Metals Preliminary

7/30/2019 Metals Preliminary

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Metals

1. Metals have been extracted and used for many thousands of years1.1Outline and examine some uses of different metals through history, including

contemporary uses, as uncombined metals or as alloys

The Copper Age was 3200BC to 2300BC. It is the period that archaeological records indicate that

copper was the first metal to be extracted from its ore. Copper was heated with charcoal and

globules of copper formed. Molten copper was used to make ornaments and domestic utensils.

The Bronze Age was 2300BC to 1200BC. It was later discovered that heating copper with tin

produces an alloy, bronze. Bronze was harder than copper and more easily melted to be molded due

to its low melting point. Bronze was used for tools and weapons.

The Iron Age was 1200BC to 1AD. Iron is more reactive than copper, so it need a higher temperature

to melt. Hematite was mixed with charcoal in primitive furnaces by blowing air and obtaining asufficiently high temperature. By 1000BC, iron had replaced bronze for tools and weapons because it

was harder and had hard tensile strength.

The Modern Age is 1Ad to present. There had been more extraction and uses of other metals such as

aluminium, chromium and metal alloys. Iron is the most widely used metal today. Many other

metals have come into common use due to the advancement in extraction technology.


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Metals1.2Describe the use of common alloys including steel, brass and solder and explain how

these relate to their properties

Alloys are mixtures of metals, which can be mixed in any proportion so they do not have aconstant composition or chemical formula

o Properties of alloys vary with the composition Alloys are giant lattice structures of metal ions surrounded by a sea of delocalized electrons

o Alloys are held by strong metallic bonds Alloys are difficult to recycle since the metals are very hard to separate The composition of some common alloys are given in the following table

Alloy Composition Properties Use(s)

Carbon Steel 99.8% Fe, 0.2% C Hard but easily

worked

Nails, Bridges, Scissors, Car

Bodies, Cables & chains

Brass 65% Cu, 35% Zn & small

amounts of other

elements eg. Pb, Sn & Al

lustrous gold

appearance, hard but

easily machined,

polishes well

Plumbing fittings, musical

instruments, decorations.

Solder 33% Sn, 67% Pb Low melting point

Adheres firmly to

other metals when

molten

Joining metals together

(plumbing and electrical)

Stainlesssteel

74% Fe, 18% Cr, 8% Ni Strong and resistscorrosion

Sinks, cutlery & Machinery

Bronze 85% copper, 15% tin Hard, resists

corrosion, attractive

appearance

Statues, medals, weapons


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Metals1.3Explain why energy input is necessary to extract a metal from its ore

Energy is needed to physically break up the ore and separate the mineral from the gangueo It is also needed to chemically decompose the mineral to extract the metalo The most active metals form the most stable bonds, thus they need the most energy

to decompose them

The easiest metal compounds to decompose (eg. Copper) can bedecomposed by roasting with sand

The next easiest (eg. Iron) can be decomposed by roasting with coke(carbon) in a furnace

The most difficult (eg. Na and Al) are decomposed by electrolysis of moltencompounds or aqueous solutions.

Only a few metals such as gold, copper and silver occur in nature. Most metals occur ascompounds in rocks

o Gangue is the unwanted waste of an ore after the mineral has been extracted Pollution can results from inappropriate disposal of this waste Rehabilitation is the process of returning mined areas to their original state

The steps in extracting metal are:o Mining the Oreo Separating the mineral from the gangue using physical processes such as froth

flotation, or chemical processes such as dissolving

o Extracting the metal by chemical processes such as roasting or electrolysis- Chemical reactions used to extract metals from its ore either absorb or release heat, which is a

form of energy (extraction reaction). Most metals need energy input for this reaction.

- Energy also needed to mine the ore and to purify/concentrate it, and maintain high temperatures

needed for the extraction reaction

1.4Identify why there are more metals available for people to use now than there were200 years ago

Al, Ti, Mg, and NA are examples of metals that were not available 200 years

- This is due to the lack of technology for separating such reactive metals- Development of technological methods of separation have enable compounds such as

Al2O3 to be separated, which cannot be decomposed by heat

Metals tend to be harder to extract the more reactive they are. This requires a large amount of

energy that could not by supply by charcoal fire or using a furnace. The discovery and implication or

electricity has allowed the process of electrolysis to be used to extract the more reactive metals

more readily leading to an increased range of metals used today.


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Metals

2. Metals differ in their reactivity with other chemicals and thisinfluences their uses

2.1Describe observable changes when metals react with dilute acid, water and oxygenReaction with Dilute Acid: Metal + acidmetal salt + hydrogen

Metals above hydrogen in the electrochemical series react with dilute mineral acids; thesemetals dissolve, produces bubbles of H2 gas and the colour of the solution changes

Metals with passivating oxide layers do not react Metals neutralise the acid All metals except silver, gold and platinum react with dilute HCl and H2SO4 form hydrogen

gas. However, Na, K Li, Ca and Ba react with the water in the acid. Tin and lead react slowly

unless acid is heated.

Acids are substances which produce hydrogen ions H+. This ion is the species which reactswith the metal.

Balanced formulae equation: Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) Net ionic equation shows actual ionic species that undergoes change in reaction:

Zn(s) + 2H+

(aq) Zn2+

(aq) + H2(g)

The chloride ions are spectator ions and are thus omitted. They are included in the complete

ionic equation

Water: Metal + watermetal hydroxide + hydrogen

Water is more energetically stable than dilute acids, thus less metals react with it incomparison to acid

The most reactive metals dissolve and produce hydrogen gas with cold water Li, Na, K, Ca, Bareact with water at room temperature

Mg, Al, Zn, Fe react with steam at higher temperatures whereas Sn, Pb, Cu, Ag, Au, Pt do notreact at all

When reactions occur with water, the products are hydrogen gas and metal hydroxide.With steam, the product is oxide, not hydroxide

Oxygen: Metal + oxygenmetal oxide

Fewer metals react with oxygen than above Many metals react, some violent and rapidly Most produce metal oxide, usually a white substance All metals except silver, platinum and gold react with oxygen to form oxides. Li, Na, K, Ca, Ba react rapidly at room temperature. Mg, al, Fe, Zn react slowly but vigorously

if heated. Sn, Pb, Cu react slowly and only if heated.

Metals which burn in air form white solids which have none of metals physical property. Al and Zn become coated with a layer of oxide which prevents further reaction. Others form a powdery surface layer of oxide which does not impede further reactions. Cu forms a black surface layer of copper oxide.


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Metals2.2Describe and justify the criteria used to place metals into an order of activity based on

their ease of reaction with oxygen, water and dilute acids

The more violent the reaction between a metal and oxygen, metal and water and metal and dilute

acids, the more reactive and thus the higher the position of the metal in the electrochemical series.

The degree of reaction with oxygen can be used to identify the most reactive metals (less reactive

metals do not react). Metals of moderate reactivity can be determined by their reactivity with water

(less reactive metals do not react). Metals that are relatively unreactive can be determined by their

reactivity with dilute acids (less reactive metals do not react).

The more reactive the metal, the more vigorously it reacts with water, dilute acid and oxygen

More reactive metals lose electrons more easily

From reaction with oxygen: {Na, K, Ca} > {Mg, Al, Fe, Zn} > {Sn, Pb, Cu} > {Au, Ag, Pt}

From reaction with water: {Na, K, Ca} > {Mg, Al, Fe, Zn} > {Sn, Pb, Cu, Au, Ag, Pt}

From reaction with dilute acid: {Mg, Al, Fe, Zn} > {Sn, Pb} > {Cu, Au, Ag, Pt}

Metal Activity Series

As we go down list, ease of losing electrons

decreases, ease of oxidation decreases and ease

of reduction increases. List order determined by

displacement reactions + reactions with oxygen,

water and acid

RULE: A more reactive solid metal will displace a

less reactive metal in a solution

e.g. Mg(s) + CuSO4(aq) (aq) + Cu(s)

Copper will appear around reaction site

(magnesium = reaction site). If copper put into

MgSO4, opposite will not happen, as copper is

less reactive

More active metals more recently discovered as their compounds are stable and the metals are

harder to extract (electrolysis- a relatively new discovery- used to extract theses metals)

Least active metals (e.g. gold) occur free in nature, and their compounds are unstable. Order of

metals discovered and extracted is around the same as least to most active in series.

More reactive metals have more stable salts, less reactive metals have less stable salts. Salts are a

combined form of the metal, in less reactive metals they wont exist for long breaks down to form

pure metal


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MetalsUses of Activity Series

o Can be used to make predictions e.g. calcium is high in list, so expected to be active

o Can be used to choose metals for particular purposes

o Reflects current and possible future developments in the use of metals

2.3Identify the reaction of metals with acids as requiring the transfer of electronsAcids: substances which in solution produce hydrogen ions. The hydrogen ion in acid reacts with the

metal (metal lose electrons to become positive, hydrogen gains electrons to become negative). Acid

molecules break up when they react with metals.

Metals form positive ions, thus they give away electrons. These electrons are attracted to hydrogen

protons from the decomposed acid molecules. Resulting is the metal combining with the non-

hydrogen element of the acid to form stable metal salt compounds while hydrogen atoms combine

in twos to form stable diatomic molecules

REDOX REACTIONS (reduction oxidation)

Redox reactions involve transfer of electrons from one species to another. Reduction & oxidation

occur simultaneously in complete chemical reactions, as there can be no overall gain/loss of

electrons.

Oxidation is losing electrons (becomes positive), reduction is gaining electrons (becomes negative)

O.I.L R.I.G

o Loss of hydrogen (oxidation), gain of hydrogen (reduction)o Gain of oxygen (oxidation), loss of oxygen (reduction)

Oxidation agent (oxidant) causes oxidation of another species

Reduction agent (reductant) causes reduction of another species

Metals are always reductants, oxygen/hydrogen is always oxidantDocumenting Reactions

1. Balanced formulae equations: show the actual chemical substances used in the reactionZn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

2. Complete ionic equations: shows all ions involved in the solutions used for the equationZn(s) + 2H

+(aq) + 2Cl

-(aq) Zn

2+(aq) + 2Cl

-(aq) + H2(g)

3. Spectator ions: ions that do not undergo any change during a reaction (Cl-in above e.g.)4. Net ionic equation: shows actual ionic species that undergoes change in reaction spectator

ions are hence omitted

Zn(s) + 2H+

(aq) Zn2+

(aq) + H2(g)

5. Half equations- Describe the oxidation and reduction process separately in terms of electrons lost or

gained

- Compose half equations from net ionic equation (omit spectator ions)Zn(s) Zn

++ 2e

-[Oxidation]


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Metals

- When electrons are on the right hand side, LOSS of electrons has occurred (oxidation)2H

++ 2e

- H2(g) [Reduction]

- When electrons are on the left hand side, GAIN of electrons has occurred (reduction)- To determine how many electrons are lost/gained, look at valency and how many atoms

there are (e.g. H valency of 1, but there are 2 atoms, so gains 2 electrons).

6. Name reductant and oxidantHydrogen is oxidant, Zn is reductant

2.4Outline examples of the selection of metals for different purposes based on theirreactivity, with a particular emphasis on current developments in the use of metals

As we go down the activity series, metal ions become easier to reduce to metal atoms (to turn ions

back to atoms, electrons must be added)

Therefore the further down the activity series the metal is, the easier it is to extract

Some situations where choice of metal is based on reactivity:

o Roof guttering for houses non-reactive but expensive (e.g. Al) or cheaper but eventuallycorroding (e.g. galvanised iron) metals?

o Water pipes: metal must not react with water (e.g. copper)iron sometimes used, as itscheaper, but eventually corrodes

o Electrical contacts for circuit boards in electronic equipment cheap copper that eventuallycorrodes or expensive, non- reactive gold?

Metal activity series can be used to choose metals for particular purposes based on reactivity

o Calcium (active) used in nuclear industry to change uranium tetraflouride to uraniumo Gold (unreactive) used in jewellery, microchips and electronics- resistant to corrosiono Copper used in wiring and electronics due to low corrosion (as its unreactive)

Many metals are being used in alloys for extra strength, especially reactive metals

2.5Outline the relationship between the relative activities of metals and their positions onthe Periodic Table

Metals generally decrease in reactivity going across from left to right of the periodic table. - - Metals

generally increase in reactivity going down the periodic table. Left-right rule is more dominant than

up-down rule (e.g. Na more reactive than Mg, even though they are in the same row, but different

columns). No trends/patterns for reactivity of transitional metals hence the rule is only GENERAL

Hence most reactive metals are located on the bottom left hand side, least reactive top right

(excluding noble gases)


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Metals

More reactive

Increase

Increase

Decrease

Decrease

2.6Identify the importance of first ionisation energy in determining the relative reactivityof metals

Maximum number of electrons in each energy level is 2n2

where n is the level number

Ionisation energy: energy required to remove an electron from element in the gaseous state.

Na(g) + energy Na+

(g) + e-

Equation describes first ionisation energy for sodium atom

Na+

(g) + energy Na2+

(g) + e-

Equation describes second ionisation energy for sodium atom

reactivity of metals increases as their ionisation energy decreases

Metals with high ionisation energies are less reactive, and vice versa

- Ionisation energy measured in kilojoules per mole

- The lower the ionisation energy, the easier it is to remove an electron

- Ionisation energies show that drive to noble gas configuration is basis for chemical bonding

First ionisation energy: energy needed to remove completely the first electron from an atom in the

gaseous state

- The higher the first ionisation energy, the less reactive (lower in activity series) the metal is

- 1st

ionisation energy increases across a period as electron shells go from empty to full

- 1st

ionisation energy decreases down a group as electrons go further away from nucleus

- Noble gases have highest ionisation energies as they have stable configurations

- 2nd ionisation energy (energy needed to remove a second electron) always greater than 1st, as

the electron is being removed from a positive ion, which has greater electrostatic attraction

- Ionisation energy will increase greatly when an electron is removed from a lower shell

- More energy needed as ionisation levels increase

More reactive- dominant

Reactivit

Ionisation Energy


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Metals elements with low ionisation energies readily form positive ions, hence these elements form

ionic compounds (e.g. Na, Ca, Al)

3.As metals and other elements were discovered, scientists recognisedthat patterns in their physical and chemical properties could be used

to organise the elements into a Periodic Table

3.1Identify an appropriate model that has been developed to describe atomic structureThe atomic model has changed drastically over time since its initial development with John Dalton

(17661844). Dalton made the atomic theory, which stated that:

- Matter is composed of indivisible atoms- All atoms of an element are the identical- Atoms of different elements differ in weight and chemical properties- Atoms are neither created nor destroyed- Atoms of different elements combine in simple, whole number to form compounds

In 1897, J.J. Thomson discovered the electron, and proposed the plum pudding model in

1904, which portrayed the atom as a sphere of positive electricity with electrons scattered within it.

Rutherford disproved Thomsons theory in 1911 through his experiment of firing alpha

particles at a sheet of gold. In doing so, he proved that an atom is mostly empty space with an

extremely small, yet massive core, known as the nucleus, with the electrons orbiting the nucleus.

In 1914, Niels Bohr introduced the concept of electron shells, as well as stability and

radiation caused by jumping electrons. In 1926, Erwin Schridinger replaced the theory of shells withclouds.

The current atomic theory is as follows:

Protons are positively charged (+ve) particles that are found in the central nucleus of anatom, and weigh 1 amu (atomic mass unit*)

Neutrons are neutral particles (no charge) and are also found in the nucleus, and weigh 1amu

Electrons are negatively charged (-ve) particles that are found outside the nucleus. Theyexist in energy levels of shells that surround the nucleus. The mass of an electron is

miniscule 0.00055 amu (1/2000)

* 1 amu = 1.661x10-27

- Bohr Atom Model: the currently accepted structure, shows a positive nucleus surrounded by

electrons shells of varying energy levels

Bohrs model is an appropriate model to describe the atomic structure. The nucleus is the

central part of the atoms which contains the protons and neutrons. It has a positive

equal charge equal to the number of protons. The electrons move through a relatively

large space outside the nucleus. The electrons are kept moving around the nucleus by


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Metalsattractive electro static forces between the positively charged nucleus and negatively

charged electrons.

About 2000 years ago, Democritus suggested the idea that all matter was made of tinyparticles

o This was dropped since Aristotle did not agree with the idea. In 1808, Dalton put forward the atomic theory, stating that:

o Elements are made of very small particles called atomso Atoms cannot be broken into anything simplero Each element has its own kind of atomso Atoms cannot be created or destroyed in a chemical reaction

Joseph Thomson showed that atoms were not solve spheres, but made of tinier particlesthemselves.

o He suggested the plum pudding model, where negative particles were scatteredthroughout the nucleus

Ernest Rutherford showed atoms were largely empty space with most of the massconcentrated in a central positively charged nucleus

o He states that each negative particle was an electron which orbited around thenucleus

Niels Bohr altered this, showing the electrons orbit in energy levelso This model of an atom is referred to as the Bohr model of an atom.

3.2Outline the history of the development of the Periodic Table including its origins, theoriginal data used to construct it and the predictions made after its construction

The first scientific discovery of an element was in 1649, phosphorus. In 1789,

Antoine Lavoisier published a table of 33 elements, which contained compounds as

well. He divided the table into non-metals and metals and started the classification of

elements. In 1829, Joltan Dobereiner arranged elements exhibiting similar

characteristics into triads. He also stated that the atomic weight in the middle

element of the triad was the average of the other two, publishing his own version of

the Periodic Table based on this.

In 1863, using 62 elements of known atomic mass, John Newlands observedthat the properties of elements varied periodically with their atomic masses. He then

published his own Periodic Table.

In 1870, Dmitri Mendeleev arranged the known elements into horizontal rows

in order of increasing atomic mass. Elements with similar properties were arranged

in vertical columns. As a result of his work, Mendeleev proposed his periodic table

law, which stated that the properties of elements are periodic functions of their

atomic masses. He made accurate predictions of elements that were not yet

discovered by leaving gaps in his table. An example of one element Mendeleevpredicted was Germanium. This is the initial basis of the modern Periodic Table.


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Metals

In 1910, Henry Mosely changed this by organizing elements in increasing

atomic number rather than mass. He stated that when the elements are listed in

order of increasing atomic number, similar chemical properties reoccur periodically.

- Antoine Lavoiser (French Chemist) 1775-1785: classified the known elements into metal and non-

metal based on physical and chemical properties

- Johann Dobereiner (German Chemist) 1829: drew attention to several groups of three elements

with very physical and chemical properties

- John Newlands (Englishman) 1864: law of octaves where elements were arranged in order of

increasing atomic weight and states that every eighth element starting from a given one will have

similar chemical and physical properties

Realised that there were still elements to be discovered, leaving spaces for future

discoveries

Identified many similarities but also required similarities which were non-existent

- Dmitri Mendeleev (Russian) and Lothar Meyer (German) 1869: had worked independently of each

other but had similarly arranged the elements in order of increasing atomic weight, and placed

elements having similar properties under one another (forerunner of the modern periodic table)

Periodic Law: properties of the elements vary periodically with their atomic weights

Made predictions which were very close to the actual characteristics of some elements

which had not been discovered at the time

- Henry Moseley (British) 1914: resolved the discrepancies within the Periodic Table, which was

arranged according to weight, by basing the arrangement on the atomic number

Modified Periodic Law: properties of elements vary periodically with their atomic numbers

- Eventually, it was recognised that properties were dependent upon the number of protons and

hence the number of electrons

In the 1800s, 30 naturally occurring chemical elements were known.

French chemist, Antoine Lavoiser classified the elements into two groups, metals and

non metals based on their physical properties.

In 1829, a German chemist, Dobereiner recognized the similarities of several groups of

three elements in which he called the triads.


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MetalsIn 1864, an Englishman, John Newlands, proposed the law of octaves where the

elements were ordered according to their atomic weight.

In 1869, Mendeleev proposed the periodic law where the properties of the elements vary

periodically with their atomic weight. He arranged the elements with increasing atomicweight and grouped them with elements with similar properties. Mendeleev knew that

there were still more elements to be discovered and left spaces in his periodic table.

In 1914, a British chemist, Henry Moseley, proposed a modified periodic law where the

properties of the elements vary periodically with their atomic numbers.

In 1787, Lavoisier classified the known elements into metals and non-metals, based on theirphysical and chemical properties

Dalton in 1808 put forward his atomic theory.o From this, he suggested that the atomic mass of hydrogen, should be given as 1

since it was the lightest element.

o These masses came to be known as relative atomic masses In 1829 Dobereiner, noticed some similarities in some sets of three elements, which he

called triads.

o Eg. He found Li, Na, and K to have similar properties In 1863, Newlands put forward his law of octaves.

o He pointed out that if elements were placed in order of increasing mass, there was arepetition every eighth element.

o Newlands work however broke down when some elements did not fit his ideas. In 1869, Mendeleev proposed the periodic law

o This states that if the elements were arranged in order of increasing atomic masses,then elements with similar physical and chemical properties (including valency)

occurred at regular intervals.

o His table was successful because he placed elements where he though they shouldgo and left gaps for elements he predicted would be discovered later.

o He could even predict the properties of these unknown elements, due to theirposition on his table.

In 1914, Henry Moseley stated the modern period law.o This states that when atoms are arranged in order of increasing atomic number

(number of protons) they show a repeating pattern of properties.

o Using atomic numbers rather than atomic mass solved some problems, such as Kand Ar, where K after Ar but has a lighter atomic mass.

The modern periodic table placed the elements in the vertical Groups and horizontal Periods,and includes the nobles gases discovered between 1893 and 1898

In 1789, Antoine Lavoisier published a table of 33 elements.o

Contained some compounds which at that time could not be broken

down into simpler substances


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Metals

o Divided his table into metals and non-metals In 1864, Meyer arranged that elements in order of increasing atomic mass and

placed them in groups based on their valency.

In 1869, Dmitri Mendeleev arranged the elements in horizontal rows in order ofincreasing atomic mass. Elements with similar properties were arranged in

vertical columns in the table. Mendeleev was also able to make accurate

predictions about elements that bad not been discovered by leaving gaps.

In 1913, Henry Moseley proposed the concept of atomic number to fix theirregularities in Mendeleevs table. It was eventually determined that atomic

number determines the chemical properties rather than atomic mass.

3.3Explain the relationship between the position of elements in the Periodic Table, and:- electrical conductivity

- ionisation energy

- atomic radius

- melting point

- boiling point

- combining power (valency)- electronegativity

- reactivity

Property Trends across a period Trends down a group

Atomic Radius Decreases Increases

Electrical Conductivity Decreases Groups 1-3 decreases

Groups 4-5 increases

Electronegativity Increases Decreases

Melting & Boiling Points Increases from Groups 1-4

then decreases

Groups 1-2 decreases

Transition metals increase

Groups 3-4 decrease

Groups 5-8 increase

Ionisation Energy Increases Decreases

- Reactivity: The most reactive metals are found in the lower left corner of thePeriodic Table. This can be deduced from the trends in the ionisation energy.

The most reactive non-metals are found in the top right corner of the Periodic

Table.


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Metals

- Valence: The most common valence of an element is its group number (if it isin Groups 1 to 4) or 8 minus its group number (if it is in Groups 5 to 7). The

transition metals usually have more than one valency, and thus, their

valencies cannot be deduced. Noble gases rarely form compounds, therefore

their common valence is zero.

- Electrical conductivity

Decreases as you go across a period (metal to non-metal)

- Ionisation Energy

Increases as you go across a period

Decreases as you go down a group

- Atomic Radius

Decreases as you go across a period (the nucleus pulls the valence electrons closer as the

positive charge is greater)

Increases as you go down a group (more shells are added)

- Melting and Boiling Point

Variable throughout the table

- Combining Power (Valency)

Increases then decreases across a period

The remains the same down a group

- Electronegativity (ability to gain an electron)

Increases as you go across a period (charge is higher so attraction is greater)

Decreases as you go down a group (distance is greater so it is harder to pull an electron in)

Chlorine is the most reactive non-metal

- Reactivity

Dependent upon electronegativity and ionisation

o Electrical conductivityAcross a period, the electrical conductivity of elements decreases because elements are

less metallic. Non metals do not have free mobile electrons in their crystal lattice. Down a


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Metalsgroup, the electrical conductivity of elements increases because they are more metallic.

Down a group, the valence shell is further away from the nucleus and can more easily

escape into the lattice.

oIonization energy

Ionization energy in the energy required to remove an electron from an atom of the

element in the gaseous state. Across a period, the ionization energy increases because

the atomic radius decreases across a period. The valence electrons closer to the nucleus

experience a stronger nuclear pull. Down a group, the ionization energy decreases

because the atomic radius is bigger and outer electrons are not as attracted to the

nucleus of atoms.

o Atomic radiusThe atomic radius is the average distance from the nucleus to the valence shell. Across a

period, the atomic radius decreases as the valence shells are closer to the nucleus. Down

a group, the atomic radius increases because the number of electron shells increases.

o Melting point and boiling pointAcross a period, the melting point increases from group I to group IV the decrease from

group IV to group VIII. The lattice changes from metallic bonding to covalent network

and then covalent molecular. Down a group, it decreases from groups I to IV and

increases from groups V to VIII

o Combining power (valency)The combining power of a group increases down the periodic table. Across the periodic

table, the combining power decreases.

o ElectronegativityElectronegativity is the tendency of an atom of an element to attract electrons. Across a

period, the Electronegativity increases as the metallic character decreases. Down a group,

the Electronegativity decreases as the metallic character increases.

o ReactivityThe reactivity of elements down a group increases and it decreases as it goes across a

period.

Across a period Down a group


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MetalsAtomic Radius Generally increases due to

shielding

Generally decreases due to increasing

protons and no shielding

Electronegativity Generally decreases due to greater

radius

Generally increases due to decreasing

radius and tendency to hold electrons

Ionisation Energy Generally decreases due to

electronegativity and radius

Generally increases due to greater

electronegativityMelting and Boiling

Points

Generally increase till group IV

before decreasing again

Observing non metal groups they

increase however decrease down

metallic groups

Electrical

Conductivity

Generally decreases from well to

negligible (metal non-metal)

Trend can be described in group 4 as

increasing down the group (non-metal

semi- metalmetal)

Valency Valency refers to number of electrons on outer shell and corresponds with its

group number (groups 1 -4) or 8 minus group number ( groups 5-8)

Reactivity Metals: generally increases down a group and decrease across period

Non-metals: trend less clear when ionic bonds form reactivity generally

decreases down the group

Atomic radiuso The radius of an atom decreases across a period due to the increasing nuclear

charge causing the valence shell to be attracted closer to the nucleus

o The radius of an atom increases down a group due to the extra shell being addedonto the atom whilst the number of electrons in the outer shell remain the same.

Boiling and melting pointo The BP and MP of an atom across a period relates to bonding:

From Group I-IV, BP and MP increases since bonding becomes stronger From Group V-VIII, BP and MP decreases since intermolecular forces

become weaker

o The BP and MP of an atom down a group varies, dependant on the bonding and typeof crystal structure

Going down Group I-IV, BP and MP decreases Going down Group V-VIII, BP and MP increases

Electrical conductivityo Conductivity generally decreases across a period as elements become less metallic(except transition metals where it increases again since they are very metallic)o Conductivity is variable

Conductivity of Group I-III decreases as you go down the group Conductivity in Group IV-VII increases as you go up, since the elements

become more metallic.

Electronegativityo It is the measure of the ability of a atom to attract electrons towards itselfo Across a period, electronegativity increases since atomic radii decreases, thus

electrons can approach closer to the nucleus to feel a stronger attraction force

o Down a group, electronegativity decreases, since atomic radii increases yet theelectrons in the valence shell remain the same, causing the electrons to have aweaker attraction force to the nucleus.


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Metals Valency

o Maximum valency relates to the number of valence electrons an element haso Maximum valency increases across a period, since the number valence electrons

increase across a period

o Maximum valency is constant down a group, since the number of valence electronsremain constant down a group.

4. For efficient resource use, industrial chemical reaction must usemeasured amounts of each reactant

4.1Define the mole as the number of atoms in exactly 12g of carbon 12(Avogadrosnumber)

A mole is defined as the amount of a substance that contains the same number of particles as there

are atoms in exactly 12g of carbon-12. Chemists have determined that the number of atoms in 12g

of carbon-12 is 6.022 x 1023. This number is called Avogadros number after the Italian scientist

Amadeo Avogadro. One mole of any substance contains 6.022 x 1023 particles of that substance.

A mole of a substance is that quantity which contains as many elementary

units (e.g. atoms, ions or molecules) as there are atoms in exactly 12 grams of the

carbon-12 isotope.

The Avogadro Constant (for which we use the symbol NA) is the number of

atoms in exactly 12 grams of the carbon-12 isotope.

NA = 6.02 x 1023

particles per mole

A mole of an element is the mass which in grams is numerically equal to the

atomic weight. A mole of a compound is the mass which in grams is equal to the

molecular weight.

Explanation: It was found that the mass of the carbon atom was 1.99x10-23

grams.

Expressed differently, there are 6.02x1023

atoms in 12 grams of carbon.

Therefore, is 6.02x1023

atoms of carbon have a mass of 12 grams, then

6.02x1023

atoms of titanium (atomic mass 48) will have a mass of 4x12=48 grams, as

each atom of titanium is four times heavier than carbon. For sulfur, 6.02x1023

atoms

(each of which is 3212=2.67 times the mass of a carbon atom) must have a mass of

2.67x12=32 grams. In general:


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If for any element we take the mass which in grams is numerically equal to the

atomic weight, then it contains 6.02x1023

atoms.

Therefore 63.6 grams of copper will contain 6.02x1023

atoms. 108 grams of

silver will contain 6.02x1023

atoms.

This concept also extends to compounds. The molecular weight of water is 18.

This means a water molecule is 1812=1.5 times heavier than a carbon atom. Hence

6.02x1023

molecules of water must have a mass of 1.5x12=18 grams. In general:

If for any compound we take the mass which, in grams, is numerically equal to its

molecular weight, then it contains 6.02x1023

molecules.

Therefore 46 grams of nitrogen dioxide, NO2, will contain 6.02x1023

molecules.

98 grams of sulfuric acid, H2SO4, contain 6.02x1023

molecules. 342 grams of sucrose,

C12H22O11, contain 6.02x1023

molecules.

We call this quantity 6.02x1023 atoms or molecules a mole. There are two ways of

looking at a mole:

Further reading:

Moles of Gaseous Elements

When we talk of a mole of oxygen (atomic weight 16), do we mean 16 grams

or 32 grams? (Oxygen gas is O2, hence a molecular weight of 32.) Sometimes, to

avoid ambiguity, we need to specify which elementary units (atoms, ions or

molecules) we are talking about.

A mole of oxygen atoms is 16 grams. A mole of oxygen molecules is 32 grams.

This mole of oxygen molecules contains 2 moles of oxygen atoms because each

molecule of oxygen contains two atoms. Similarly, a mole of chlorine gas (molecules)

is 2x35.5=71 grams, whereas a mole of chlorine atoms is 35.5 grams.

Converting between mass, moles and numbers of atoms or molecules

Mass to moles:Number of moles = mass mass of one mole

= mass molar mass

= mass atomic or molecular weight in grams

Moles to number of atoms or molecules:Number of atoms or molecules = number of moles x 6.02x10

23

Moles and chemical equations

4P(s) + 5O2(g) 2P2O5(s)

The above equation tells us that:1. Qualitatively, phosphorus reacts with oxygen to form diphosphorus pentoxide

A

A number of particles

(atoms, ions, molecules)

6.02x1023

A mass

(the atomic or molecular

weight in grams)


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2. Four atoms of phosphorus react with five molecules of oxygen to form to

molecules of diphosphorus pentoxide

3. Four moles of phosphorus react withfive moles of oxygen to form two moles of

diphosphorus pentoxide.

That is, the chemical equation can be read in terms ofmoles as well as in

terms of atoms and molecules. Therefore:

4x31.0 grams of phosphorus react with 5x32.0 grams of oxygen to form

2x(2x31.0 + 5x16) grams of diphosphorus pentoxide

Or, 124.0 grams of Phosphorus react with 160.0 grams of oxygen to form

284.0 grams of P2O5.

Therefore,

A Chemical equation tells us the ratios by mass in which substances react or are

formed in a reaction.

The mole is a chemical counting unito 1 mole is defined as the number of atoms in exactly 12 grams of the carbon-12

isotope

o This number of atoms is: 6.022 x 1023(Avogadros number)o Number of moles can be calculated by:

Where:o n is number of moles (mol)o m is mass of substance (g)o M is molar mass of substance

4.2Compare mass changes in samples of metals when they combine with oxygen

The total mass of a system remains constant during a chemical reaction. Masses are

transferred from molecules. Stoichiometry is the qualitative description of portions

by moles of the substance in a chemical reaction. The stoichiometry of a reaction is

the relative qualities as they are represented by the coefficient of the balanced

equation.

e.g. 2Mg(s) + O2(g) 2MgO(s)

During this magnesium and oxygen reaction, 48.62g of magnesium is present,

as well as 32g of oxygen. At the end of the reaction, two separate magnesium oxide

molecules will be formed, with a total mass of 2x(24.31+16)=80.62 grams. Thus, it

obeys the law of conservation of mass.


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- Oxygen has a mass of 16 amu and so a molar mass of 16 g

- For every oxygen atom that is added to the compound with a metal, the formula mass increases by

16 amu

Metals exhibit very varied reactivities in their reactions with oxygen. An example, lithium

and sodium and potassium tarnish rapidly when exposed to air and must therefore be

stored in liquid paraffin old. Also other metals react with oxygen and explode. Due to the

different reactivities of metals, the less reactive a metal is, then the more the metal

weighs. If a more reactive metal reacts with oxygen, then it can result in an explosion andit is less weight.

Law of conservation of mass: matter can neither be created nor destroyed in achemical reaction

2Mg(s) + O2(g) 2MgO(s)(using the periodic table to find the elements masses)o 224.305 + 2x15.999 2(24.305 + 15.999)o 80.608 80.608o LHS = RHS, therefore it obeys the law of conservation

The total mass of the system remains constant in a chemical reaction. When metals react

with oxygen in the air, they generally form metal oxides. The oxygen combines with the

metal and adds mass to the original metal. The metal is a limiting agent as oxygen is usually

in abundance relative to the metal sample.

4.3Describe the contribution of Gay-Lussac to the understanding of gaseous reactions andapply this to an understanding of the mole concept

Experiments measuring the combining volumes of gases led to an improvement in our

understanding of the formulas of substances and the development of the mole concept. The French

chemist Joseph Gay- Lussac found that there was a simple relationship between the volumes of

gases involved in chemical reactions. For example, in the reaction of hydrogen and oxygen gas to

produce water, the relationship is as follows:


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MetalsHydrogen gas oxygen gas steam

2H2 (g) + O2 (g) 2H2O (g)+

2L of + 1L of 2L of steam

hydrogen oxygen

Similar simple ratios are found to exist for other reactions involving gases. An example is the

reaction of hydrogen gas and chlorine gas to produce hydrogen chloride.

Hydrogen gas chlorine gas Hydrogen chloride gas

H2 (g) + Cl2 (g) 2HCl (g)

1 volume of + 1 volume of 2 volumes of hydrogen

hydrogen chlorine chloride

Gay- Lussacs law of combining volumes can be stated as follows:

the ration of the volumes of gases involved in a reaction, if measured at the same temperature and

pressure, are expressed by small, whole numbers.

After studying the volumes in which gases reacted, Gay-Lussac, in 1808, proposed

the law of combining volumes:When measured at constant temperature and pressure, the volumes of gases

taking part in a chemical reaction show simple whole number ratios to one another.

For example:

1. 100 mL of hydrogen reacts with 100 mL of chlorine to form 200 mL of hydrogen

chloride (onevolume reacts with onevolume to form twovolumes)

2. 100 mL of hydrogen reacts with 50 mL of oxygen to form 100 mL of steam

(gaseous water) at temperatures above 100C (twovolumes react with onevolume

to form twovolumes). Avogadro put forward the following hypothesis due toconfusion over diatomic molecules such as hydrogen and oxygen:


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Equal volumes of all gases, measured at the same temperature and pressure, contain

equal numbers of molecules.

e.g. 2H2(g) + O2(g) 2H2O (g)

This means that in 2L of hydrogen gas, there must be the same number of

molecules as there are in 2L of oxygen gas. It also means that the greeter the volume

of gas, the greater the number of molecules it will contain.

Simple number ratios occur due to the molecules of gases combining in simple

number ratios, as seen with the co-efficient in the equations. It is possible to extend

Avogadros hypothesis to a consideration of the number of moles of gases. Since

equal volumes of gases contain an equal number of particles, it can be deduced that

these equal volumes of gases also contain the same number of moles of gas.

Gay Lussac released the Law of combining gases which states that:o When two gaseous elements combine, the volumes of the gases that react are in

simple whole number ratio to each other and also to the volume of the product of it

is a gas. (All volumes measure under same conditions)

o It is a summary of experimental results Gay-Lussac found that the volume of gases (under the same conditions) was not conserved

in a chemical reaction such as one between hydrogen and oxygen gas. However he found

that the ratio of reactants and products appeared as simple whole number. Avogadro later

explained this by proposing that elements could form molecules. This led to Avogadros law

which stares Equal number of molecules occupies the same volume at the same

temperature and pressure.

4.4Recount Avogadros law and describe its importance in developing the mole concept

Avogadro summarized his ideas into what is now called Avogadros Law which states:o Under the same conditions of temperature and pressure, equal volumes of all gases

contain the same number of molecules

ie. 100ml of H2 has same number of moles as 100ml of O2o He proposed that the particles in gaseous elements were not single atoms as Dalton

believed by were made of more than one atom He coined these particles molecules

4.5Distinguish between empirical formulae and molecular formulae

An empirical formula shows the ratio in which atoms are present in acompound

A molecular formula shows how many of each type of atom are present


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Metals The empirical formula of a compound is the simplest whole number ratio of the

numbers of atoms of each element in the compound. The molecular formula

specifies the actual number of atom of each element in a molecule. E.g. the

compound, hydrogen peroxide has the molecular formula of . The molecule

contains two hydrogen atoms and two oxygen atoms bonded together. Theempirical formula of hydrogen peroxide would be HO.

5. The relative abundance and ease of extraction of metals influencestheir value and breadth of use in the community

5.1Define the terms mineral and ore with reference to economic and non-economicdeposits of natural resources

Minerals are naturally occurring inorganic substances, usually compounds with a

particular chemical composition and a definite crystal structure. Examples of minerals

include hematite, magnetite, gibbsite, boehmite, malachite and chalcopyrite.

Ores are naturally occurring deposits that are mixtures of minerals from which a

substance, usually a metal can be economically extracted. Examples of ores include

bauxite and iron ore.

Minerals are any useful naturally occurring elements (eg gold) or compounds (eg bauxite)from the earth

Ores are metal bearing substances (mixtures) from the earth with commercial valueMetal Ore Name Ore composition

Iron Haematite Fe2O3

Copper Cuprite Cu2O

Aluminium Bauxite Al2O3

Zinc Sphalerite ZnS

5.2Describe the relationship between the commercial prices of common metals, theiractual abundances and relative costs of production

The commercial price of metals depends on a few factors including their relative

abundances and the cost of production.


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MetalsThe greater the abundance of a metal the lower the commercial price of the metal would

be.

The cost of production of the metals depends on where it is located and the amount of

energy input. If the location of the ore is located in a high population zone, the miningprocedure would be difficult because there would be damages done to the environment

and increase the cost of production.

If an ore is located in remote places, then the cost of production would increase because

it would cost money to transport the raw materials to refinery plants.

The more reactive the metal is, then the higher the energy input is needed for extraction

and it would increase the cost of extraction.

5.3Explain why ores are non-renewable resources

Ores are non-renewable resources as they were formed when the earth wasformed and there is no way of forming more ores.

Also, the rates at which humans are using these ores are much faster thanthe rate that these ores are formed in the earth. Therefore our consumption

of these ores is much greater than their production, resulting in the

classification of ores as non-renewable resources.

Ores are deposits of naturally occurring minerals which were formed during the

evolution of the universe and the planets; therefore they are non renewable resources.

Abundance:o A mineral must be sufficiently concentrated in the ore body to make it

economically viable to extract. If the concentration is low, it is usually

not economical, as the costs of extraction and production are greater

than the value of the mineral obtained.

Cost of production:o Companies need to determine the cost of mining, milling and

extracting a metal from its ore. This would impact upon the final


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price; transportation costs may also vary due to the remoteness of

some ore bodies. The location can also affect the price as there may

be a lot of energy/water or the site may be environmentally protected.

Common metals:o Some metals are more expensive to extract than others due to higher

energy costs involved. For example, aluminium is more expensive to

extract than copper due to the high cost of the electrolytic process

used to make aluminium.

5.4Describe the separation processes, chemical reactions and energy considerationsinvolved in the extraction of copper from one of its ores

There are 4 types of copper oreso Malachite (CuCO3)o Chalcopyrite (CuFeS2)o Copper Sulfide (Cu2S)o Cuprite (Cu2O)

Stages of copper extractiono First, the ore is crushed into small particleso The crushed ore is then subjected to a physical process such as froth flotation (See

8.2 Chemical Earth - 1.9)

This produces a copper concentrate which is about 25-30% pure coppero Next the copper concentrate is heated with sand (SiO2). The sand combines with

iron oxide to form iron silicate, a liquid slag which is discarded

FeO + SiO2 FeSiO3o The resulting copper sulfide is then heated on its own while air bubbles through it,

reducing the copper sulfide to copper metal, and producing sulfur dioxide

Cu2S(l) + O2(g) 2Cu(l) + SO2(g) This remaining copper is blister copper, which is about 98% pure. This process is known as smelting.

Smelting is the extraction of metals by heating substances to hightemperatures to produce a molten material from which metal can

be obtained

o The blister copper is then subjected to electrolysis to purify it. The impure copper is used as an anode (positive since it is an electrolytic

cell) while the cathode is a thin sheet of pure copper in a sulfate solution

The anode copper gives up electrons to the circuit and goes into the solution Cu Cu2+ + 2e-

The cathode copper ions take electrons from the circuit, which deposit ascopper metal

Cu2+ + 2e- + Cu The cathode is then removed containing 99.95% pure copper

Throughout the process, there is a considerable input of energy, mainly in the form of heat.


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5.5Recount the steps taken to recycle aluminium

o Aluminium is collected from recycling binso The aluminium is transported to a recycling plant/facilityo Differ the aluminium from its alloys and aluminium metalo Aluminium products are deposited in a furnace where they are

subjected to very high temperatures. This process purifies the metal,

as it removes the impurities based on melting points.

o Analyse the purity of the aluminium and adjust it composition beforecasting it into ingots.

o The solution, pure aluminium, is remoulded into the desired productand sold to companies that need it.

Documents

Metals Preliminary