IB DESIGN TECHNOLOGY
MICROSTRUCTURES and
MACROSTRUCTURES
MICROSTRUCTURES
Structure of
Matter
Look
around you and count how many different materials you can see. Add to your list
other materials you can think of, such as oil, bone, helium, alcohol. The list
could, given time, be almost endless, yet all the materials known to us are
composed of just 92 or so elements. Centuries ago it was thought that there
were just four basic elements, earth, water, air and fire. Modern science, by which
we mean discoveries made from about the 17th century, easily dismissed such
primitive ideas by showing for example that air was composed of at least two
elements (oxygen and nitrogen). However, many steps over many years were
necessary before it was shown that there are just about 92 elements.
Substances
or material may be said to be made from matter, all matter is composed of particles.
Most
substances are in the form of compounds, which means that they can be
broken down into other substances. An element is a substance, which cannot be
broken down into anything else. For example, water is a compound because it can
be broken down into oxygen and hydrogen. But oxygen and hydrogen are both
elements because they are impossible to break down. Many of the elements are
metals, but not all metals are elements. For example, copper and tin are both
elements but brass is a mixture, called an alloy, of copper and zinc.
Suppose
you could cut a metal element into smaller and smaller pieces, cutting each
small piece into even smaller pieces. Eventually you would be down to the
smallest possible pieces, which could be identified with the metal used: single
atoms. Cut the piece any smaller and the atom would be broken into electrons,
protons and neutrons, which could have been produced, from any element.
Atoms are the smallest particles of an
element, which can be identified as being from that element. All atoms have the same basic
structure.
Atoms
join together to form molecules which are the 'building blocks' of all
substances. For example, each molecule of water is composed of two hydrogen
atoms and one oxygen atom. Another example is oxygen gas, which has molecules
each, composed of two oxygen atoms. Imagine some of the possible combinations
of just a few types of atom; with 92 or so different types of atoms known, the
range of possible combinations is huge. Hence the enormous variety of
materials.
Molecules are the smallest
particles of an element or compound, which can exist independently.
For example, oxygen atoms at room
temperature do not exist independently. They join in pairs to form oxygen
molecules. So the symbol O2 represents a molecule of oxygen which is composed of two oxygen atoms.
Each element has its own symbol; each compound has the symbols of the elements
into which it can be broken down, usually written to show how many of each type
of atom there are in each molecule. So the symbol for carbon dioxide is written
as CO2, indicating that each
molecule has one carbon atom and two oxygen atoms joined together.
The nuclear model
of the atom
The idea that atoms could not be
divided or changed was a key starting point for scientists in the 19th century.
They supposed that atoms were indestructible, capable only of being joined
together to form molecules. In 1897 J.J. Thomson discovered that matter
contains tiny negatively charged particles which we now call electrons. He
showed that electrons from different elements were identical and so concluded
that all atoms contained electrons. Experiments by Ernest Rutherford over the
years 1908 to 1912, showed that every atom contains a point-like positively
charged nucleus where most of the mass of the atom is concentrated. We know now
that the nucleus contains two types of particles, protons and neutrons, and that the electrons
are outside the nucleus. The electrostatic force between each electron and the
nucleus prevents each electron from leaving the atom.
The charge of the proton is equal
and opposite to that of the electron. Neutrons are uncharged. So an uncharged
atom has the same number of electrons around its nucleus as there are protons
in its nucleus. Atoms become charged by adding or removing electrons usually.
Charged atoms are called ions.
The
mass of the proton is approximately the same as the mass of the neutron. The
mass of an electron is much less at about 1/2000 of the proton mass.
The
Periodic Table lists the elements in order of increasing atomic mass, so it
starts with hydrogen, then helium, then lithium and so on. The order number of
each element is called its atomic number Z; so for hydrogen, Z = 1. For helium
Z = 2, for lithium Z = 3, etc. Rutherford proved that the atomic number is
equal to the number of protons in the nucleus.
Each
hydrogen atom has only one proton in its nucleus. Each helium atom has two
protons in its nucleus. Yet the mass of a helium atom is about four times that
of a hydrogen atom. Why? Electrons are too light to account for the difference.
Neutrons are responsible. A helium atom with a mass about four times that of a
hydrogen atom must have two neutrons in its nucleus as well as two protons.
We use the atomic number Z and the
mass number A to identify each type of atom. So the symbol X identifies an atom
of element X which has Z protons and (A - Z) neutrons. An uncharged atom would
have Z electrons round its nucleus. The electrons of each atom are arranged in shells around the nucleus, each
shell able to hold up to a certain number of electrons. Each shell is an
electron energy level, and the electrons normally occupy the innermost shells
available since these are the lowest energy levels. The maximum number of
electrons, which each shell can hold, is worked out from the Periodic Table.
The innermost shell, which is nearest, the nucleus can take up to two
electrons; the next shell out can take up to eight electrons, etc. So an
uncharged sodium atom (Z=11) which has eleven electrons would have two
electrons in the innermost shell, eight electrons in the next shell, then a
single electron in the otherwise empty third shell. Sodium is very reactive because
that single electron can easily be removed.
States of
matter
Ice
is a solid. Its molecules are locked together in a rigid structure. When ice is
heated sufficiently, it changes to water. The energy supplied by the heating
proces enables the molecules to break away from each other, so its rigid
structure falls apart. When water is heated sufficiently, its temperature rises
to boiling point when the water changes to steam. Once again the heating
process supplies energy which enables the water molecules to break from each
other. Ice, water and steam are examples of the three states of matter.
The solid state Solids have fixed volume
and fixed shape.
The liquid state Liquids have fixed
volume but take the shape of their container.
The gaseous state Gases have no fixed
volume and no fixed shape.
To
change a solid to a liquid or to change a liquid to a gas, energy must be
supplied to break the bonds, which hold the molecules together. When a gas
changes to a liquid or a liquid changes to a solid, energy is released because
bonds are formed to hold the molecules together. For example, consider what
happens when salt is heated. To melt salt, it must be heated very strongly; at
room temperature, the atoms in salt are locked together in a very rigid
structure with strong forces holding the atoms in place. To enable the atoms to
break the grip of these strong forces, the salt must be strongly heated; this
makes the atoms vibrate so much that they break free.
The
forces, which hold atoms and molecules together, are due to the charged
particles in each atom. In other words, the forces are electrostatic in origin.
Electrons are mostly responsible, but there are several ways in which they
cause bonding forces.
Some IB definitions:-
Atom The smallest
part of an element that can exist chemically.
Molecule Two or more atoms which
are normally bonded together covalently.
Ion A positively or negatively charged
atom or molecule caused by the
loss
or gain of electrons from an atom or atoms.
Element A substance that cannot be
decomposed into simpler substances.
Compound A substance formed by the combination of elements in fixed
proportions. They may be bonded ionically or covalently.
Pure
substance A
substance made of only one element or compound.
Mixture A
substance made of two or more substances that can be separated by physical
means, ie not chemically bonded together.
Alloy A
mixture that contains at least one metal. This can be a mixture of metals or a
mixture of metals and non-metals.
Composite A mixture composed of two or more substances (materials)
with one substance acting as the matrix or glue.
A
graph showing the equilibrium position of a particle in a bond using a general
potential energy vs separation curve.
Bonding
A
bond is
a force of attraction between particles.
Ionic bonding
Crystals
of common salt (i.e. sodium chloride) are bonded in this way. An uncharged
sodium atom has a single electron in its third shell, the inner shells being
full. Each uncharged chlorine atom has seven electrons in its third shell;
again the inner shells are full. Now the third shell is full with eight
electrons, and since atoms prefer full shells, then a chlorine atom likes to
gain an extra electron whereas each sodium atom likes to lose an electron. So when
sodium and chlorine atoms form a sodium chloride crystal, each sodium atom
gives up an electron to a chlorine atom. The sodium atoms become positive ions
and the chlorine atoms become negative ions. The electrical forces between the
ions causes them to become regularly arranged, as shown in figure below. The
force between adjacent oppositely charged ions is called an ionic bond.
Put
an ionic crystal in water and it will dissolve. The effect of the water is to
weaken the electrical forces between the ions. The ions break off and the
crystal dissolves to form a solution. Most inorganic crystals are ionic.
Covalent
bonding
When atoms are unable to
gain electrons to complete part-filled shells, they can share electrons. Shared
electrons act as bonds between the atoms; this is referred to as a covalent
bond. For example, oxygen molecules are each composed of two oxygen atoms
joined by a covalent bond. Each uncharged oxygen atom has 8 electrons (Z = 8),
arranged with two in the innermost shell and six in the second shell. If an
oxygen atom shares two of its outer shell electrons with another oxygen atom,
which also contributes two electrons for sharing, each of the two oxygen atoms
has a full outer shell. Each covalent bond requires one electron from each atom.
So the two oxygen atoms form two covalent bonds since each atom contributes two
electrons for sharing.
Molecules
of organic compounds are held together by covalent bonds between their atoms.
Organic compounds contain carbon atoms, and an uncharged carbon atom (Z = 6)
has two electrons in its innermost shell and four in its second shell. The
second shell can take up to eight electrons, so to fill it a carbon atom forms
four covalent bonds with other atoms. For example, methane gas molecules each
have a carbon atom joined to four hydrogen atoms; each hydrogen atom forms a
covalent bond with the carbon atom to satisfy the full shell requirement of the
hydrogen atom. So by forming four covalent bonds, the carbon atom fills its
second shell. (See methane molecule)
Metallic
bonding
In
a metal the atoms have lost their outermost electrons which move freely inside
the metal. The metal atoms therefore become positive ions; they are held in
place in regular order by the electrical forces between the ions and the free
electrons. The atoms are arranged in an order, which can differ from one metal
to another. So when a metal solidifies, lots of tiny crystals, called grains,
are formed inside the metal.
All
metals conduct electricity. The reason is that they all contain free electrons.
When a potential difference is applied across a metal, the free electrons
inside the metal move towards the positive terminal. So an electric current in
a metal is due to the movement of free electrons.
Hence a metal is an ordered array of positively charged ions through which the free electrons move in all directions at high speed, as shown. The binding forces that hold a metallic crystal together are the forces between the positive ions and the cloud of moving electrons.
The
movement of the free electrons means that metals are good thermal and
electrical conductors. The metallic bonds are not fixed in position and can
therefore allow metals to deform without the bond breaking. This ductility
given by the metallic bond allows useful forming of metals to take place.
In an ionic bond the opposing charges of
the ions hold the crystal (eg NaCl) together in a lattice. The ions can often
be separated easily in water but the electrons stay attached to their
respective ions inside the crystal.
In a covalent bond the outer electrons of
some atoms can come close enough to overlap and be shared between the nuclei,
thereby forming a covalent bond. Each pair of electrons shared is called a
covalent bond. Mention of sigma, pi, double or triple bonds is not required.
Metallic bonding involves outer electrons
but these are freer and they can flow through the crystalline structure. The
bonding is caused by attraction between the positively charged metal atom
nuclei and the negatively charged cloud of free electrons, and is spread
throughout the lattice—"Positively charged nuclei in a sea of
electrons". Specific arrangements of metal atoms in crystals are not
required.
Iconic,
covalent and metallic are called primary bonds and their relative
strengths
are ionic > metallic > covalent.
Secondary bonds
In addition to the three types of
bonding already mentioned, there are also weaker secondary bonds, which exist in
substances such as water and many ceramics and plastics. The water molecules
have a positive charge near the hydrogen atoms and a concentration of electrons
causing a negative charge at the other side of the molecule. This distribution
of charge causes attraction between the water molecules. The primary covalent
bond between the oxygen and hydrogen atoms in each molecule is very strong but
this secondary hydrogen bond between the molecules can be broken by heating which
results in the water vaporising.
Another type of secondary bond is
the van der Waals bond. This is formed when the fluctuating electrostatic charge in
adjacent atoms of different molecules produces a weak electrostatic force
between the molecules. Van der Waals bonding is also often present as a
secondary bond between the long-chain molecules of polymers. Although the molecules
within each polymer chain have strong primary bonds it is the van der Waals
bonds that bind the chains together. When the polymer is stretched the van der
Waals bonds break easily, allowing significant deformation of the material.
They can then easily form again between new neighbouring atoms when the
material is released.
Solid
structures
The atoms or ions in a solid may be
thought of as hard spheres. If we examine the manner in which they interlock to
form the final solid, we can appreciate how this structure determines the
different characteristics of ceramics, plastics and metals. A structure which
has a regular arrangement of atoms repeated in all directions is referred to as
crystalline. This structure is present in all metals. Where the atoms are arranged
randomly, as in glass, the substance is said to be amorphous. The atomic structure of a
material mainly determines its microstructure and chemical and physical
properties.
Diamond and sand (SiO2)
molecules form in a network covalent (giant) structure.
In diamond each carbon is covalently
bonded to four other carbon atoms, tetrahedrally arranged. The carbons at the
edges are attached to hydrogen atoms. In sand (silica, SiO2) the
arrangement is also tetrahedral. Both sand and diamond are very hard.
|
C
C C C
C
C |
O
O
O
|
A crystal
is a
regular arrangement of particles (atoms, ions or molecules).
A crystal
is an homogeneous portion of matter that has a definite, orderly atomic
structure, and an outward form bounded by smooth, plane surfaces, symmetrically
arranged. Crystals are produced whenever a solid is formed gradually from a
fluid, whether the formation results from the freezing of a liquid, the
deposition of dissolved matter, or the direct condensation of a gas into solid
form. The angles between corresponding faces of any two crystals of the same
substance, regardless of size or superficial difference of form, are always
identical.
Most solid matter displays orderly atomic arrangement
and is of crystalline structure. Solids that have no crystalline structure,
such as glass, are called amorphous. In structure they show greater similarity
to liquids than to solids, and are known as supercooled liquids.
Amorphous solids
In
the simplest solids, the atoms have no regular pattern but are arranged
randomly in a given space.
Where the arrangement of the atoms
extends only to each atom's nearest neighbours we have an example of a
short-range order material. In a molecule of silica (SiO~), for example, the
silicon atom is covalently bonded to four oxygen atoms in a tetrahedral
structure similar to that of a diamond molecule and the molecules of silica are
then randomly bonded together. Other examples of amorphous solids are glass and
pitch. These amorphous solids are generally unstable and glassy in appearance.
The manufacture of amorphous ceramics (ceramic glasses) and polymers is easy and
inexpensive, and provides us with many useful transparent materials. Recent
developments have been made in producing amorphous metals.
Amorphous materials do not have
regular structures or crystal patterns. Short range order may occur as far as
next neighbour atoms. The general appearance of amorphous materials is glossy
and they can occur in ceramics, polymers and metals.
Fibre structure A fibre structure is used to describe the elongation
of the crystals in a cold worked metal, or any type of filament material from
which yarns and fabrics are manufactured by spinning, weaving, knitting,
bonding, or any filament material/molecule in food products.
Fibres have a
length-to-thickness ratio of at least 80 but are generally much longer. Textile
fibres and food are made up of polymers.
Melting When a pure solid
substance is heated sufficiently it will melt into a liquid. The atoms or molecules of a solid are
tightly bonded together and the solid consequently has a definite shape and
size. In a liquid the bonding
forces are not so strong and therefore the atoms or molecules can move,
changing the shape of the liquid to fit the container although its volume
remains fairly constant.
Boiling When a pure liquid
substance is heated sufficiently it will boil or vaporise into a gas. The atoms or molecules of a gas gain
sufficient extra energy during the vaporization process to overcome the atomic
bonding forces, thereby enabling them to move freely in the available
space. Thus the gas has no fixed
shape or volume.
Pure substances melt at a fixed temperature
but mixtures soften over a range of temperatures before melting. Therefore the working or deformation of materials in a plastic condition is possible
because of the range of temperatures over which a mixture softens.
Plastic
state – when a material is pliable and may be shaped by moulding. The plasticity of a material is
often increased by the application of heat.
Plastic deformation - The permanent deformation of a solid subjected to a
stress.
IB
properties/bonding matrix
The properties specified in the IB properties/materials matrix (see below) can
be organized into a properties/bonding matrix.
This shows the relative values of
the materials in the IB properties/bonding matrix in terms of their bonding
characteristics. The matrix allows
the designer to identify or evaluate the properties of materials in a given
design context.
IB properties/bonding matrix
|
|
Ionic |
Simple |
Network |
Metallic |
|
Chemical
properties |
|
|
|
|
|
Solubility
in water |
high |
low |
v. low |
v. low |
|
Solubility
in organic solvents |
v. low |
v. high |
v. low |
v. low |
|
Physical
properties |
|
|
|
|
|
Electrical
resistivity |
v. high |
high |
v. high |
v. low |
|
Thermal
conductivity |
v. low |
low |
v. low |
v. high |
|
Thermal
expansion |
low |
low |
v. low |
v. high |
|
Hardness |
high |
v. low |
v. high |
low–high |
|
Mechanical
properties |
|
|
|
|
|
Tensile
strength |
v. low |
low |
v. high |
high |
|
Stiffness |
high |
low |
v. high |
high |
|
Toughness |
low |
medium–high |
low |
v. high |
|
Materials ˆ |
ceramic |
plastics |
ceramic |
metals |