The chemical Earth




The chemical Earth


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The chemical Earth


Module 1- The Chemical Earth
Week 1/2:

  • Homogeneous substances are substances of uniform composition e.g. Pure water
  • Heterogeneous substances are substances of variable composition e.g. Tap water
  • Impure substance is a substance contaminated with other substances (mixture)
  • A compound is a pure substance that can be decomposed into smaller substances
  • They are broken down into elements (which cannot be decomposed)
  • Lithosphere (crust+ top part of mantle)- contains rocks, sand, soils, minerals
  • Hydrosphere (water of the Earth's crust)- contains water, sea water
  • Atmosphere (layer of gas above the Earth)- contains gases, nitrogen, oxygen, argon
  • Biosphere (inhabited portion of Earth)-contains living matter, organisms, plants
  • Sieving can separate solids of different sizes
  • Filtration, sedimentation, decanting can separate solids and liquids
  • Evaporation can separate dissolved solids in liquids
  • Distillation can separate two liquids if they have a different boiling points.  Fractional distillation is utilised when boiling points are <40 degrees in difference.
  • Separating funnels can be used to separate liquids of different densities
  • A solvent can be added to a mixture where one substance is soluble
  • Gases can be separated by using different boiling points or solubilities

Week 3:

  • Common properties of purse substances are colour, physical state, melting/boiling points, density electrical conductivity, solubility in different liquids and mechanical properties.
  • Distinct colours e.g. Deep brown- liquid bromine, solid copper- reddish brown etc.
  • States of matter are solid, liquid or gas
  • Changes from solid to liquid to gas and vice versa are called changes of state
  • The melting point of a solid is the lowest temperature at which the solid changes to a liquid
  • Melting points are sharp unless it is a mixture.
  • Opposite of melting point is freezing point which is exactly the same
  • Lowest temperature at which bubbles of vapour form is called the boiling point.
  • Boiling points are quite sharp can test purity of a substance
  • Converting a gas to a liquid is condensation or liquefaction
  • Density is defined as mass per unit volume e.g. Kg/meter cubed
  • Gravimetric analysis is determining the masses of substances in a sample
  • Free elements include oxygen and nitrogen, including noble gases
  • Most elements form as compounds because they are reactive
  • The more reactive an element is, less chance of finding it as an element
  • Metals are elements which are solid, shiny, good conductors and malleable/ductile
  • Semi-metals display characteristics of both metals and non-metals.
  • Elements are used due to their physical properties e.g. Aluminium- light but strong, copper-high electrical conductivity, liquid nitrogen- low freezing and boiling points.
  • Vertical columns of a Periodic Table are called groups (Valence electrons)
  • Horizontal rows of a Periodic Table are called periods (number of shells)
  • Two elements are liquids (mercury/ bromine), 11 are gases (noble gases + fluorine+ chlorine) while the rest of them are solids at room temperature.

Week 4:

  • Matter is made up of small particles
  • In solids, particles are packed tightly, liquids, less tightly and in gases, they are free
  • When heated, energy is transferred to the particles, causing them to vibrate
  • Atom is the smallest particle of an element
  • A molecule is the smallest particle of a substance that can exist alone.
  • Symbols are used to represent elements- lowercase for elements and uppercase for compounds
  • Combination of symbols is called formulae.
  • A pair of atoms bonded in a molecule is called diatomic (O2) and one atom is called monatomic, usually the noble gases.
  • In an atom, there exists a small dense nucleus and several electrons rapidly moving around
  • Electrons move randomly in a volume described as the electron cloud
  • Atomic number of an element is the number of protons, mass number is protons + neutrons
  • Each energy level can hold a certain amount of electrons, where 1st= 2, 2nd= 8, 3rd= 8 or 18 etc.
  • Most elements like to obtain its nearest noble gas electron configuration- losing or gaining electrons to become an ion.
  • Electrons in the highest energy level (valence shell) are called valence electrons.
  • Formulae that give the ration of atoms by elements are called empirical formulae
  • Positive ions are called cations.  Negative ions are anions
  • Elements in the same group tend to either gain or lose the same amount of electrons, giving it similar properties (except for Group 4)
  • A covalent molecule is when an element shares electrons to obtain noble gas configuration
  • Substances made from covalent molecules are called covalent molecular substances
  • Lewis dot structure is used to show electron configuration visually (see right)
  • Forces between pairs of molecules are called intermolecular forces (usually weak)
  • Covalent network solids are solids where covalent bonding occurs (covalent lattices)
  • Delocalised electrons hold positive ions in a lattice- making it malleable and ductile yet hard.

Week 5:

  • Physical change- no new substance is formed
  • A change in which at least one new substance is formed is called a chemical change
  • Also called chemical reactions
  • Usually defined by a gas being produced, a precipitate, change in colour, temperature change, disappearance of a solid or an odour is produced.
  • In a chemical reaction, starting substances are reactants and products are formed
  • Mass and the number of atoms are conserved (Law of conservation of matter)
  • Physical properties e.g. melting and boiling points, chemical properties e.g. reactivity, effect of light and decomposition of heat
  • Substances can be decomposed by heating them e.g. copper nitrate or through electricity e.g. electrolysis of water, or through light e.g. silver salts
  • In physical change- particles are not changed, only in a chemical reaction
  • Direct combination reactions is when the elements react to form a compound e.g. magnesium oxide.
  • The stronger a chemical bond is, the more energy is required to break the bonds.
  • Word equations are used to represent a reaction with symbolic equations- they must be balanced.

Week 6:

  • The number of positive charges are equal to the negative charges in ionic compounds.
  • The valency is the numerical value of the charge that the ion of the element carries.
  • Binary compounds consist of two elements only.  Ionic binary compounds consist of two ionic substances.  Named with positive ion than negative ion.
  • Compounds with three or more atoms are called polyatomic ions such as sulfur.
  • The valency of a covalent compound is the number of covalent bonds the element can form.

Module 2: Metals/ The Mole
Week 1:

  • Alloys are homogeneous mixtures of a metal with one or more other elements.
  • Common examples are brass (copper+ zinc), bronze (copper+ tin), solder (tin+ lead), stainless steel (iron+ chromium or nickel), steel (iron+ carbon)
  • Mild steel contains less than 0.2% carbon and is soft, structural steel contains up to 0.6% carbon and is hard, while high-carbon steel contains up to 1.5% carbon and is very hard.
  • Electrical conductivity is the current passing through a metre cube when a voltage of 1 volt is applied across the opposite faces (megohm-1m-1)
  • Thermal conductivity is the energy transmitted per second through a metre cube of the substance when there is a 1 degree temperature difference (Js-1m-1K-1)
  • Hardness can be measured through Mohs scale: from 1-10
  • Tensile strength is a measure of how well a material resists bending, twisting or stretching.
  • Iron (usually in mild steel or structural steel) can be used for buildings, as it is cheap and malleable, however it corrodes so it needs to be painted or galvanised with zinc.
  • Aluminium alloyed with copper, magnesium or manganese is light and strong, and does not corrode.  It is used for aeroplanes, windows, door frames, household utensils and drink cans.
  • First metal to be extracted was copper (copper oxide heated with charcoal) but it was soft.  Mined from 3000BCE, called the copper age.
  • When bronze was discovered (tin and copper ores) it became a favourite metal (hard, low melting point and malleable) for ploughs and weapons (2000BCE)
  • Iron was extracted from iron oxide using a high temperature from 1200BCE which were used for tools and weapons.  This was known as the Iron Age.
  • Steel and iron increased during the Industrial Revolution with new alloys and metals in the Modern Age e.g. cobalt, nickel, zinc, manganese, tungsten, titanium and chromium.
  • A mineral is a pure crystalline compound that occurs in the Earth's crust.
  • An ore is a compound or mixture of compounds from which it is economic or commercially profitable to extract a desired substance such as a metal.

Week 2-3:

  • The reactivity of metals is ranked on an activity series, with oxygen or other substances. 
  • Group 1 metals react rapidly with air, Group 2 react slowly but burn rapidly in air or oxygen.  Group 4 metals react slowly with air but only if heated forming ionic compounds.
  • Similarly the metals react with water in different ways as shown above.
  • A net ionic equation shows the actual ionic species that undergo change in the reaction.
  • A complete ionic equation shows all ions involved in the solution, including spectator ions which do not undergo any change during a reaction.
  • Metals that react with oxygen, water and dilute acids lose electrons to form positive ions, known as oxidisation reactions.  The opposite is a reduction reaction known as electron transfer reactions.
  • When there is a oxidisation and reduction reaction, it is called a redox reactions or a electron-transfer reaction.
  • Half-reactions are reactions that describe the oxidation and reduction processes separately in terms of electrons lost or gained.
  • An activity series is a list of metals in order of decreasing reactivity with oxygen, water or dilute acids. 
  • The first ionisation energy of an element is the energy required to remove an electron from a gaseous atom of the element, usually measured in kilojoules per mole of atoms. 
  • Reactivity of metals concerns the environment it is used in e.g. piping, body implants, roof guttering- all of which use inert metals so they do not react with water etc.
  • Aluminium used to be an expensive metal due to its extraction process.  Now Aluminium can be converted to alumina (Al2O3) through the Bayer Process and melted with cryolite and then electrolysed, the aluminium ions would be reduced to metal at the cathode.  However high cost of electrolysis and melting made it expensive.  Now it is much cheaper
  • Alternative is to recycle aluminium, it uses less than 10% of the production energy costs.

Week 4-7:

  • Stoichiometry is the study of quantitative aspects of formulae and equations- stoichiometric equations are calculations based on this aspect.
  • Isotopes have a different mass of neutrons in the nucleus hence have different masses.
  • Atomic weight is the average mass of the element in the naturally occurring element
  • Molecular weight of a compound is the sum of the atomic weights of the atoms
  • Formula weight is the sum of the atomic weights in a compound.
  • A mole of an element contains 6.022 x 10^23 atoms, same for a compound.  Also known as the Avagadro constant, number of atoms in 12 grams of carbon.
  • Molar mass is the mass of a mole, e.g. oxygen's molar mass is 16grams/mol.  Be careful of molecules and atoms, oxygen molecules is 32g/mol, oxygen atoms are 16g/mol
  • No. of moles = mass/ molar mass
  • No. of atoms = no. of moles x Avagadro constant.
  • Percentage compositions are used to determine the composition of an ore e.g. In Fe2O3, 70% is iron.  % of iron = mass of iron in one mole/ mass of one mole of Fe2O3
  • Empirical formula tells us the ratio in which the atoms are present (simplified ratio).  Can be calculated by finding the moles of each substance, dividing by the smallest amount and then rounding them or making them whole numbers.
  • The molecular formula is the unsimplified empirical formula e.g. C6H12O6 is molecular, while empirical would be CH2O.  To calculate have molecular mass/ empirical mass x Empirical formula.
  • Write a balanced chemical equation- then find how much moles are needed, using stoichiometry.
  • Look for the limiting reagent- a substance that runs out before the other does.  Calculated from the stoichiometry. 
  • The yield of a metal is the mass of metal that is obtained from an ore.  In theoretical calculations we assume 100% yield, but there are real life problems like precipitating all of the mineral out, and we can't achieve 100% efficiency.  Using the yield percentage at the end of the solution from your theoretical yield to your actual yield.
  • Gay-Lussac and Avagadro combined- one volume reacts with one volume to form two volumes, works between hydrogen and chloride, two volumes reacts with one volume to form two volumes e.g. 200ml hydrogen + 100ml oxygen = 200ml water

Week 8:

  • Dobereiner developed triads of elements that had similar properties
  • Newlands proposed a law of octaves of increasing atomic weight.
  • Mendeleev was the forerunner of the modern Periodic Table as was Meyer who arrange in order of increasing atomic weight, those that had similar properties were called the periodic law.  Mendeleev was very successful, recognising gaps were elements were not yet discovered.
  • However with atomic weight there were discrepancies, Moseley rearranged them by atomic number and proposed the modified periodic law, properties of elements vary periodically with their atomic numbers.
  • Atomic radius decreases across a period (as more protons slightly shrink the radius) while it increases going down a group (another shell is added).
  • Melting and boiling points peak at Group 4 while they are at their lowest at noble gases
  • Ionisation energy is the minimum energy required to remove an electron from an atom of an element in a gaseous state.  Peaks at noble gases, reaches lowest at Group 1.  The more electrons lost, the more energy required until it finishes the shell and starts on the one underneath, in this case the energy jumps significantly.  This trend can identify how many valence electrons an element has.  Ionisation energy decreases going down a period since it is more reactive, loses electrons more readily.
  • Valency goes from 1-4.  Anything with more than 4 valence electrons is subtracted form 8
  • Metal to non metal across a period, more metallic going down a group.
  • Reactivity decreases across a period, more reactive going down a group
  • Electronegativity is the ability to attract an electron when it forms compounds.  It is greatest near Group 7 and decreases going down a group.

Module 3: Water
Week 1:

  • Water is a raw material, a necessity for all living things, a solvent, a transport medium and a thermal regulator. 
  • Surface tension is the resistance of a liquid to increase its surface area.  Higher surface tension means it beads in a spherical drop rather than spread out.
  • Viscosity is the resistance of a liquid to flow through a tube, higher the viscosity, the less it flows e.g. honey
  • Specific heat capacity of water is 4.18 J/K/g, it is how much energy required to heat one gram of a substance by one kelvin. 
  • Water has a bent structure/ shape of around 100 degrees.= as does hydrogen sulfide.  Ammonia has a pyramidal shape and methane has a tetrahedral structure. 
  • Heteroatomic molecules like water have electron pairs distributed unevenly, forming polar covalent bonds.  These creates areas of charge called dipoles and thus these polar molecules have a net dipole.
  • Electronegativities of atoms determine whether they are polar bonds and the shape determines whether it has a net dipole.
  • Dipole-dipole forces are intermolecular forces present in polar covalent molecules like hydrogen sulfide while dispersion forces are apparent in non-polar molecules like methane.
  • The strongest of these bonds, the hydrogen bond occurs between hydrogen, oxygen, nitrogen or fluorine atom.  This means water which has hydrogen bonds has the highest boiling point.
  • Surface tension occurs when the surface molecules have unbalanced forces since there are no attractions above it, causing a inward force.  Water is strong (H- bonds) and hence has a high surface tension.
  • Viscosity depends on how long the molecule chains are, the longer the higher, and the strength of the intermolecular bonds, more stronger, the more viscous. 
  • The tetrahedral structure of atoms in ice leaves space in between which is why it expands on freezing.

Week 2:

  • The solvent dissolves the solute to make a solution, occurs when a component is in larger amount than the other. 
  • Ionic compounds are soluble in water, they break up into ions.  Polar water molecules are strongly attracted to the charged ions.  Hence the ions become hydrated, and water is a polar solvent, because it consists of polar molecules.
  • Water of crystallisation is when water becomes bound to a pure substances, known as hydrates
  • Most molecular substances are non-polar and do not form hydrogen bonds with water, and are hence insoluble or have poor solubility.  Sugars can form hydrogen bonds hence it dissolves.  However non-polar substances dissolve in non-polar substances such as iodine. 
  • Covalent lattices or large molecules do not dissolve since they are so big and have many strong covalent bonds.

Week 3:

  • A solid is produced (precipitate) by two soluble solutions.  This is called a precipitation reaction.
  • Net ionic equation does not show spectator ions, complete ionic equation shows all ions while neutral species equation shows normal balanced chemical equation.
  • A saturated solution is when a given amount of solvent in a temperature cannot dissolve any more solute, e.g. too much salt in water- can't dissolve.
  • Concentration is the amount of solute present in a solvent or solution.
  • Different measures of concentration include mass/100ml (w/v), volume/100ml, (v/v), %(w/v and % (v/v) are percentages of before.  Also %(w/w) and ppm, grams of solute per million grams of solution. 
  • Measuring cylinders, pipettes, burette or volumetric flasks can be used
  • Dilution can change the concentration of a substance c1v1= c2v2
  • Molarity is no. of moles/ vol. of solution
  • Specific heat capacity can be measured by q= mc change in temperature.
  • Exothermic reactions release energy and endothermic reactions absorb energy.



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The chemical Earth



The chemical Earth


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