Chemistry: Elements of Life (EL)

a Atoms contain:

Test yourself
• A nucleus, containing:
  - Protons, +1 charge and a relative mass of 1
  - Neutrons, 0 charge and a relative mass of 1
  - Electrons, -1 charge and a relative mass of 0.0005

Notation and the periodic table:

• Atoms are sometimes written with the notation where:
  - The mass number (A) is the total relative mass of the atom
  - The atomic number (Z) is the number of protons
  - x is the symbol of the element
• For example, in bromine, the atomic number A = 35 and
  
  

Properties relating charge, atomic number and relative mass:

• If the atom is neutral, the total charge is 0 so there will be an equal number of protons and electrons
• Because neutrons and protons both have equal relative masses, the number of neutrons = mass number - atomic number
• One mole of a substance (see b) contains about 6.02 × 10²³ atoms. This is the Avogadro constant, (given in the data sheet)

Isotopes:

• Most of the mass numbers on the periodic table are not whole numbers
• This is because there is usually a variation in the number of neutrons in naturally occurring elements
• For example, about 10% of lithium atoms contain three neutrons, and most of the rest contain four (you don't need to memorise this example)
• These atoms are called isotopes

Charged atoms:

• Atoms which are not neutral (same number of protons and electrons) are called ions
• Positively charged ions contain fewer electrons than protons and are called cations
• Negatively charged ions contain more electrons than protons and are called anions

Writing masses:

• Relative mass is the mass of something compared to ^th of a carbon-12 isotope (which contains 6 protons and 6 neutrons)
• Relative atomic mass, , is the mean relative mass of an atom (taking into account the different masses of its isotopes)
• Relative isotopic mass is the relative mass of a specific isotope. This is always an integer (whole number)
• Relative formula mass, , is the mean relative mass of an ionic or giant covalent substance
• Relative molecular mass, also , is the mean relative mass of a molecule

b Moles:

Test yourself
• A mole () is a unit to measure the amount of a substance
• One mole of a substance contains as many particles as there are atoms in 12 g of carbon-12
• Molar mass, M (measured in ), is the mass of one mole of something, so it has the same value as relative mass
• number of moles n () = mass m (g) molar mass M ()
• Aqueous solutions are mostly water, so only use the above equation for solids, liquids and gases
• If you need to work out the number of moles from the number of particles (e.g. atoms, ions, molecules), use
  number of moles = number of particles

Formulae:

• The molecular formula of a molecule shows the number of atoms of each element: for example, ethanol (CH₃CH₂OH) has the molecular formula C₂H₆O because it has 2 carbon atoms, 6 hydrogen atoms and 1 oxygen atom
• The empirical formula is the molecular formula but simplified down to the lowest possible whole numbers
  - For example, C₂H₄ would have the empirical formula CH₂
• Empirical formulae can be calculated from percentage compositions of each atom:
  - Calculate the mass of each atom if there's 100 g of the substance
  - Find the number of moles of each with the number of moles formula above
  - Divide the number of moles of each type of atom by the smallest
  - This gives the ratio for the empirical formula
  - However, you probably won't get exactly whole numbers due to rounding errors, so you need to round your final answer

Percentage yield:

• Theoretical yield is the amount of product if there are no losses. This can be calculated from the masses of reactants in a balanced equation
• Actual yield is the mass of product, measured by weighing it. It will be less than the theoretical yield because some will be lost, some of these causes are listed below:
  - Evaporation or losses to the environment
  - Chemicals not fully reacting
  - Solution being left in apparatus (e.g. small droplets on the inside of a beaker)
• percentage yield = actual yield theoretical yield × 100

Water of crystallisation:

• Hydrated compounds contain water of crystallisation, water molecules embedded inside (whereas anhydrous compounds don't)
• One mole of a particular hydrated compound always contains the same number of moles of water (but a mole of another hydrated compound may have a different number of moles of water)
• The number of moles of water is nearly always a whole number
• When heated, many hydrated compounds lose their water of crystallisation
• A hydrated compound is shown with a ·, like MgSO₄·7H₂O
• Example: 2.53 g of hydrated magnesium chloride (MgCl₂·xH₂O) was heated to constant mass. 1.17 g of solid remained. What is the formula of the hydrated compound?
  - The mass of water is equal to 2.53 - 1.17 = 1.36 g
  - This is equal to of water
  - The number of moles of anhydrous solid is
  - Divide both values by the number of moles of anhydrous solid:
  
  - Therefore, the number of moles of H₂O for one mole of hydrated solid is 1:6
  - This must always be rounded to the nearest whole number
  - So the formula is MgCl₂·6H₂O for a 1:6 ratio

Measuring the mass of a solid:

• Zero the balance (preferably a 2-3 decimal place one)
• Place a weighing bottle or boat onto the balance, add in the solid
• Record the weight of both the weighing bottle and solid
• Empty the solid into the container where it will be used (e.g. evaporating basin)
• Reweigh the weighing bottle (after zeroing the balance)
• Subtract this mass from the combined weight to find the mass of solid measured (excluding remnants on the weighting bottle)

c Concentration:

Test yourself
• In an aqueous solution, number of moles n () = concentration c () × volume V ()

Measuring volumes of liquids:

• A volumetric pipette is used to accurately dispense a fixed volume of a liquid:
  - Ensure that it is clean by rinsing it with water and then the solution to be measured
  - Dip the pipette into the solution and use a pipette filler to draw liquid into it
     - Stop once the bottom of the meniscus is touching the line on the pipette (view at eye level)
  - Transfer to the glassware the solution is being transferred to (e.g. conical flask)
  - There will be a small amount left at the tip, this must be left in the pipette
• A burette is used to accurately dispense a solution and record the exact volume:
  - First, clean the burette, rinsing with water and then a small volume of the solution being used
  - Use a funnel to empty the solution into the burette
  - Run a small amount into a beaker until there are no air bubbles
  - Record the volume to the nearest
  - Slowly open the tap, letting the solution out until the end point is reached
  - At the end point, close the tap and read again to the nearest
  - Subtract this from the first reading. This volume is called a titre
  - The total uncertainty will be 0.05 + 0.05 =

Preparing a standard solution:

• A standard solution has an exact, known concentration (e.g. )
• Making from a solid:
  - Calculate the solute mass required. Weigh (using the correct technique for weighing solids)
  - Pour of deionised water into a beaker
  - Empty in the solid and record the mass of solute used
  - Stir in the beaker until all of the solid has dissolved
  - Transfer to a clean volumetric flask
  - Rinse the beaker and pour into the volumetric flask
  - Add deionised water, swirling, until the level is about below the mark
  - Use a dropping pipette for the final , until the bottom of the meniscus lines up with the mark
  - Insert the stopper and invert several times, shaking to ensure an even mix
• Making from a more concentrated solution:
  - Rinse a beaker with the existing solution (stock solution)
  - Half-fill this beaker
  - With a pipette filler, rinse a pipette with some of the stock solution
  - Fill the pipette to (assuming we want a dilution factor of 10)
  - Empty into a volumetric flask
  - Add deionised water to mix the contents, and continue as with making from a solid

Acid-base titrations:

• Titrations can be used to accurately measure the concentration of an acid:
  - Rinse and fill a burette with the acid
  - Measure the start volume
  - Fill a pipette with the alkali solution and empty into a conical flask
  - Add 2-3 drops of an indicator (not universal, there needs to be a clear colour change at neutralisation)
  - Run the acid into the flask and swirl the flask until a colour change is seen
  - Record the end volume and note this as a trial titration
  - Keep repeating this until three titres are reached which are concordant
     - Concordant results must be within of each other
     - Find the mean of these - the mean titre

d Ionic equations:

Test yourself
• In an ionic equation, only the reacting ions and their products are shown
• Charges, as well as ions, must be balanced
• You should also include state symbols (unless the question states otherwise), which are written next to the ions/molecules with the same height. These are
  - aqueous (dissolved in solution), (aq)
  - liquid, (l)
  - solid, (s)
  - gaseous, (g)

e, f Shells:

Test yourself
• The first electron shell is called n=1, the second (farther from the nucleus than the first) is n=2, etc
• The core of an atom contains the nucleus as well as all electron shells except the outermost one
• Each shell can only hold a certain number of electrons, the first four are:
  - n=1: 2 electrons
  - n=2: 8 electrons
  - n=3: 18 electrons
  - n=4: 32 electrons
• The lowest energy shells are filled first (generally the ones closest to the nucleus)

Sub-shells:

• Each shell is split into one or more sub-shells
• There are four types of sub-shell, each with a maximum number of electrons:
  - s sub-shell: 2 electrons
  - p sub-shell: 6 electrons
  - d sub-shell: 10 electrons
  - f sub-shell: 14 electrons
• The sub-shells are filled the above order (s, then p, then d, followed by f)
• Generally, energy increases as shell and sub-shell type increases, although 3d has a higher energy level than 4s and 4s is therefore filled before 3d

Orbitals:

• Sub-shells are divided into orbitals
• Each orbital can hold either one or two electrons
• The maximum number of orbitals for each sub-shell type are as follows:
  - s sub-shell: 1 s orbital
  - p sub-shell: 3 p orbitals
  - d sub-shell: 5 d orbitals
  - f sub-shell: 7 f orbitals
• One full orbital is written as ⇅, since one electron will orbit clockwise and the other counter-clockwise. They fill singly first and must be in opposite directions; spin-pairing
• Electronic configuration is written like 1s²2s²2p⁶3s¹, with the large number meaning the shell number. The small superscript number is the electron count and the letters are the sub-shell type
• An orbital can take different shapes. An electron in an s orbital can be anywhere in a spherical region. Unlike many diagrams infer, electrons do not 'orbit'. Instead, they can be found in any location, most commonly in a certain area (close to the nucleus of a hydrogen atom, in a spherical zone)
• A p orbital has two regions of electron density, each side of the nucleus. Up to three can fit, one on each plane (x, y and z):
  

g The development of the model of the atom:

Test yourself
• Dalton: suggested that everything is made up of small solid spherical particles
• Thomson: discovered the electron, and created the 'plum pudding' model - negatively charged electrons embedded inside a positively charged sphere
• Rutherford: disproved the plum pudding model with the nuclear model - a nucleus surrounded by electrons with evidence from the gold foil experiments (see below)
• Mosely: noticed that the charge increased by exactly one for each nucleus - and proposed the idea of protons
• Chadwick: discovered the neutron
• Bohr: proposed a modified model to the nuclear model; electrons in fixed orbits, each with a fixed amount of energy, explaining why the noble gases are inert and how many atoms react
• The gold foil experiments, conducted by Rutherford, Geiger and Marsden, found that the mass of an atom is concentrated in a small space - the nucleus. They did this by firing alpha particles (positive helium nuclei) through thin gold foil. Many passed straight through (which couldn't happen if the plum pudding model were correct)

Evidence for electron shells:

• First ionisation energy is the energy required to remove the weakest electron from an atom. The trends (see q) provide evidence towards Bohr's model
• Atomic emission spectra vary with different atoms, showing that their electronic structures are different
• Additionally, the new model explains why the noble gases are inert (unreactive) due by full electron shells

h Nuclear fusion:

Test yourself
• Nuclear fusion is where two light atomic nuclei join together to form one or more different nuclei, and sometimes protons or neutrons
• This releases a lot of energy as gamma radiation. This is electromagnetic radiation (i.e. photons)
• However, fusion only happens at high temperatures and pressures
• These conditions are needed so that the repulsion of nuclei is overcome by the high kinetic energy of atoms
• Many elements are formed by nuclear fusion reactions in stars
• Here is an example of a fusion reaction:
• In this reaction, a hydrogen nucleus and a deuterium nucleus fuse together to form a helium nucleus
• The sum of proton numbers and mass numbers must be equal on both sides
• Sometimes, protons or neutrons are formed, these can be written as
  - proton:
  - neutron:
• Sometimes, the energy released is written as (gamma)

i Covalent bonding:

Test yourself
• Electrostatic attraction is the attraction between a positive and negative particle (e.g. proton and neutron)
• Electrostatic repulsion is the repulsion between two particles of the same charge (e.g. two electrons)
• Covalent bonds are shared electrons between two non-metal atoms
  - Although the electrostatic attraction is strong within covalently bonded molecules, the inter-molecular forces are low, resulting in relatively low melting and boiling points
  - The nuclei repel, so there must be a balance between these electrostatic repulsion forces and the electrostatic attraction
• A single bond is a covalent bond made up of one electron from each atom
• A double bond is two covalent bonds (two pairs of electrons, e.g. the bonds in oxygen molecules)
• This trend continues (e.g. 3 electron pairs for a triple bond)
• In a dative bond, both bonding electrons are from the same atom
• Lone pairs are outer shell electron pairs not involved in bonding. For example, the O in H₂O has two lone pairs because only two of its outer shell electrons are involved in covalent bonds

Ions:

• Ions have a different number of electrons to the their proton count, resulting in the particle having an overall charge
• For example, Cl^- has 18 electrons and 17 protons, resulting in an overall charge of 17 - 18 = -1

Dot-and-cross diagrams:

• Dot-and-cross diagrams can be used to show how the electrons are shared in a covalently bonded molecule
• Use dots to represent electrons from one atom and crosses to represent electrons from another
• If your molecule has more than two atoms, assign each dots or crosses based on how clear it would be
• You only need to show the outer shell electrons
• The shells should overlap to show that electrons are shared
• The dot-and-cross diagram for a Cl₂ molecule is shown below:
  

Dot-and-cross diagrams for ionic substances:

• Use square brackets around each ion with its charge. The dot-and-cross diagram for CaO is shown below:
  

j Ionic bonding:

Test yourself
• Ionic bonding is when ionic compounds join together with alternating positive and negative ions in all three dimensions to form a lattice structure
• It happens between metals and non-metals
• The high electrostatic forces give ionic lattices high melting and boiling points
• They can conduct electricity when molten or dissolved in water
• They are arranged in a positive-negative pattern because this reduces repulsions and maximises attractions
• The structure of a small LiCl lattice is shown below:
  

Metallic bonding:

• Metals also have lattice structures. Closely packed positive metal ions (cations) are attracted to 'free' delocalised electrons (negative) which are able to move and carry a charge (hence why solid metals conduct)
• This strong attraction causes a high melting point for all metals except mercury
• Metals are insoluble due to the strength of the metallic bonds
• They are very ductile because the metal ions can slide over each other under stress
• A diagram for a group 1 metal is shown below:
  
• Group 2 metals will have two delocalised electrons per cation

Covalent networks:

• Covalent networks have strong covalent bonds between the atoms, leading to high melting and boiling points. Covalent lattices don't conduct electricity generally because the electrons are locked in the network structure (except graphite which does have delocalised electrons)
• They are good thermal conductors as vibrations travel easily due to the lattices' stiffness
• They are not soluble in polar solvents like water, because the atoms are more attracted to each other than the solvent. This shows that they don't contain ions
• Examples include graphite and diamond (carbon) and silicon dioxide (SiO₂)
• These are generally only formed from group 4 elements

k Outer shells:

Test yourself
• Generally, atoms in covalent molecules have full outer shells of 8 electrons, the octet rule
• However, there are some exceptions
  - Boron in BF₃ only has 6 electrons in its outer shell (it's electron deficient)
  - Sulfur in SF₆ has 12 electrons in its outer shell

Bond angles:

• Electrons are all negatively charged, so they repel
• Lone-pair—lone-pair angles are the largest, followed by lone-pair—bonding-pair and bonding-pair—bonding-pair
• The main bond angles you need to know are shown below:
• Two bonds around the central atom: Linear, 180°, e.g. BeCl₂, CO₂
• Three bonds: Trigonal-planar, 120°, e.g. BF₃
• Four bonds: Tetrahedral, 109.5°, e.g. CH₄
• Five bonds: Trigonal-bipyramidal, 120° and 90°, e.g. PF₅
• Six bonds: Octahedral, 90°, e.g. SF₆
• Two bonds and one lone pair: ~120°, e.g. SO₂, NO₂ (because the extra electron density in the double bonds cancels out the repulsion from the lone pair)
• Two bonds and two lone pairs: Bent, 104.5°, e.g H₂O
• Three bonds and one lone pair: Pyramidal, 107°, e.g. NH₃
• Four bonds and two lone pairs: Square planar, 90°, e.g. XeF₄

l Ionic compounds:

Test yourself
• Some ionic compounds consist of two atoms, one positive and one negative
  - For example, sodium chloride, NaCl, has a sodium ion (+1 charge) and a chloride ion (-1 charge). This is because the sodium atom has lost an electron to the chlorine atom
  - With some compounds, such as magnesium fluoride, the magnesium atom needs to lose two electrons for a stable outer shell, so it ionically bonds to two fluorine atoms for MgF₂
• Although atomic radius decreases across the periods due to higher proton-electron attraction, a chloride ion is larger than a sodium ion because sodium loses an electron and shell whereas chlorine gains an electron and keeps three shells
• Ionic compounds can have more than two elements, for example oxygen and hydrogen can covalently bond together and gain an electron from sodium to form NaOH
• The charges must always sum to zero
  - For example, the charge of Na is +1 and OH is -1, so the sum = +1 + -1 = 0

m The periodic table:

Test yourself
• The periodic table is a list of elements in order of proton number
• It groups elements according to their properties
• Groups are columns and the group number is the number of outer shell electrons of elements in that group
• Periods are rows and the period number is the number of electron shells the atoms in it have
• The s-block elements are those in groups 1 and 2, with electronic structures ending in s¹ or s²
• The p-block elements are those in groups 3-8, with an electronic structure ending in a p
• The d-block elements are those in the transition metal block, with an electronic structure ending in a d

n Melting point trends in periods 2 and 3:

Test yourself
• The periodic table has many trends, called periodicity
• In periods 2 and 3, melting points increase across the period, up to group 4 (carbon/silicon) where it starts to decrease
• The increasing melting point trend is due to the metal—metal bonds becoming stronger due to higher charge density (more delocalised electrons and a smaller ionic radius)
• Carbon and silicon form giant covalent structures, so they have an even higher melting point
• The simple molecular structures (e.g. N₂) have lower melting points because their intermolecular forces are weak
• Larger molecules (e.g. S₈ and P₄) have stronger intermolecular forces and therefore higher melting points than smaller molecules such as O₂, resulting in a decrease in melting point from groups 5 to 8
• Noble gases (e.g. Ne) are monatomic (one atom in a molecule), leading to a very low melting point
• The trend is shown below:
  

o Ions and the periodic table:

Test yourself
• Group 1 elements form +1 ions
• Group 2 elements form +2 ions
• Group 6 elements form -2 ions
• Group 7 elements form -1 ions
• This is due to the number of electrons in their outer shells

Commonly cations (+ve):

• +1:
  - Hydrogen ions (H⁺), same for lithium (Li⁺), etc
  - Ammonium (NH₄⁺)
• +2:
  - Magnesium (Mg²⁺)
  - Calcium (Ca²⁺)
  - Barium (Ba²⁺)
  - Iron(II) (Fe²⁺)
  - Copper(II) (Cu²⁺)
  - Zinc (Zn²⁺)
  - Lead (Pb²⁺)
• +3:
  - Aluminium (Al³⁺)
  - Iron(III) (Fe³⁺)

Common anions (-ve):

• -1:
  - Flouride (F^-), same for chloride (Cl^-), bromide (Br^-) etc
  - Nitrate (usually NO₂^- or NO₃^-, depending on the oxidation state given (see ES [year 1] and CI [year 2]))
  - Hydrogencarbonate (HCO₃^-)
• -2:
  - Oxide (O^2-)
  - Carbonate (CO₃^2-)
  - Sulfate (SO₄^2- unless an oxidation state is given (you'll see this in ES))

Naming ionic substances:

• Find the names above and combine them, giving the cation (+ve) first and anion (-ve) last
• For example, CaO contains Ca²⁺ and O^2-. Therefore, the name is calcium oxide
• The same applies if an oxidation state Roman numeral is given; e.g. FeCl₃ is iron(III) chloride

p Group 2 oxides and hydroxides:

Test yourself
• Oxides of group 2 have the general formula MO
• Hydroxides of group 2 have the general formula M(OH)₂
• In water, oxides and hydroxides form alkaline solutions, M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g), although they are not very soluble
• Metal oxides are formed when group 2 metals react with oxygen, 2M(s) + O₂(g)→ 2MO(s)

Group 2 carbonates:

• Group 2 carbonates have the general formula MCO₃. When they decompose, the oxide and CO₂ is formed; MCO₃(s) → MO(s) + CO₂(g)
• The carbonates become more difficult to decompose down the group, explained by thermal stabilities (see r)

Going down group 2 from Mg to Ba:

• The reactivity with water increases
• Thermal stability of the carbonate increases
• The pH of the hydroxides in water increases
• Solubility of the hydroxides increases
• Solubility of the carbonates decreases

q Ionisation enthalpy:

Test yourself
• First ionisation enthalpy is the energy required to remove the most loosely held electron from one mole of gaseous atoms, producing one mole of gaseous ions with a +1 charge
• It is measured in
• It is always positive (endothermic)
• The general equation for the first ionisation process is: X(g) → X⁺(g) + e^-
• For oxygen it would be O(g) → O⁺(g) + e^-

Factors affecting ionisation enthalpy:

• Atomic radius: larger atoms have lower ionisation enthalpy (the electrons are easier to remove)
• Nuclear charge: if there are more protons and electrons, ionisation enthalpy increases due to higher proton-electron attraction
• Electron shielding: inner electron shells partially shield the outer one from the positive charge from the nucleus, reducing ionisation enthalpy

Periodic trends in ionisation enthalpy:

• First ionisation enthalpies decrease down the groups, because the attraction between the nucleus and outer electron is weaker
• They increase across the periods, because there are more protons, and no change in electron shielding as all elements in a period have the same number of electron shells

r Charge density:

Test yourself
• Charge density is the measure of the concentration of charge on the cation
• Across a period from groups 1 to 3, the overall charge increases and atomic radius slightly decreases, so the charge density increases
• Down a group, the atomic radius increases and overall charge remains the same, so charge density decreases
• Therefore, group 2 metals with a high charge density are at the top of the group

Group 2 carbonate thermal stability:

• The smaller the +2 metal ion, the higher the charge density and therefore the higher the distortion or polarisation (see the diagram below)
  
• Smaller ions with higher charge density distort the large carbonate ion, so the compound decomposes at lower temperature - it has a lower thermal stability

Testing this experimentally:

• Heat the carbonate in a test tube at a constant rate
• Pass the gas collected through limewater, and determine how long it takes to go cloudy (e.g. put a pencil cross under the limewater test tube and record the time for it to become invisible)
• Limewater cloudiness indicates the presence of carbon dioxide
• Further down group 2, it will take longer to turn cloudy
• Ensure to use:
  - The same volume of limewater
  - The same heating conditions
  - The same number of moles of carbonate used
• Diagram:
  

s Solubilities:

Test yourself
• Most ionic solutions dissolve in water, such as:
  - Lithium, sodium, potassium and ammonium salts
  - Nitrates
  - Chlorides, bromides and iodides (except lead and silver halides)
  - Sulfates (except barium, calcium, lead and silver sulfates)
  - Lithium, sodium, potassium, strontium, calcium, barium and ammonium hydroxides
• Additionally, all metal carbonates are insoluble in water

Salts:

• Salts are formed with the following equations
  - acid + base → salt + water
  - acid + carbonate → salt + water + carbon dioxide
  - acid + metal → salt + hydrogen
• When ionic substances dissolve, the ions are surrounded by water molecules and are distributed evenly in the solution. Aqueous salt solutions therefore conduct electricity because of these hydrated ions which are free to move and carry a charge
• When writing ionic equations, only the reacting ions are included - all others are not written and are called spectator ions
• Precipitates are suspended solid particles (not dissolved) in a solution

Testing for metal ions:

• Add sodium hydroxide (containing OH^-)
• Ag⁺ forms Ag₂O (brown) and water
• Ca²⁺ forms Ca(OH)₂ (white)
• Cu²⁺ forms Cu(OH)₂ (blue)
• Pb²⁺ forms Pb(OH)₂ (white)
• Fe²⁺ forms Fe(OH)₂ (green)
• Fe³⁺ forms Fe(OH)₃ (red-brown)
• Zn²⁺ forms Zn(OH)₂ (white), which redissolves with another OH^- to Zn(OH)₃^- (colourless)
• Al³⁺ forms Al(OH)₃ (white), which redissolves with another OH^- to Al(OH)₄^- (colourless)

Testing for halides:

• Add dilute nitric acid (HNO₃) and then silver nitrate solution (AgNO₃):
  - Chlorides: white AgCl
  - Bromides: cream AgBr
  - Iodides: yellow AgI
• First, ensure that a sulfate is not present as this will also produce a precipitate with silver nitrate. So test for sulfates before testing for halides
  - Alternatively, add a dilute acid first to remove any unwanted anions

Testing for lead ions:

• Pb²⁺ + potassium iodide → lead iodide (yellow)

Testing for sulfates:

• Add dilute HCl and then barium chloride (BaCl₂) or barium nitrate (Ba(NO₃)₂)
• Sulfates (SO₄^2-) form a white barium sulfate precipitate
• First, ensure that there are no carbonate or sulfite ions (SO₃^2-), since barium carbonate and barium sulfite are also insoluble. Therefore, always test for carbonates before testing for sulfates

Litmus paper:

• Litmus paper comes in two forms:
  - red litmus paper turns blue in alkaline solutions
  - blue litmus paper turns red in acidic solutions

Testing for ammonium (NH₄⁺):

• Add sodium hydroxide and heat:
  NH₄⁺ + OH^- → NH₃ + H₂O
• Ammonia, NH₃, is alkaline and is a gas
• Therefore, red litmus paper will turn blue
• Ensure that the litmus paper is dampened so that the ammonia can dissolve

Testing for hydroxides:

• Dip red litmus paper into the solution. It will turn blue if a hydroxide is present
  - However, it will also turn blue if another alkali is present, so test for those first

Testing for nitrates:

• Warm the solution and add sodium hydroxide and aluminium
• Aluminium reduces the nitrate ions to ammonium ions which react with the hydroxide ions, creating ammonia gas and water
• Now test the gas given off with the ammonium test above

Optimal testing order:

• To avoid false positives, aim to use the following order:
  - test for carbonates
  - test for sulfates
  - test for halides
• It can also be useful to add a dilute acid to the solution, to remove any anions that could interfere

t Acids and bases:

Test yourself
• An acid is a compound that produces hydrogen ions (H⁺) in water. They donate ions to bases, making them proton donors
• Examples of mineral acids (acids derived from inorganic compounds):
  - HCl, hydrochloric acid
  - H₂SO₄, sulfuric acid
  - HNO₃, nitric acid
• A base is a compound that reacts with an acid to produce water and a salt. A base is a proton acceptor
• An alkali is a base that dissolves in water to produce hydroxide ions (OH^-)
• A neutralisation reaction is when a base reacts with an acid to form a salt

Making a salt from an acid and alkali:

• 1. Carry out an acid-base titration to find out how much acid is required to neutralise of the alkaline solution
• 2. Measure another of alkali. Using the burette, add the amount of acid that is required to neutralise the alkali
• 3. Transfer the neutralised solution to a clean evaporating basin and heat lightly (to avoid spitting) to evaporate the water
• 4. Leave the mixture to cool in the evaporating basin
• 5. Filter the mixture, and wash the solid with cold distilled water
• 6. Transfer the residue to a watch glass and leave it to dry. Measure its mass at regular intervals; you can stop once it remains constant (i.e. no more water is evaporating)

Making a salt from an acid and a base:

• 1. In a beaker, warm excess of the insoluble base in dilute acid
• 2. Continue to heat, then add universal indicator to see if the solution is neutral. If required, more solid base can be added
• 3. Leave to cool
• 4. Filter off the excess base and transfer the filtrate to a clean, dry evaporating basin
• 5. Heat the evaporating basin until salt crystals begin to appear on the sides of the basin

Making water-insoluble salts with precipitation reactions:

• 1. Add the desired solutions into a beaker to form a precipitate of the salt
• 2. Filter the precipitate
• 3. Wash the precipitate with deionised water
• 4. Transfer to a watch glass and dry in an oven (below the melting point of the salt), or in air
• 5. At regular intervals, weigh
• 6. Once the solid has dried to a constant mass, you can stop heating

u Group 2 oxides and hydroxides:

Test yourself
• Oxides and hydroxides are bases
• They react with acids to from salts and water:
  - MO(s) + 2HCl(aq) → MCl₂(aq) + H₂O(l)
  - M(OH)₂(s) + H₂SO₄(aq) → MSO₄(aq) + 2H₂O(l)

v The electromagnetic spectrum:

Test yourself
• In A Level chemistry, the three parts of the electromagnetic spectrum from lowest to highest (shortest to longest) wavelength are:
  - ultraviolet
  - visible [violet → red]
  - infrared
• Due to the equation in w, you can calculate that the above list shows the frequency and energy from highest to lowest

w Light:

Test yourself
• Although light is a wave, it can also be considered as being made up of particles, photons
• Waves have a frequency and a wavelength, wave speed c () = wavelength (m) × frequency ( or )
• energy of a photon Δ E () = Planck constant h × frequency ( or )

Bohr's theory:

• When heated, Bohr's theory explains that:
  - The atom gains energy and its electrons become 'excited'
  - This causes electrons to jump to higher energy levels, further from the nucleus than at their neutral states
  - Later the electrons will drop back to lower levels, giving off a single photon of energy, often as visible light
• Bohr's theory is that electrons exist in fixed, quantised energy levels. The ground state is n = 1, the closest to the nucleus. As n increases, the energy values converge (get closer together) because at higher energy levels, there is less nucleus—electron attraction (since there's an inverse square law; doubling the distance reduces the force by 4×)
• Because the energy for each state is different for different atoms, they can be plotted in emission spectra and used to identify atoms

Emission spectra:

• Emission spectra show vertical lines of colour from the emissions on a black background
• An absorption spectrum is similar, but contains black lines (in the same places) on a coloured background. It represents the wavelengths of electrons absorbed by electrons increasing their energy levels
• Frequency increases from right to left (red has a lower frequency and higher wavelength than violet), and wavelength decreases from right to left
• The lines converge on the left of the spectra for the reasons in the above section
• The spectrum for hydrogen is shown below (each line is labelled with its corresponding wavelength):
  

Flame tests for common metal ions:

• Dip a nichrome wire loop in concentrated hydrochloric acid and then into a sample of the compound. Hold this in a flame to test the compound. One of the following colours will be produced:
• Li⁺: bright red
• Na⁺: yellow
• K⁺: lilac
• Ca²⁺: dark red
• Ba²⁺: green
• Cu²⁺: blue-green

x Calculating relative atomic mass from mass spectra:

Test yourself
• The y-axis of a mass spectrum is the % abundance, and the x-axis is mass/charge or m/z. This is usually just the mass however, as charge is mostly +1 (the mass spectrometer ionises the sample)
• The relative atomic mass can be calculated by multiplying the relative isotopic mass by relative abundance (%) for each ion mass detected (for each bar on the graph). The sum of these should then be divided by 100 to get the average mass of each atom of this type