Metals

*
Properties of metals
* Distinguish between metals and non-metals by * Deduce an order of reactivity from a given set
their general physical and chemical properties. of experimental results.

* Identify and interpret diagrams that represent Extraction of metals
the structure of an alloy. * Describe the use of carbon in the extraction of
* Explain why metals are often used in the form of some metals from their ores.
alloys. * Describe the essential reactions in the
Reactivity series extraction of iron in the blast furnace.

* Place in order of reactivity: potassium, sodium, * Relate the method of extraction of a metal
calcium, magnesium, zinc, iron, hydrogen and from its ore to its position in the reactivity
copper, by reference to the reactions, if any, of series.
the elements with Uses of metals
* water or steam, * Explain the use of aluminium in aircraft
* dilute hydrochloric acid (except for alkali manufacture in terms of the properties of the
metals). metal and alloys made from it.

* Compare the reactivity series to the tendency of * Explain the use of zinc for galvanising steel,
a metal to form its positive ion, illustrated by its and for sacrificial protection.
reaction, if any, with: * Explain the use of aluminium in food containers
* the aqueous ions of other listed metals, because of its resistance to corrosion.

* the oxides of the other listed metals.

* Group III-The NO-NAME Metals
They have 3 electrons in their outer
Group II-The Alkaline Earth Metals shell, most show increasing
They are similar to the Alkali metals but reactivity as you go down except for
are not as reactive. They also form alkaline Boron (which is a metalloid). Boron
oxides. is generally unreactive.

Group I-The Alkali
Metals
They react vigorously
with water – hydrogen
bubbles off, leaving
solutions of their
hydroxides which are
alkalis (pH of more
than 7).

The Transition Elements
They are much less reactive and do not
react with water or oxygen. They also show
no clear trends in reactivity unlike groups
I, II, and III. Most can form ions with
variable charges (Fe3+ and Fe2+).

*
Which of these properties belong to metals?

shiny low density
weak malleable

sonorous high melting and
ductile conductor boiling points
brittle
dull insulator
strong
low melting and boiling thin
high density
points

Remember metals are on the
LEFT side. Over 80% of the
elements are metals

Physical Chemical
* They are strong * They react with oxygen to form
* They are malleable oxides
* They are ductile * Metal oxides are bases
* They are sonorous (make a * Metals form positive ions when
ringing noise when you strike they react
them) * For the metals in groups I,II, and
* They are good conductors of III the charge on the ion is the
electricity and heat same as the group number but
* They have high melting and the transition elements have
boiling points variable valency
* They have high density

*
An alloy is a mixture of metals that has different (improved) properties
from the metal elements that make it. The pure metal is usually too soft to be
useful.
Often an alloy is composed mainly of one
metal (the parent metal) with small
amounts of other metals added.
The other metals replace positions of the
parent metal in the metal structure, or
sometimes fit into the spaces in the metal
structure.

The red balls are the parent metal. The blue and yellow balls are other metals.
The blue lines represent free electrons.

The other metals in the structure can change the properties of the alloy by
preventing the metal ions from sliding over each other. This can make the alloy
tougher and stronger than the parent metal.

Which one of these diagrams represents a pure metal?

Diagram b is a pure metal – the atoms are arranged in a regular lattice. When
pressure is applied (by hammering the metal) the layers can slide over each
other – this is why metals are malleable and ductile.

But when the metal is an alloy (diagrams a, c, and d) new atoms enter the
lattice and the layers can no longer slide easily. So the alloy is stronger than
the original metal.

Some examples of alloys:
(remember this is a small list- there are thousands of them

Alloy Made from Special properties Uses
70% Cu Harder than copper, Musical instruments, ornaments,
Brass
30% Zn does not corrode door knobs, etc…
90.25% Al
Aluminium 6% Zn Light but very strong,
Aircraft
alloy 2.5% Mg does not corrode
1.25% Cu
Stainless Cookware, cutlery, hardware,
steel (there are 80% Fe
Strong, does not surgical instruments, major
many types of 12% Cr
corrode appliances, industrial equipment,
stainless steel – 8% Ni
this is one) and construction

A very old video
from 1938 about
alloys- start at
3:40

* What does
reactive mean?

A reactive element has a strong drive to become a compound, so that its
atoms GAIN stable outer shells. So the metal reacts readily with other
elements and compounds.
Potassium (K) Please
Sodium (Na) Send
How did we determine which metal is Cats
Calcium (Ca)
more reactive than another? R Monkeys
Magnesium (Mg)
Well- by using experimental Aluminium (Al) E And
observations of course! (Carbon) A Coloured
Zinc (Zn) C Zebras
The elements that are bolded are Iron (Fe) T In
the ones that YOU will have to place Lead (Pb) I Large
in order of reactivity by reference to (Hydrogen) V Hired
the reactions with water, steam, or Copper (Cu) I Cages
dilute hydrochloric acid. Mercury (Hg) T Make
Silver (Ag) Y Sure
Gold (Au) Guarded
Platinum (Pt) Properly

* WATER
Some metals react chemically with cold water

metal + water → metal hydroxide + hydrogen

Metal hydroxide is Hydrogen gas is
formed liberated

Example: Potassium with cold water
YOU WRITE THE EQUATION

2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
This is typically the
apparatus that is used

This is a violent reaction! The
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g) hydrogen gas produced usually
explodes because of the heat
released!

Aqueous potassium hydroxide is
produced – it is an alkaline solution.

The test for an
alkaline solution:
Red litmus paper
will turn blue in
the solution

The test for
hydrogen gas:
A lighted splint goes
off with a “pop”
sound

Reaction of magnesium with cold water
Reaction is
much slower!

2Mg(s) + 2H2O(l) → 2MgOH(aq) + H2(g)

Reaction of zinc with cold water

but does it happen??

Zn(s) + H2O(l) → No reaction

Reactions of metals with cold water
Metal Observations/Equations
• Reacts very violently; explodes with cold water
• Enough heat is produced to ignite the hydrogen
Potassium gas produced
• Hydrogen gas burns in air
2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)
• Reacts violently
Sodium • Hydrogen gas formed may catch fire and explode
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Reactivity
of metals • Reacts readily
increase Calcium • Hydrogen gas formed
2Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
up the
series • Reacts very slowly with cold water
Magnesium
• A few bubbles of hydrogen gas produced only
Zinc
Iron
Lead • No reaction occurs
Copper
Silver

* STEAM
Some metals react chemically with hot water – aka STEAM

metal + steam → metal oxide + hydrogen
Metal oxide is formed
Hydrogen gas is
liberated

Metals like zinc and iron do
not react with cold water
but do react with steam

Example: Magnesium with steam

This is typically the Mg(s) + H2O(g) → MgO(s) + H2(g)
apparatus that is used

Reaction
Reaction of metals with steam is
Mg(s) + H2O(g) → MgO(s) + H2(g) vigorous!

• Reacts violently with steam
• Magnesium oxide (white power) and hydrogen
Magnesium produced
• Bright white glow produced
Mg(s) + H2O(g) → MgO(s) + H2(g)
• Burns in steam
• Zinc oxide (yellow when hot and white when cold)
Zinc
and hydrogen produced
Zn(s) + H2O(g) → ZnO(s) + H2(g)
• Reacts slowly with steam
• Iron(IV)oxide and hydrogen produced
Reactivity Iron • Reaction is easily reversible; so iron needs to be
of metals heated constantly for reaction to take place
increase 2Fe(s) + 3H2O(g) → Fe2O3(s) + 3H2(g)
up the Lead
series Copper • No reaction occurs
Silver

* dilute acid
Some metals react chemically with dilute hydrochloric acid (HCl)

metal + dilute acid → salt + hydrogen
Metal chloride is formed (salt) Hydrogen gas is liberated

You are given 3 unknown metals X, Y and Z and each is added to a test
tube filled with 1/3 dilute hydrochloric acid. What do you observe?

1. Arrange the reactivity of
the metals (starting with
the most reactive) with
reference to your
observations.
2. Using the following table
predict a suitable
identity for metals X, Y,
and Z

Reactions of metals with dilute hydrochloric acid
• Explosive reaction- should not be carried out in
Potassium school lab.
Sodium 2K(s) + 2HCl(aq) → 2KCl(aq) + H2(g)
2Na(s) + 2HCl(aq) → 2NaCl(aq) + H2(g)
• Reacts violently to give hydrogen and calcium
Calcium chloride
2Ca(s) + 2HCl(aq) → 2CaCl2(aq) + H2(g)
• Reacts rapidly to give hydrogen and magnesium
Reactivity Magnesium chloride
of metals 2Mg(s) + 2HCl(aq) → 2MgCl2(aq) + H2(g)
increase • Reacts moderately fast to give hydrogen and zinc
up the Zinc chloride
series 2Zn(s) + 2HCl(aq) → 2ZnCl2(aq) + H2(g)
• Reacts very slowly to give hydrogen and iron(II)
Iron chloride – and only if acid is concentrated
Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)
Lead
Copper • No reaction occurs
Silver

You are given 3 unknown metals X, Y and Z and each is added to a test
tube filled with 1/3 dilute hydrochloric acid. What do you observe?

1. Arrange the reactivity of the metals (starting with the most reactive) with
reference to your observations.
2. Predict a suitable identity for metals X, Y, and Z

How do you know
the gas produced
is hydrogen?

A lighted splint
goes “pop”

Magnesium Copper Iron

Summary of observations of the reactivity of the metals
from experimental observations
Remember the ones Reaction with
in bold you have to Reaction with dilute
Metal water or
know! acids
steam
Potassium Explosive reaction
Reactive React with
Sodium
metals Calcium
cold water Violent reaction

Fairly Magnesium Moderately fast reaction
React with
reactive Zinc
steam Slow reaction
metals Iron
Lead
No reaction
Unreactive Copper No reaction
with either
Silver
metals
Displacement of
hydrogen is a REDOX
• Hydrogen gas is produced in all three reactions reaction! H+ gained e-
• Reactive metals tend to react with non-metals to form compounds
• Unreactive metals have a greater tendency to remain uncombined

* carbon
• The reaction with metal oxides and carbon can be used to compare the
reactivity of metals. In other words when metals compete with each
other and with carbon to form a compound what will happen?
• The more reactive a metal is the more difficult it is to split its oxide
up.

We can compare these two reactions:
Magnesium oxide was not reduced
MgO(s) + C(s) → no reaction
2CuO(s) + C(s) → 2Cu(s) + CO2(g)
Copper oxide reduced to copper
Can be tested
using limewater Magnesium being a more reactive metal
than copper in the reactivity series did
not decompose when heated

K > Na > Ca > Mg > Zn > Fe > Pb > Cu > Ag

Why is this important?
• Metals below aluminum in the reactivity series can be reduced by
carbon
• Metals can be extracted from their ores by reduction with
carbon
• This includes metals from zinc to gold in the reactivity series
• Oxides of metals above zinc in the reactivity series Potassium (K)
cannot be reduced by carbon Sodium (Na)
• These metal oxides are very stable Calcium (Ca)
• They can be reduced by electrolysis Magnesium (Mg)
Aluminium (Al)
(Carbon)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
(Hydrogen)
Copper (Cu)
Mercury (Hg)
This is typically Silver (Ag)
the apparatus Gold (Au)
that is used Platinum (Pt)

*
• The reaction with metals, a metal oxide, and oxygen can also be used
to compare the reactivity of metals. In other words when metals
compete with each other and with oxygen to form a compound what
will happen?
• A metal will reduce the oxide of a less reactive metal.

Powdered iron is heated with copper(II)oxide. Can the iron grab the oxygen
from the copper(II)oxide?

This reaction gives out heat and
Fe(s) + CuO(s) → FeO(s) + Cu(s) starts to glow. The iron has “won”.

Iron removes the oxygen from the oxide of
the metal below it. It is a more reactive
metal than copper in the reactivity series.


*
Reactivity of a metal depends on the ease with which it forms metal
ions.

Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
Zinc “wins” against copper. It displaces the
copper from the copper(II)sulfate solution

All of these are REDOX reactions – we can also
call them SINGLE DISPLACEMENT reactions
because zinc and copper displaced each other.

A metal that is higher up in the reactivity
series has a higher tendency to form its
positive ion.


*
Remember the ones
in bold you have to
MOST REACTIVE

Potassium (K)
know!
Sodium (Na)
Calcium (Ca)
Magnesium (Mg) Metals above blue line: carbon
Aluminium (Al) can’t reduce their oxides
(Carbon)
Zinc (Zn)
Iron (Fe) Metals above red line: they displace
Lead (Pb) hydrogen from acids
LEAST REACTIVE

(Hydrogen)
Copper (Cu)
Mercury (Hg)
Silver (Ag) Non-metals carbon and
Gold (Au) hydrogen are included
Platinum (Pt) here for reference only!

*
1. The reactivity series is a list of the metals in order of their drive to
form positive ions (therefore a stable outer shell). The more easily

its atoms give up electrons, the more reactive the metal is.

2. A metal will react with a compound of a less reactive metal (an
oxide, an aqueous salt) by pushing the less reactive metal out of the
compound and taking it’s place.

3. The more reactive the metal, the more stable its compounds are.
They do not break down.

4. The more reactive the metal, the more difficult it is to extract from
its ores (they are stable). For the most reactive metals you need the
toughest method of extraction – electrolysis!

5. The less reactive the metal, the less it likes to form compounds.
This is why copper, silver, and gold are found as elements and not
potassium, sodium, or calcium. These elements are always found as
compounds.

* How do you extract a
metal from it’s ore?
* The most unreactive metals (Ag and
Au) occur in their ores as elements.
All you need to do is separate the Potassium (K)
metal from sand and other impurities Sodium (Na)
– it does NOT involve chemical Calcium (Ca)
reactions. Magnesium (Mg) electrolysis
Aluminium (Al)
* The ores of all other metals are (Carbon)
compounds - they have to be reduced Zinc (Zn) Heating with reducing
to get the metal. Iron (Fe) agent – ex Carbon

* The compounds of more reactive Lead (Pb)
metals are very stable so electrolysis Silver (Ag) Occur naturally as
Gold (Au)
is needed. elements – no
Platinum (Pt) chemical reaction
* The compounds of less reactive needed
metals are less stable so a reducing
Which method is
agent can be used. more expensive?

* The blast furnace!

The charge contains
three things: A mixture called the charge, containing
Iron iron ore, is added through the top of the
ore, limestone, and furnace.
coke
Hot air is added through the bottom.
Stage 1: The coke burns, giving off heat
Iron forms and
trickles down Stage 2: Carbon monoxide is made
Carbon monoxide
forms and rises Stage 3: The iron(III) oxide is reduced

Carbon dioxide After a series of reactions, liquid iron
forms and rises collects at the bottom of the furnace.

*
Metal Used for… Properties that make it suitable

Overhead electricity A good conductor of electricity (not as good
cables (with a steel as copper, but cheaper and much lighter);
core for strength) ductile, resists corrosion
Aluminium
Cooking foil and food Non-toxic, resistant to corrosion, can be
cartons rolled into thin sheets

Drink cans Light, non-toxic, resistant to corrosion

Can be deposited as a thin film, shiny surface
Coating CDs and DVDs
reflects a laser beam

Protecting steel from Offers sacrificial protection to the iron in
Zinc rusting steel

Coating or galvanising Resists corrosion, but offers sacrificial
iron and steel protection if coating cracks

For torch (flashlight) Gives a current when connected to a carbon
batteries pole, packed into a paste of electrolyte

*
Galvanization is the process of applying a protective
zinc coating to steel or iron, in order to prevent
corrosion (rusting).
Sacrificial zinc anodes are used to protect steel tanks, rails and ship hulls from
rusting. Since zinc oxidizes more easily than iron (it is higher in the reactivity
series), it corrodes first. When the anode is mostly consumed, it can simply be
replaced.
Any time you have two different metals that are physically or electrically
connected and immersed in seawater, they become a battery. The electrons that
make up that current are supplied by one of the metals giving up bits of itself - in
the form of metal ions - to the seawater. This is called galvanic corrosion and,
left unchecked, it quickly destroys underwater metals.

The way to counteract galvanic corrosion is
to add a third metal into the circuit, one
that is quicker than the other two to give up
its electrons. This piece of metal is called a
sacrificial anode, and most often it is zinc.

*
Aluminum is one of the most widely used metals in making planes. Why?
-high strength/weight ratio
-corrosion-resisting qualities
-comparative ease of fabrication
-nonmagnetic
-an excellent conductor of electricity

Aluminium-magnesium alloys are both
lighter than other aluminium alloys
and much less flammable than alloys
that contain a very high percentage
of magnesium.

Check this blog out about Aluminium alloys
in planes: http://blog.aluminium-
messe.com/aluminium-remains-the-number-
one-material-in-aircraft-construction/182/

Metals

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