  2p^6
  • D-orbitals: From n=3, each shell contains five d-orbitals 3d^10
  • F-orbitals: From n=4, each shell contains seven f-orbitals 4f^14

Sub-shells: is made up of one type of atomic orbital only. The sub-shells within a shell have different energy levels, 4s has a lower energy level then 3d so it is filled first. The lowest available energy level is filled first, and ecah orbital is filled singly before pairing starts.

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Ionic bonding & Polar bonds

Ionic bonding: the electrostatic attraction between oppositely charged ions. Electrons are transferred from the metal atom to the non-metal atom (metal ion positive & non-metal ion negative).

  • Giant ionic lattices: each ion is surrounded by oppositely charged ions, these ions attract each other from all directions, forming a three-dimensioned giant ionic lattice. All ionic compounds exist as giant inoic lattices in the solid state.
  • High melting & boiling points- a large amount energy is needed to break the strong electrostatic bonds that hold oppositely charged ions together in the solid lattice. The greater the charge the stronger the electrostatic forces.
  • Electrical conductivity- in a solid ionic lattice: ions are held in fixed positions, so do not conduct electricity. However when an ionic compound is melted or dissolved, the solid lattice breaks down and ions are free to move so conducts electricity.
  • Solubility- an ionic lattice dissolves in polar solvents. The polar solvents have polar bonds , which can break down the ionic lattice by surrounding each ion to form a solution.

Polar bonds: if bonding atoms are different e.g. in a covalent bond.The bonding atom with a greater attraction for the electron pair is said to be more electronegative (electronegativity increases across the periodic table)

  • A permanent dipole- is a small charge difference across a bond that results from a different in the electronegativitites of the bonded atoms.
  • If a molecule is symmetrical, polar bonds cancel and there is no overall dipole
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Covalent bonding

Covalent bond: a bond formed by a shared pair of electrons between nuclei. The negatively charged shared pair of electrons is attracted to the positve charged or both nuclei, the attraction overcomes the repulsion between the two positvely charged nuclei.Dot & cross diagrams are used to represent covalent bonds.

  • Single bond- is one shared pair of electrons
  • Multiple covalent bonding- atoms can share more than one pair of eletrons
  • Lone pairs- a concentrated region of negative charge around the atom.

Dative covalent bonding: a bond formed by a shared pair of electrons that has been provided by one of the bonding atoms only, also known as a coordinate bond. A-->B. 

Octet rule: states that elements gain or lose electrons to attain an electron configuration of the nearest noble gas. A better rule than the octet rule would be: unpaired electrons pair up, the maximum number of electrons that can piar up is equivalent to the number of electrons in the outer shell.

  • Simple molecular lattice /structures: made from small simple molecules held by strong covalent bonds, held by weak intermolecular forces. Have low mp/bp because intermolecular forces are weak and non conductors no free charged particles. Generally soluble in non-polar solvents, weak london forces form covalent bonds between these solvents.
  • Giant covalent structures; have high mp/bp because high temperatures are needed to break the strong covalent bonds. they are non-conductors except for graphite. They are insolublein any solvents, because the covalent bonds are too strong to be broken.
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Shapes of molecules & ions & intermolecular forces

Electron pair repulsion theory: the shape of a molecule is determined by its bonded pairs and lone pairs.

  • Linear : 1 or 2 bonded pairs (180 degree bond angle)
  • Non-linear: 2 bonded pairs & 2 lone pairs (104.5 degree bond angle)
  • Trigonal planar: 3 bonded pairs (120 degree bond angle) 
  • Pyramidal: 3 bonded pairs  & 1 lone pair (107 degree bond angle)
  • Tetrahedrel: 4 bonded pairs (109.5 degree bond angle)
  • Trigonl bipyramid: 5 bonded pairs (90 & 120 degree bond angles)
  • Octahedral: 6 bonded pairs (90 degree bond angle)
  • Each lone pair reduces the bond angle by about 2.5 degrees. Strengths of repulsion: lone pair/lone pair > bonded pair/lone pair > bonded pair/bonded pair

Intermolecular forces: is an attractive force between neighbouring molecules and can occur due to constant random movements. They are much weaker than chemical bonds.

  • Van der Waals' forces: dipole-dipole interactions (permanent dipole can induce a slight charge on a non-polar molecule) (molecules with permanent dipoles are attracted to other permanent dipoles) , London- dispersion- forces (random movements of electrons in atoms shells, inducing  dipoles in neighbouring molecules, more electrons bigger the force).
  • Hydrogen bonding: strong permanent dipole- permanent dipole attraction between an electron deficient hydrogen atom and a lone pair of eletrons on a highly electronegative atom.
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The periodic table & Ionisation energies

  • Periods: horizonatal row in the table, with each successive period following the same pattern (periodicity-repeated trends across a period)
  • Groups: vertical column, which contains elements with similiar properties. Each group has the same number of electrons in outer shell & same type of orbitals.

The first ionisation energy of an element is the energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ions. Successive ionisation energies are higher as each electron is removed there is less repulsion & a higher charge.

Factors affecting ionisation energy is: atomic radius (larger the atomic radius the smaller the nuclear attraction experienced by the outer electrons), nuclear charge (the higher the nuclear charge, the larger the attractive force) &  electron shielding/ screening (the repelling effect, the more inner shells there are the smaller the nuclear attraction).

  • Trends across a period: ionisation energies generally increase across each period because the number of protons/charge of the nucleus increases, electrons are added to the same shell & electron shielding doesnt change. There is a small decrease between group 2 and 13 because group 13 have p-orbital with higher energy so easier to remove. There is also a similiar decrease between group 15 and 16 because group 16 has a spin paired p-orbital.
  • Trends down a group: ionisation energies decrease as the number os shells increase, shielding increases & atomic radius increases.
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Group 2 elements (alkali earth metals)

Physical properties: have high melting & boiling points, have low densities & form colourless white compounds. 

Electron configuration & ionisation energy: have two electrons more than the configuration of a noble gas, reactivity increases down the group & ionisation energy decreases. They are reactive metals & strong reducing agents.

  • Group 2 elements & oxygen (redox reaction)- react vigorously & form ionic oxide
  • Group 2 elements & water (redox reaction)- moving down the group each metal reacts ore vigorously forms hydrogen & metal hydroxide
  • Group 2 elements & dilute acid- more vigorous moving down thr group & forms salt & hydrogen
  • Group 2 oxides & water- form metal hydroxides, that are soluble in water & form alkaline solutions with water, releasing OH- ions. The solubility of hydroxides increase down the group

Uses: calcium hydroxide ( Ca(OH)2 ) is used as lime to reduce acidity of soil , 'milk of magnesia' contains magnesium hydroxide which neutralises excess stomach acid & metal carbonates are useful for building materials.

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Group 7 (the halogens)

Properties: have low melting & boiling points & exist as diatomic molecules. They have 7 electrons in their outer shells. Moving down the group boiling point increases, and physical state changes from gas to liquid to solid, because there are more intermolecular forces as you move down the group.

They are also very reactive and electronegative, so they are strong oxidising agents. Reactivity decreases down the group as atomic radius increases, electron shielding increases & the ability to gain an electron decreases.

  • Redox reactions: occur between aqueous solutions of halide ions and aqueous solutions of halogens. A more reactive halogen will oxidise & displace a halide of a less reactive halogen (displacement reaction). The halogens form different-coloured solutions so colour changes indicate the redox reaction has occures, they are usually shaken with cyclohexane to distinguish between bromine and iodine.
  • Chlorine oxidises both Br- & I- (Bromine orange) (Iodine brown in water & violet cyclohexane)
  • Bromine oxidises I- (Iodine brown in water & violet cyclohexane)`
  • Disproportionation: is a reaction in which the same element is both reduced & oxidised
  • Reaction of chlorine with water (water purification)- forms hydrochloric acid & chloric(I) acid (HClO)
  • Reaction of chlorine with cold dilute aqueous sodium hydroxide (bleach formation)- forms NaCl + NaClO + H2O
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Testing for ions

Carbonate ions (CO3 2-): add a dilute strong acid, collect any gas formed & pass through limewater. Limewater goes cloudy & fizzing/ colourless gas is produced.

Sulfate ions (SO4 2-): add dilute hydrochloric acid & barium chloride. White precipitate of barium sulfate formed.

Halide ions (Br-, Cl-, I-): dissolved halide in water, add an aqueous silver nitrate. Silver chloride (white precipitate) / silver bromide (cream) & silver iodide (yellow). Add aqueous ammonia if colour hard to distinguish (iodide insoluble in ammonia).

Ammonium ions (NH4 +): add sodium hydroxide & warm gently. Turns red litmus paper blue.

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