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Demonstrate understanding of aspects of selected elements
Chemistry A.S. 1.4
2013
Periodic Table• We will concentrate on the following elements K,
Na, Li, Mg, Ca, Al, Cu, Fe, Zn, Pb, Ag, Au, C, N, O, S, Cl, Br, I
• Elements can be classified as Metals or non-metals.
• Complete the exercise “structure of metals”
Metals• Metals are found towards the left and bottom of the PT.• 1,2 or 3 valence electrons which are lost to form cations.• Metals form solids with metallic bonds.
• The physical and chemical properties of metals are explained by the metallic solid structure.
Physical properties of metals
• Electrical conductivity
• Thermal (heat) conductivity
• Density (gml-1)
• Lustre
• Malleability
• Colour
• State
• Ductility
• MP and BPThe metals you need to know about and be able to relate their chemical and physical properties to their uses are…
Na K Li Ca Mg Al
Zn Fe Pb Cu Ag
.
high
high
high >3 – except Na and Li
high – when freshly cut
high– more or less
silver – except Cu and Au
solid – except Hg
High
Vary but BP usually >1000oC
Physical Properties
Hardness – How easily a material can be cut or scratched with a knife.
Many pure metals, like Iron, are too soft for engineering so mixtures are formed which increase their hardness – these are called alloys.
“Alloys”
Alloys
• 2 or more molten metals are mixed and cooled to form a new mixture which has more useful properties than the parent metals.
• When atoms of a different size are added the regular lattice structure is distorted so it is harder for layers of atoms to move, therefore the alloy is harder.
Common AlloysName Use Parent
metalProperties
Brass Musical instrument
Cu, Zn Won’t tarnish, conducts like Cu
Solder Bonding metals Pb, Sn Melts at lower temp than Pb and Sn
Pewter Drinking vessels Pb, Cu Won’t poison water, stronger than Pb
Duralumin Aircraft bodies Al, Mg Low density, stronger than Al
Bronze Statues, marine fittings (cannons)
Cu, Sn Corrosion resistant, harder than Cu
Cupronickel “gold” coins Cu, Ni Harder than Cu, lasts longer
Steel
• Steel is an alloy made of metal and non-metal. (Fe and C)
• It has high tensile strength. (can be bent and returned to its original shape without breaking)
• Better corrosion resistance than Fe.
• Other metals can be added to make specialist steels. Eg. Cr and Ni = stainless
Uses of Metals
metal use property
aluminium
overhead power cablesgood electrical conductor, low density
drinks cans does not react with water
cooking pots good heat conductor
As Duralumin aircraft partshigh strength, low density
copperwater pipes does not react with water
electrical wires good electrical conductor
Iron as steelconstruction high strength
“Tin” cans high strength
“Wordsearch” / pg82 ESA
Chemical Reactions of Metals
• When metal atoms undergo chemical reactions new substances are formed.
• All metals have low numbers of electrons in their valence shells so they
have similar chemical properties.
• Metal + Acid Metal Salt + Hydrogen– Hydrochloric acid forms CHLORIDE salts
– Sulfuric acid forms SULFATE salts
• Metal + Oxygen Metal Oxide
• Metal + Water Hydrogen + Metal Hydroxide or Oxide
• The ease of removal of electrons determines how reactive the metal is.
Reactivity Series• Sodium• Lithium• Calcium• Magnesium• Aluminium• Zinc• Iron• Lead• Hydrogen• Copper• Silver• Gold In
cre
ase
s
Chemical reactivity generally involves atoms gaining or losing electrons. Why?
IONIC COMPOUNDSMetals lose valence electrons to form cations (+)
Non metals gain valence electrons to form anions (-)
The ions that have been formed are now attracted to each other.
So Mg2+ will be attracted to Cl-. This forms an ionic compound.
Naming rules:
- The positive ion is first, and the negative second.
- The negative ion ends in –ide, e.g. The sulfur atom becomes the sulfide ion.
Exceptions:
NO3- (nitrate) SO4
2- (sulfate)
CO32- (carbonate) HCO3
- (hydrogen carbonate)
Ionic Bonding
• Na + Cl
2,8,1 2,8,7
Strong electrical attractions (+/-) between oppositely charged ions in a 3D structure.
IONIC FORMULAESo Mg2+ will be attracted to Cl-.
Because Mg is 2+ and Cl is only 1-, Mg will attract 2 Cl’s.
The compound formed will be MgCl2. The subscript shows that the are 2 Cl’s for each Mg.
If the starting ions were Cu2+ and S2-, the 2 ions have the same charge. So each Cu will only attract 1 S.
The compound formed will be CuS. There is never any charges on the final product - they balance out
DIFFICULT ONESWhat is the formula for Magnesium Nitrate?
Find the 2 ions on your table of ions…
Mg2+ and NO3- This means that there are 3
oxygens attached to the Nitrogen – don’t let them get lost!So for each Mg we will need
2 NO3’s:
MgNO3
NO3
The shorthand way of writing this is:
Mg(NO3)2
The brackets are needed to show that we want 2 of the whole thing – you always need them if you have 2 of an ion with more than 1 bit. (polyatomic ions)
NH4+ and S2-
Fe3+ and OH-
Al3+ and SO42-
= (NH4)2S
= Fe(OH)3
= Al2(SO4)3
Metal reactions• Metal + Oxygen → Metal oxide
• Aluminium + Oxygen → Aluminium oxide
• 4Al + 3O2 → 2Al2O3
• Metal + Water → Metal oxide + Hydrogen
• Aluminium + Water → Aluminium oxide + Hydrogen
• 2Al + 3H2O(g) → Al2O3 + 3H2
• Metal + Acid → Metal salt + Hydrogen
• Magnesium + Hydrochloric acid → Magnesium Chloride +
Hydrogen
• Mg + 2HCl → MgCl2 + H2
General Equation
Word Equation
Formula Equation
Now write them as formulas and balance them!
Complete the following reactions:1) Lithium + water
2) Lithium + hydrochloric acid
3) Silver + oxygen
4) Magnesium + sulphuric acid
5) Copper + oxygen
6) Aluminium + oxygen
7) Zinc + water (g)
8) Sodium + sulphuric acid
9) Lithium + oxygen
10)Aluminium + hydrochloric acid
Lithium hydroxide + hydrogen
Lithium chloride + hydrogen
Silver oxide
Magnesium sulphate + hydrogen
Copper oxide
Aluminium oxide
Zinc oxide + hydrogen
Sodium sulphate + hydrogen
Lithium oxide
Aluminium chloride + hydrogen
METAL REACTIONS1. METAL + OXYGEN
Metals react in air to give metal oxides. Heating increases the rate of this reaction. (it may burn or change colour)
Metal oxides are Basic, but only the first 2 groups of the periodic table are alkalis (bases that dissolve in water).
An example:
Magnesium is reacted in the air (with O2) to produce a white powder, which turns litmus paper blue. Write the word and balanced symbol equation for the reaction.
Magnesium + Oxygen Magnesium oxide
Mg + O2 MgO
Metal + Oxygen Metal OxideGeneral eqn:
22
Ionic bond
Sodium + Water H2
Na + H2O H2
Metal + Water H2 + a HydroxideGeneral eqn:
+ Sodium Hydroxide
+ NaOH2
2. METAL + WATER
Some metals react in water to give Hydrogen gas (H2) and a Metal Hydroxide or oxide.
Reaction speed depends on the reactivity of the metal. Reactive metals react with cold water, others need steam .
An example:
Sodium reacts violently when placed in cold water and the gas produced sometimes explodes, but the reaction of Magnesium is only visible with steam and produces an oxide.
2 2
Group 1 metal reactions
3. METAL + ACID
Many metals react in acid to give Hydrogen gas (H2) and a metal salt.Reaction speed depends on the reactivity of the metal.
These ones react with acids.
An example:Magnesium fizzes when placed in a test tube with Hydrochloric acid. The gas produced explodes with a squeaky pop.
Magnesium + Hydrochloric acid H2
Mg + HCl H2
Metal + Acid H2 + a SaltGeneral eqn:
+ Magnesium Chloride
+ MgCl22
“Equations practice1”
ALUMINIUMAluminium is high on the reactivity series, but never seems to do anything. Why?
Aluminium forms an oxide coating very quickly.
Aluminum oxide is shiny and silver so it looks like the metal but it doesn’t react.
That is why aluminium is used for many things even though it is reactive.
This is worth remembering. Examiners love to ask about it.
Non-Metals
• Top and right of the PT• 4,5,6 and 7 valence electrons • Electrons are usually gained to form anions• Non-metals form ionic compounds with metals and
covalent bonds with other non-metals.
Physical properties of non-metals
The non-metals you need to know about and be able to relate their chemical and physical properties to their uses are…
C N2 O2 O3 S Cl2 Br2 I2
• Non-metals are poor conductors – except for graphite
• Colour• State
Physical propertiesElement Symbol State Appearance Hardness
Carbon C solid Grey solid/ colourless crystal
Soft/ very hard
Nitrogen N2 gas Colourless
Oxygen O2/O3 gas ColourlessO3 strong smell
Sulfur S solid Yellow, faint odour Brittle crystal
Chlorine Cl2 gas Green/yellow strong smell
Bromine Br2 liquid Red/brown strong smell
Iodine I2 solid Purple crystal sublimes
If 2 or more forms of the same element exist in the same state, but with different arrangements of atoms they are called allotropes.
“Diamond Vs Graphite”
Allotropes of Carbon1) Diamond – very hard, doesn’t conduct electricity, very high melting point , used for jewellery, cutting tools
2) Graphite – soft, shiny, does conduct electricity, very high melting point, used for lubricant, pencils, electrodes
3) Fullerenes – high tensile, conducts electricity, high ductility, dissolves in oil used for nanotechnology (molecular sponge)
Graphene• A single layer of graphite is
called graphene.• It has the ability to conduct
electricity.• It can be rolled to form
nanotubes.• It’s thinness and conducting
ability make it a “super material”
• Used for medicine delivery, carbon sequestration, wires in smaller electrical circuits.
Allotropes of Oxygen
• An electric spark can convert small amounts of 3O2 2O3
• O3 in the lower atmosphere is a respiratory pollutant.
• In the upper atmosphere it blocks uv radiation
• It can be used to purify air and water by killing microbes
Ozone hole
Sulfur• Yellow brittle crystal found
near volcanoes
Reactions• Burns with a blue flame
• S + O2 SO2
Uses• “Sulfuric acid production”-
contact process• Manufacture of fertiliser,
medicine, fabric, explosives
Contact Process
1) S + O2 SO2 (sulfur is burnt)
2) 2 SO2 + O2 2SO3 (lowers activation energy)
3) SO3 + H2SO4 H2S2O7 (dissolved to make oleum)
4) H2S2O7 + H2O 2 H2SO4 (gives off heat)
4000C, vanadium pentoxide
Read and take notes on pg 96,97 ESA study guide, properties of Sulfuric acid
Sulfur DioxideEnvironmental effects
– Acid Rain – Respiratory problems(choking, sharp smell)
•S(s) + O2(g) SO2(g)
•White smokey fumes, burns with a blue flame
•Dissolves a bit in water
•SO2 + H2O H2SO3
Sulfurous acid
Uses
•Bleach (oxidant)
•Preservative of fruit
•winemaking
Chlorine
• Found as NaCl and combined with other metals.
• Cl2 is poisonous yellow/green gas
Use• Kills microbes in water
Cl2 + H2O HCl + HOCl
HOCl H+ + OCl- (hypochlorite ion oxidises/ kills bacteria)
• Household bleach 2NaOH + Cl2 NaOCl + NaCl + H2O (hypochlorite ion sterilises)
• Monomer component of PVC• Bleaching agent in the paper industry
Makes water acidic
Bromine
• Found combined with other metals eg NaBr
• Br2 is poisonous red/brown liquid
Use• Kills microbes in water• Medications• Light sensitive agent in photography
Iodine
• Blue / Black solid at room temperature
• Sublimes to form violet/ pink gas
• Very rare but found as a soluble iodide salt in seawater and many seaweeds.
Use• Medicines, dyes, catalysts,
essential in diet
Trends down group 17
• Atoms in the same group have the same number of valence electrons and therefore react similarly.
• Atoms higher in the group have less “shells” so the valence electrons are closer to the positively charged protons in the nucleus.
• The smaller atoms therefore have greater attraction for electrons and are more reactive.
Nitrogen
• Colourless, odourless atmospheric gas• Sodium nitrate (saltpetre) found in Earth’s crust
Reactions
• N2 + O2 2NO
• 2NO + O2 2NO2
Ammonia, NH3 is a strong smelling gas which is water soluble.
• NH3 + H2O NH4+ + OH-
• Ammonia is a base and is therefore neutralised to make ammonium compounds.
• NH3 + HCl NH4Cl
Makes water basic
AmmoniaUses
•Making HNO3 (making explosives)
•Liquid refrigerant•Agricultural supplement•Industrial “ammonia production”•Lab production
•Ca(OH)2 + 2NH4Cl 2NH3 + 2H2O + CaCl2