General Introduction Welcome to Dental Materials Science

LECTURE NOTES

Dental Material Science I

DENTAL TECHNOLOGY DIPLOMA General Introduction

Welcome to Dental Materials Science.

Please have the notes with you during lectures, when the material will be further explained. Although you may find taking some notes is useful to give an extra view on some points, these notes cover all the material you will need to pass the module. You will not need to take any notes during the lecture, which will go over the material in these notes again by a Powerpoint presentation, or on the board.

You may also find it useful to purchase the recommended materials science textbook, or to read the reference textbooks in the library. These books have been listed in the resources section which follows shortly.

These notes have also been designed to allow you to study and learn the material away from class. Being able to cover the material at your own speed and with your own pattern of learning is beneficial for many students. To help with this, there are questions typical of some of those asked in the examinations at the end of each section. These allow you to check your knowledge of each section as you proceed.

Module Overview

This module has been designed to build your knowledge of atoms, material structure, chemical bonding, and properties of materials. This knowledge leads to a better understanding of the chemical, physical and mechanical properties of materials. In later sections of the course this knowledge will be valuable in understanding the reasons for using a particular dental restorative material, and the techniques necessary to fabricate it.

i Recommended Textbook

Title: Dental Materials properties and Manipulation. Author: R.C. Craig, J.M.Powers, & J.C. Wataha. Publisher: Mosby ISBN: 0-323-02520-X

Reference Textbook

Title: Dental Materials: Properties and Selection. Author: O’Brien, William, J. Publisher: Quintessence Publishing Company Pub Place: Chicago Pub Date: 1989 ISBN: 0867151994

Your college may have a copy that you can borrow or you can purchase it yourself. iii

LEARNING OUTCOME 1

CLASSIFICATION OF MATTER.

Assessment criteria: You will have achieved this learning outcome when you can:  Distinguish the different states of matter.  Classify matter as elements, mixtures or compounds. TOPIC 1 - Classification of Matter.

Recommended Time - 2 hrs

Introduction In explaining to you the properties and use of various dental materials, we need first to understand what material, or matter, is. Scientists describe matter as belonging to different types in order to help understand its properties, For example, we often divide solid materials into metals, plastics (polymers) or ceramics, a classification we will learn more about later. Topic 1 introduces you to some of the ways scientists describe and classify matter. You will learn about:

 States of matter (ways in which matter can exist)

 Classification of matter as elements, mixtures or compounds.

Matter Matter is defined as anything which has mass and occupies a volume. Mass is the amount of material present. For example, when you see bubbles in a liquid, the bubbles have a volume, and the mass of the air can be determined. At the same time, you can observe differences between a gaseous and a liquid state of matter. This simple observation shows that the same matter can to exist in different states. Changing from one to another state of matter is a reversible, physical change. We can change water (a liquid form of hydrogen oxide) to ice (a solid form of the same compound) by cooling it sufficiently. We can change the ice back into water by heating it

States of Matter

Matter can exist in one of 4 forms: (i) Solid. (ii) Liquid. (iii) Gas. (iv) Plasma

Plasma is a rare state of existence for matter on this planet and needn’t bother us much for dental work, but it is important to understand the other three.

The differences in the behaviour of matter as solid, liquid and gas is caused by the behaviour of its atoms in that state. For example, water can exist as a solid (ice) below 00C, as a liquid above 0oC, and as a gas if heated above 1000C. What causes the properties of these different states of matter is the how mobile the atoms are each state. In the solid state, the atoms or molecules are fixed in position due to strong forces between molecules. Because the molecules or atoms remain in these fixed positions, the only movement possible is vibration. Every solid has a fixed volume and a fixed, definite shape. However, when the solid is heated, the atoms or molecules react to the extra energy by vibrating with increased frequency and amplitude. They are still held firmly in place, however, and cannot break free of the forces holding them in place until their energy becomes great enough. At this point, the matter becomes a liquid. We say that it has reached its melting point.

The forces between molecules are much weaker in liquids so the particles have greater mobility. Liquids are able to flow, a property due to the constant motion of their particles relative to one another. This is why they have no definite shape. The particles have only limited movement, however. They cannot move apart much, so that liquids have a constant, or fixed volume at any one temperature. Most particles of a liquid are held within the liquid due to forces of attraction between molecules (surface tension), but particles can gain enough energy to escape and form a vapour. This is called evaporation. As a liquid is heated, more and more particles evaporate until the temperature reaches the boiling point and a complete change of state from liquid to gas occurs.

In the gaseous state the particles are in constant motion and free to move in any direction. As a result, gases are not only fluid (like liquids, they have no fixed shape), they also have no fixed volume. As the particles in a gas regularly collide and rebound from each other at high speed, they move around until they fill the whole container. If they are not contained they will fly away due to their highly mobile state. As gases are heated, the heat energy is transferred to increased motion (velocity) of the particles, and can be observed as an increase in pressure.

At room temperature, if a substance appears as a liquid then it has a melting point below room temperature. If a substance is a gas at room temperature then it has a boiling point below room temperature.

Kinetic Theory of Matter

This theory explains how matter can be changed from one state to another. We know that all matter is be made up of tiny particles (atoms or molecules.) At any temperature above absolute zero, these particles are in a state of constant motion.

The amount of motion of a substance in any state of matter is due to its particles responding to the available energy. Heat is a form of energy. If a substance is liquid at room temperature, its particles are mobile, or in the liquid state. Its melting point is below room temperature. For water, with a melting point of 0 0C, there is enough energy at 200C to keep water particles mobile and able to flow. As the temperature of water is decreased the mobility of the molecules decrease until they cannot move fast enough to overcome the forces of attraction between them. The substance then changes state to become a solid. The molecules are no longer free to move, but they are still able to vibrate in positions fixed relative to each other.

If we decrease the temperature further, particle vibration decreases until -273 0C is reached. This is referred to absolute zero, since at this point all vibration stops.

The temperature at which we observe changes of state are exact for pure substances. If we change the purity of the substance, the melting point and boiling point will also change.

Homogenous and Heterogeneous Matter

We use these terms describe the composition and properties of matter. For example, if we have pure water in a container, then a sample taken from anywhere in the liquid will have the same composition, and properties. Because of this, only one value of a property is needed to fully describe pure matter. It is said to be homogenous, its properties are identical at any point in it.

If two elements or substances are mixed together, such as oil and water, they will quickly separate. We are able to easily separate the two distinct parts by physical means. This is called heterogeneous matter. The properties will not be the same at any point in the body. They will depend on which of the mixture components has been sampled.

Elements, Mixtures and Compounds

A pure substance is composed of only one type of atom or molecule, and it will have homogeneous properties. An atom is the smallest part of an element that still has the properties of that element

A pure substance composed only of one sort of atom is called an element. A pure substance made from one sort of molecule is called a compound.

A compound is a pure substance made by combining two or more elements in a fixed proportion by weight. A molecule is the smallest part of a compound that still has the properties of that compound

When a second element is introduced to a first, the chemical and physical properties will change. For example, pure water melts at 00C and boils at 1000C. If we dissolve salt in the water, its composition changes. It is no longer pure. If we re-measure its melting and boiling points, they will have changed also. The difference in composition and properties between pure water and salt water explains the difference between a pure substance and a mixture. The properties of a mixture are a mixture of the properties of the individual substances. As the amounts of each substance vary, so will the properties.

Mixtures do not have fixed melting points or boiling points; they change with composition. There is no chemical reaction involved in their formation so they are still easily separated. In the case of salt water, the water can be evaporated to form pure water leaving pure salt water behind.

If two or more pure elements are made to react together chemically, then the result will be the formation of a compound. The compound formed will be a pure substance and will be different in composition and properties to the original elements. An example of forming a compound is the combustion of pure hydrogen in the presence of oxygen to form pure water. The compound formed is in a different state of matter at room temperature, does not resemble its gaseous reactants, and cannot easily be reversed to hydrogen and oxygen. A chemical reaction is needed to do that.

Another example is mixing pure iron and sulphur together. Only when the mixture is heated will a chemical reaction occur to form a new compound, iron sulphide. This has different composition and properties to its reactants. ELEMENTS, MIXTURES & COMPOUNDS – Practical Assignment I

AIM To examine some of the properties of elements, mixtures and compounds.

METHOD (1) (i) Weigh out 5g of iron (Fe) and 3g of sulphur (S).

(ii) Place these two elements in a mortar and grind them into a fine powder with a pestle.

(iii) Spread this powder on a piece of filter paper and examine it with a magnifying glass. Is it uniform throughout (homogeneous) or is it non-uniform (heterogeneous)?

(iv) Pass a magnet under the filter paper and observe what happens. Is there any separation? Does this indicate a homogeneous or heterogeneous material?

(v) Fill a clean test tube to a depth of approximately 10 mm with the ground material and shake it up. Is there any separation? Half fill the test tube with water and shake. Is there any separation?

METHOD (2) (i) Place 10mm of the ground material into a Pyrex test tube. Heat it over a bunsen flame in a fume cupboard. Note any reactions.

(ii) Allow the mass in the test tube to cool. Remove it form the test tube and grind it up in the mortar and pestle. Place some of the ground powder on a filter paper and examine it with a magnifying glass. Compare the results to those obtained in 1(iii).

(iii) Test the powder with a magnet as in 1(iv) and compare the results.

(iv) Place some of the material obtained in part 2 in a test tube to a depth of 10 mm and repeat 1(v). Compare the results.

RESULTS It would be helpful to record all your observations of the results in a tabulated form.

CONCLUSIONS Consider the results you have written down. Do they indicate that we have an element, a mixture , or a compound at each stage of the experiment? What do the results show you about In light of your discussion draw some conclusions about elements, mixtures and compounds. Check Your Progress Self Evaluation Questions

Listed below are questions which will help you to review Topic 1,

Write your answer to each question on the lines below the question.

You can check your answers with the ones given at the end of this topic.

Q1. List three (3) different states of matter. (i) ______

(ii) ______

(iii) ______

Q2. Describe the motion of particle of matter in each state you listed in Q1. (i) ______

(ii) ______

(iii) ______

Q3. Briefly describe the Kinetic Theory of Matter. ______

______

Q4. Define the following terms:

(i) Homogeneous 

(ii) Heterogeneous 

(iii) Element 

(iv) Mixture 

(v) Compound  Suggested Responses for Topic 1

Q1. (i) solid (i) liquid

(ii) gas

Q2. (i) Vibration of particles only – rigidly restrained. (ii) Particles have constant motion and can flow. (iii) Particles are free to move in any direction.

Q3. All matter consist of particles which are in a constant state of motion. Changes in temperature increase or decrease motion and cause changes of state.

Q4. Define the following terms:

(i) Homogeneous  chemical properties are uniform  only one physical distinct property throughout the material. (ii) Heterogeneous  show two or more different property which allows separation easily by physical means (iii) Element  A pure substance composed of only one type of atom. (iv) Mixture  A substance containing with two or more elements or compounds combined in no fixed proportions. It shows the properties of each of its components.  The components are easily separated by physical means (heterogeneous) (v) Compound  This is a pure substance made by chemically combining atoms of two or more elements in fixed proportions by weight. A compound has its own chemical and physical properties, different from those of its components. It is still homogeneous. LEARNING OUTCOME 2

DISTINGUISH BETWEEN METALS AND NON-METALS, USING KNOWLEDGE OF ATOMIC STRUCTURE.

Assessment Criteria:

You will have achieved this learning outcome when you can:  Name and describe sub-atomic particles.  Describe the arrangement of sub-atomic particles within the atom.  Relate the atomic number of an atom to its structure.  Classify atoms as metals or non-metals based on their structure.  Name, write the symbol for and describe the atomic structure of the first twelve elements.  Name, and write the symbol(s) for the elements within the periodic table that are of interest to dental restoration study. TOPIC 2 - Atomic Structure.

Recommended Time - 3 hrs

Introduction We now understand that materials can be classified according to their different behaviour, both physical and chemical. To understand why materials are different from each other, we first must understand some basic chemistry about how matter is made, and how this influences its properties. This topic starts the process of understanding the basic structure of materials by introducing you to the structure of atoms, under the following headings:

 The type and nature of sub atomic particles.  The arrangement of sub atomic particles in an atom.  The relationship between atomic number and structure.  The classification of atoms, including the division into metals or non metals.  Important properties of the first twelve elements.  Identifying elements that are present in dental materials.

History of Atomic Theory (just read this, not examinable) Precisely what goes to make up matter or substances is a problem that has fascinated scientific philosophers for centuries.

Early philosophies considered matter to be made from of four elements: earth, fire, air and water. Around 400 BC Greek philosophers proposed that matter consisted of tiny indivisible particles called atoms.

In the early 1800’s John Dalton proposed a revolutionary new approach:  All elements are made up of atoms  Atoms cannot be created or destroyed (they are indivisible)  Atoms of different elements may combine with atoms of another element in definite ratios.  Atoms of one element are different from atoms of another element.

By 1820, laboratory experiment had found the presence of smaller particles in atoms, which suggested the presence of sub-atomic particles. Atoms could be taken apart, contrary to Dalton’s ideas.

In 1911, Ernest Rutherford confirmed the presence of sub-atomic particles, and made the following conclusions:  Each atoms has a nucleus which is positively charged.  Most of the atomic mass is contained in the nucleus.  The nucleus is surrounded by an almost empty space that makes up the rest of the atom.  Negatively charged electrons are present in this space around the nucleus. The negative charge on the electrons balances the positive charge of the nucleus.

In 1913, Niels Bohr suggested that the nucleus contains two different types of sub atomic particles. This gave rise to the modern atomic theory.

Modern Atomic Theory

Bohr conclude his atomic theory as:

 Atoms consist of subatomic particles

 The nucleus contains protons (+ charge) and neutrons (no charge).

 A cloud of electrons (- charge) orbits the nucleus.

 The volume of the nucleus is extremely small compared to the volume of an atom.

 The atom is electrically neutral since the number of electrons = number of protons.

Properties of Subatomic Particles

Subatomic Symbol Charge Mass Mass Location Particle (grams) a.m.u Electron e- -1 9.07x10-28 0.00055 outside nucleus Proton p+ +1 1.672x10-24 1.0073 inside nucleus Neutron n 0 1.672x10-24 1.0087 inside nucleus

All elements are made up of different combinations of these subatomic particles. The number of each type of sub-atomic particle in an atom can be determined from the information about that particular element contained in the Periodic Table. This is a table of all the known elements and their basic properties, arranged in order of atomic number. e

e n n n p n Electron n orbits n np p e n Nucleus

The arrangement of particles in an atom. PROTONS (+ charge), NEUTRONS (no charge) and ELECTRONS (- charge).

Atomic Number and Mass Number

What makes the difference between elements? To determine this we have to look at the arrangement of subatomic particles that make up each atom. This arrangement is different for each element. The information can be determined from the Atomic Number and the Mass Number:

Atomic number = number of protons in the nucleus

Mass number = number of protons plus neutrons in the nucleus

As each atom must be electrically neutral, the number of electrons must be equal to the number of protons, which is the atomic number. If you look at the table of elements, you will see, for instance;

Atomic Number Mass Number Carbon 6 12

From this. we can work out that that carbon has: 6 protons (atomic number) 6 neutrons (mass number minus atomic number) 6 electrons (number of electrons = number of protons) Element Atomic Number Mass Number (Z) (A) Hydrogen 1 1 Helium 2 4 Lithium 3 7 Beryllium 4 9 Boron 5 11 Carbon 6 12 Nitrogen 7 14 Oxygen 8 16

We know now that the central nucleus contains the protons and neutrons, but we don’t know how the electron orbits are arranged. For example, carbon has 6 protons (atomic number) and to be electrically neutral, must have 6 electrons. Logically there must be some consistent arrangement of these electrons, because they are moving around the nucleus without crashing into each other,

Electron Structure

Bohr described electrons as moving in fixed circular orbits around the nucleus, rather like planets orbiting around a sun. This is why his description is often referred to as the “Planetary Model” of an atom. These orbits, shells, or orbitals are different distances away from the nucleus, so each electron in a different shell must have a different energy. The electrons in the furthest orbit form the nucleus have the most energy, while those closer to the nucleus have less energy.

Bohr identified each electron shell with a number, n. The shell closest to the nucleus had n=1 which is the lowest energy level. The next shell, n=2, has a higher energy level and so on for n=3,4,5,6.

It seems logical that the outer shells, being larger, could hold more electrons. In fact, it turns out that the maximum number of electrons which can fit in each shell is governed by the formulae:

Maximum Number of electrons in any shell “n” = 2n2 n can be 1,2,3,4,5,6.

For example in the first electron shell, n=1.and n2 = 1. The maximum number of electrons which can be fitted into that shell is 2, because 2x(1)2=2.

For the second shell n=2. n2 = 4 The maximum number of electrons in this shell is 8, because 2x(2)2 = 8

For the third shell n = 3. n2 = 9. The maximum number of electrons in the third shell is 18 because 2x(3)2 = 18 Consider the case of carbon. As we discussed earlier, it has six electrons. Its electron configuration can thus be calculated as

2 electrons in the 1st shell (n=1) 4 electrons in the 2nd shell (n=2)

The second shell could hold a maximum of eight electrons, but carbon only has six. After putting two into the first shell there are only four left, so the second shell can only have the remaining four electrons in it.

Under normal conditions electrons in their shells are referred to as in their “ground state.” If atoms are heated, electrons gain energy and they may jump to higher energy levels. When dropping back to the ground state, they may re-emit the same amount of energy.

If we gave each shell a number to identify it, this could become confused with the number of electrons in the shell, so chemists have identified each shell with a letter instead The closest shell to the nucleus is called K. The next is L followed by M, N etc.

The following table shows the electron configuration for the first twelve elements. Remember, the maximum number of electrons is expressed by 2n 2.

Element Number of Maximum Electron (chemical symbol) electrons number in shell Configuration K , L , M , N Hydrogen (H) 1 2 1 Helium (He) 2 2 2 Lithium (Li) 3 2, 8 2, 1 Beryllium (Be) 4 2, 8 2, 2 Boron ( B) 5 2, 8 2, 3 Carbon (C) 6 2, 8 2, 4 Nitrogen (N) 7 2, 8 2, 5 Oxygen (O) 8 2, 8 2, 6 Fluorine (F) 9 2, 8 2, 7 Neon (Ne) 10 2, 8 2, 8 Sodium (Na) 11 2, 8, 18 2, 8, 1 Magnesium (Mg) 12 2, 8, 18 2, 8, 2

Names and symbols of elements As you can see from the table above, each different element has been given a name by its discoverer, and a symbol of one or two letters made up from the element’s name or from its Latin name. The elements are named after greek or roman gods, scientists, countries, or anything else which took their discoverer’s fancy at the time. The symbols are used as a type of short-hand to represent the elements in chemical formulae and equations. For the purposes of this course, you should know the names and symbols of the first twelve elements, and their symbols, and those of another fourteen common elements or those of interest in dental work. These are set out in the next table

Element Symbol Element Symbol Element Symbol Hydrogen H Gold Au Chlorine Cl Helium He Silver Ag Sulfur S Lithium Li Palladium Pd Phosphorus P Beryllium Be Platinum Pt Mercury Hg Boron B Iron Fe Zinc Zn Carbon C Cobalt Co Calcium Ca Nitrogen N Nickel Ni Oxygen O Chromium Cr Fluorine F Tin Sn Neon Ne Copper Cu Sodium Na Aluminium Al Magnesium Mg Lead Pb

Periodic Table

You will notice that the first twelve elements in the table have been listed in order of their Mass Number, which is the way that they were listed as the number of elements being discovered grew. The normal way of considering all the elements at once is the Periodic Table. Scientists in the 1800’s discovered that some elements had very similar chemical properties to each other, even though their atomic and mass numbers were different. To find out why this was so, they arranged groups of elements with similar properties in columns and rows and looked for patterns in their properties or the numbers. In this way they made early versions of a periodic classification.

The modern Periodic Table lists the elements in a series of boxes arranged in columns and rows.. In the horizontal row, elements increase in atomic number from left to right. Each box contains important information about the element:

Atomic number, or “Z” Mass number, or “A” Chemical symbol Electron configuration Originally, elements with similar chemical properties had their boxes arranged in a vertical column. This arrangement is still used. For example, Helium, Neon, Argon, Krypton and Xenon are all inert gases and are found in a column on the right hand side of the periodic table.

Look at the periodic table shown below. Although it is more than seventy years old, little has changed except for the addition of a few extra heavy elements with atomic weights above 105. Note that elements which have very obvious metallic properties are found on the left of the table. The less “typical” metals are found in the middle, the transition metals. Just before the non-metals on the right, separated by the stepped vertical line, there are elements which have some properties of both metals and non-metals, the “metalloids” such as arsenic, or silicon. Why should these obvious groupings occur? What determines the difference between metals and non-metals, so that arranging the elements in order of atomic number and similar chemical properties will reveal it?

It turns out that the chemical properties of any element are controlled mainly by the number of electrons in its outer shells. This is a most important fact.