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