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The Physics and Chemistry of Matter

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David

Forums Super Moderator
Joined
Feb 20, 2001
Introduction

The purpose of this thread is to explain some of the Physics and Chemistry behind matter, which may prove potentially useful in cooling applications. I have seen several threads regarding various topics touching on Physics and Chemistry such as a thread where it was suggested some miriacle reaction could absorb all the heat from a CPU for months; threads on why water is used as a coolant and not something else; threads on what else we could make waterblocks and heatsinks from.

While not being an expert in the cooling field, I feel at the very least the material here may provide some useful info for those just looking for a basic primer on the science of matter, and may resolve some "why is water the best coolant" type threads.

This was mostly done from memory, so there may be slight mistakes and typos - feel free to post any corrections/additions and I will fix/add them

Atoms, Ions and Molecules
  • Atoms and Ions: descriptions and structures
  • Molecules: description and structures
    [*] Electron Arrangement: arrangement within an atom.
    [*] Magnetism: why and how?

The States of Matter
  • Solids: structure and properties
  • Liquids: structure and properties
  • Gases: structure and properties
  • Supercritical fluids

Inter- and Intra-Molecular Forces
  • A general note on inter/intra - molecular forces.
  • Intra molecular - Ionic Covalent, pi and sigma bonds.
  • Inter molecular - London forces, dipoles, Ions and Hydrogen bonding.
  • Metallic bonding

Heat Capacities and Latent Heats
  • Heat Capacity: definition and uses.
  • Latent heats: What and Why?

Thermodynamics, Entropy and Enthalpy.
  • The Laws of Thermodynamics
  • Entropy
  • Enthalpy
  • Gibbs Free Energy and reaction feasbility
 
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Atoms, Ions and Molecules

Atoms, Ions and Molecules

Atoms and Ions, description and structure

Put simply, every chemical in the world can be broken down into atoms. These are units from which molecules, complexes, structures and matter are made from.

Atoms consist of a positively charged nucleus and negatively charged electrons. The electrons orbit the nucleus in set defined orbits. These orbits can be calculated using the Schrodinger equation. In an atom, the magnitude of the charge on the nucleus is the same as the magnitude of the sum of the charges on the electrons.

The nucleus is made up of two types of particle. These are the proton and the neutron. The proton is postively charged. In an atom there are always equal numbers of protons and electrons. Protons and Electrons have equal charge. Neutrons have almost the same mass as Protons but are not charged. These lead to various isotopes of elements.

The mass of an atom is expressed in terms of Atomic Mass Units (AMU) where 1 AMU ~= 1.66 x 10^(-27) kilograms. The mass of an element can also be expressed as grams per mole where one mole is equal to 6.02 x 10^(23) atoms. Thus 1g/mole implies that 1 atom has a mass of 1 AMU, 18g/mole implies 1 atom has a mass of 18 AMU etc.

Notes:
(1) Neutrons and Protons have a mass of ~1AMU. However electrons are approximated to a mass of 0, as their mass is around 500 times less than 1 AMU so for simple calculations we can ignore them.

(2) The number of atoms in a mole is known as Avagadros number, written as L in older textbooks or N-subscript-a in more modern ones.


Example:
Helium exists in two stable forms. One form is Helium-4 and one is Helium-3. Helium 4 has an atomic mass of 4 AMU. As it is Helium it has 2 protons and 2 electrons. However this gives a mass of 2AMU. Thus it must have 2 neutrons as well. Helium 3 also has 2 protons, so it must have 1 neutron to give it its mass of 3 AMU.

Atoms with the same atomic number (number of protons) but different mass are known as Isotopes. From above - He-4 and He-3 are known as isotopes of Helium.

It is the number of protons in the nucleus which determines what a element an atom is. For any given element, the number of neutrons and electrons in its atoms may vary.


Atoms are extremely small, usually around 1 Angstrom in diameter. 1 Angstrom is 10^(-10) - ten to the power of minus ten metres. To give some sort of size comparison, the parts in a processor are approimately 1000 times larger (of the order of 10^(-7) metres)!
Most of this is actually empty space, the nucleus itself is around about 1/100,000 th of the size of the atom - of the order of 10^(-15) metres.


Ions are charged atoms. They either have more electrons than protons (negative ion, such as Chloride(1-) which has 17 protons and 18 electrons) or more protons than electrons (postive ions, such as Sodium(1+) which has 11 protons and 10 electrons).

The charge on an ion is simply the net charge - if there are more electrons than protons, the charge number is the number of extra electrons. If there are more protons than electrons this is the number of extra protons.

Example: Aluminium (3+) (or Al3+) has 13 protons and 10 electrons. Thus it is postive, by a magnitude of 3.


Molecules, description and structures

Molecules are the discrete packages* of an element or compound and can exist in many forums, depending on the compound.

Not entirely always true but this is used to simplify things

  • Discrete monoatomic molecules: In this case a molecule is the same as an atom. Elements like Argon, Kyrpton, Helium, Neon and Xenon exist in this fashion.
  • Discrete diatomic molecules: These types of molecules include Hydrogen, Nitrogen, Fluorine, Chlorine, Oxygen, Bromine and Iodine which exist as two atoms bonded together (H2, N2, F2, Cl2, O2, Br2, I2 respectively). Also included are many covalent (explained later) compounds like HCl (Hydrochloric acid).
  • Discrete polyatomic molecules: These are multi-atom molecules, which exist as seperate discrete molecules. Examples include H2O (water) and NI3 (Nitrogen Triiodide).
  • Ionic lattices: These are structures like sold Sodium Chloride. In this case you have a lattice of ions arranged in a Crystaline structure.
  • Dimers: Dimers are formed when two molecules attach together. Some carboxylic acids (organic acids) do this.
  • Polymers: These structures include many proteins and plastics. A string of monomers join together in a chain to form a long chain molecule. For example, ethene can polymerise (loads of ethene molecules join together in a big chain) to form poly-ethene (polythene).


The Arrangement of Electrons Within an Atom

The electrons within an atom are arranged within certain areas of electron density. We cannot pinpoint the location of a specific electron at a specific time however using mathematics and quantum theory we can predict the areas in which the electron is most likely to be found. These are termed orbitals.

There are 4 types of orbitals:

  • S-orbitals are spherical and may hold two electrons each. These are the main bonding orbitals for hydrogen, helium and the alkali (earth) metals.

  • P-orbitals are in three types: px, py and pz and point along the x y and z axes as below. These are the orbitals used in bonding for the p-block elements such as carbon, oxygen, nitrogen, silicon and the halides. Each holds two electrons and they are in sets of three: (Taken from http://www.chemcomp.com/journal/molorbs.htm)
    ao.gif

  • D-orbitals are the main bonding orbitals used by the d-block elements (or transition metals) such as Iron, Tungsten, Platinum, Gold, Iridium, Chromium, Nickel, Copper etc. They come in 5 varieties: dx2-y2, dz2, dxy, dyz and dxz. These all hold two electrons each and in an unbonded atom all have the same energy. However in transition metal complexes an effect called Crystal Field Stabilisation Energy means these orbitals are of different energy, depending on whether the complex is square planar, tetrahedral, octahedral etc.
  • F-orbitals are posessed by the Actinides and Lanthanides (f- block) and are as below (Taken from http://library.tedankara.k12.tr/chemistry/vol3/Atomic orbitals/h14.htm). They do not experience Crystal Field Splitting and are always degenerate (same energy). There are 7 of them.
    h14.jpg

Each "shell" has a number of these orbitals. The first shell is the first line in the periodic table, and moves downwards. For example, Hydrogen has 1 shell - and only an S orbital. This orbital has 1 electron.

The first shell has only an S orbital.
The second shell has S and P orbitals.
The third shell has S, P and D orbitals.
The fourth shell has S, P, D and F orbitals.

The orbitals are not always filled in this order. The lowest energy ones are filled first. ie: 1s before 2s before 2p before 3s before 3p. However the 4s is of lower energy than the 3d. Thus for example Potassium (K) is 1s22s22p63s23p64s1. The first d-block element is Scandium which is 1s22s22p63s23p64s23d1. This is also the case for the 4f orbitals - Pryseodynium is 1s22s22p63s23p64s23d104p65s24d105p66s24f3
 
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The States of Matter

Solids: structure and properties

A solid is an organised, rigid, layout of molecules which can only be slightly reduced in volume when compressed. When heated, these molecules vibrate (the temperature of a material being related to the kinetic energy of its molecules).

Solids can be organised or disorganised. Organised solids are those with a crystilline structure such as Sodium Chloride. Less organised solids consist of a less organised arrangement of molecules - such as Sulphur which usually exists as rings of 8 Sulphur atoms.

A solids melting point is determined by how much energy is required to overcome the forces holding the molecules together. See the section on Intermolecular forces.


Liquids: structure and properties

Liquids are similar to solids, apart from the molecules are constantly moving, and are seperated a bit more.

Brownian motion is the name given to the constant movement of liquid molecules. If you take a completely still glass of water and float a few small flakes of rice or something similar, even though the glass is still the particles still move.

Liquids are less organised than solids. The boiling point of a liquid is determined by how strong the forces between the molecules are.


Gases: structure and properties

Gases consist of molecules spread out and moving quickly about space - they take up more space than the equivelant number of molecules in liquid state. For example, 0.018 litres of Water produces around 22.4 litres of water vapour.

Due to gases being more spread out, they are easier to compress. The pressure, volume and temperature of a gas are governed by the following relation****:

pV = nRT

Where p is pressure (in Pascals), V is volume (litres), n is the number of moles of gas (see above), R is a constant (~8.3) and T is temperature. Thus if volume is held constant (in a fixed vessel perhaps) and the temperature is increased, then so is the pressure.

The following note contributed by Moto7451
Note: P can be in any unit so long as you have the appropiate constant. When doing these calculations make sure you either have the right constant or convert your units to match your constant. There are english system constants as well I believe. Useful list of constants here: http://wine1.sb.fsu.edu/chm1045/notes/Gases/IdealGas/Gases04.htm



The following image is a useful one - displaying the difference between the arrangement of molecules in a sold, liquid and gas. This is water - the first image being the crystalline structure of ice. The second picture is of water molecules moving freely. The latter is of water vapour.
sf2x10b.jpg

This image is from this webpage


Supercritical Fluids

At certain combinations of temperature and pressure (usually very high), a substance can become a supercritical fluid. In this state, the substance has some characteristics of both liquids and gases and is simply called a fluid.

Supercritical fluids can have many uses - including removing caffeine from coffee. The supercritical fluid used dissolves the caffeine and later on this caffeine can be recovered and sold.


Changes in State
A change in occurs as a result of a change in pressure and/or temperature.

Solid -> Liquid occurs when heat is gven to the solid to increase its temperature to its melting point. The reverse happens when the liquid cools below its melting point and gives out energy.

Liquid -> Gas occurs when heat is gven to the liquid to increase its temperature to its boiling point. The reverse happens when the gas cools below its boiling point and gives out energy.

Solid -> Gas occurs usually at relatively low pressure, when heat is gven to the solid to increase its temperature to the appropriate point. The reverse happens when the gas is then cooled.
 
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Inter- and Intra-Molecular Forces

A general note on inter/intra - molecular forces

Intermolecular forces are forces between molecules. These hold substances together, and can determine structures, melting points, boiling points, heat capacities and latent heats.

Intramolecular forces are forces inside molecules, holding their constituent parts together. These determine if/when a substance will decompose, for example.

I have described a number of the inter and intra forces, their range, strength and situations in which they may be found.

Note: Electronegativity is the affinity something has for electrons. Substances like Flourine love electrons so have a high electronegativity - electrons will spend a lot of time around Fluorine compared to less electronegative compounds. Caesium has one of the lowest electronegativities.



Intramolecular Forces
  1. Covalent bonding: Put simply, covalent bonding is the sharing of electrons. When two atoms with similar electronegativities bond together, they will usually just share electrons.
    For example, consider CH4, which has 4 C-H bonds. Carbon has 4 electrons in its outermost "layer" and Hydrogen has 1. Ideally, atoms want either 2 (for Hydrogen and helium) or 8 electrons in their outermost layer. Thus, 4 hydrogens each share their electrons with carbon. Carbon in turn shares one electron with each hydrogen. The end result? Each hydrogen has 2 electrons and each carbon 8. Thus the atoms are happy and the molecule is stable. Covalent bonds tend to result in discrete molecules - Methane exists as many individual CH4 molecules.

  2. Polar Covalent Bonds: These are Covalent bonds which are not entirely "fair". If there is a sufficient difference in electronegativity between the two atoms in the bond, there will be an uneven distribution of electrons. For example, in C=O (carbonyl) bond, as the Carbon nucleus has a +6 charge and the Oxgen has a +8 charge. Thus the electrons spend slightly more time around the Oxygen than round the carbon, meaning the Oxygen is slightly negative and the Carbon is slightly positive. This is one of the main reasons carbonyl chemistry is such a vast and varied topic.

  3. Ionic: If there is a great difference in electronegativity then electrons will actually move completely between atoms. For example, NaCl (Sodium Chloride) is an ionic compound. One electron moves from the Sodium atom (which is then left with 8 in its outermost layer) to the Chlorine atom (which now has 8 electrons also). This gives Na (Sodium) a + charge and Cl (Chlorine) a - charge. This charge difference holds all the ions together.
    In solid form, compounds like NaCl form lattices, of alternating Na and Cl atoms. In solution or in liquid, the ions move freely and thus conduct a current. The formula of such compounds is more a ratio of the number of one ion to the number of another rather than a formula for a number of discrete molecules.

It is also possible for there to be different types of bond formed.
  • Sigma bonds are the first bonds formed between molecules and are the strongest type of bonds. These have good overlap and are formed from molecular orbitals that point directly at each other.

  • Pi bonds are the second and third bonds formed between atoms. These are weaker, bring the atoms closer together, and are easier to break. Such bonds are formed from atomic orbitals with poorer overlap.

  • Delta bonds are the fourth, very weak bond formed when there is very poor orbital overlap, between d orbitals in two transition metal complexes. For example in some Rhodium and Molybdenum complexes where the two metal centres are bonded together.

The difference between these bonds would be difficult to explain fully without going deep into types of orbitals, sp/sp2/sp3 hybridisationa and so forth. I may add this at a later date but until then this web page may be of interest.




Intramolecular Forces
  1. London Forces affect all molecules. Put simply, the electrons orbit the atom but due to their motion and the laws of probability there must be, at some point in time, more electrons on one side of the atom/molecule than on the other. This leads to a very weak, very short lived dipole. If another molecule happens to have at some point, to have a dipole as well, the two molecules may be attracted. The attraction is weak, short lived and near irrelevant if other bonding is involved. This is often the only intermolecular bonding present in elements like the Nobel gases (He, Ne, Ar, Kr, Xe). These elements are gases, as these interactions are so weak.

  2. Hydrogen Bonding is present in compounds involving hydrogen bonded to either Oxygen, Nitrogen or Fluorine. Due to the electronegativity differences between Hydrogen and the above elements, the bond is particularly polar, and the Hydrogen becomes slightly positive and the N/O/F becomes slightly negative. Thus a slightly positive hydrogen can become attracted to a slightly negative Oxygen, Nitrogen or Fluorine on another molecule. This can also cause dimer formation in Carboxylic acids. As below:
    acetic_acid_dimer.jpg

    Image taken from here

  3. Permenant Dipoles: Same idea as Hydrogen Bonding, but not specifically NH/OH/FH bonds. In molecules where there is an electronegativity difference between two atoms the resulting dipole may cause molecules to be attracted to each other.

  4. Metallic Bonding: In metals, the metals share their outermost layer of electrons with each other, resulting in a "pool" of electrons around all the atoms. The fact that the electrons exist in this way is what makes metals so good at conducting electricity. You can apply a potential to the metal, if you push electrons in at one end, electrons will be pushed out of the other.

    A diagram of this is below.
    metalbond.GIF
    Image taken from here
 
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Heat Capacities and Latent Heats

Heat Capacity: definition and uses.

Put simply, the Heat Capacity of a substance is the amount of energy required to raise the temperature of a substance by a given temperature.

If only a small amount of energy is required to raise the temperature by 1C then it has a low heat capacity. If a large amount of heat is required, it has a high heat capactity.

Conversely, if a substance cools by a certain delta T (change in temperature, also often written dT) the heat capacity is the heat released.

Heat capacity (at constant pressure) is written as Cp and is measured in Joules per kilogram per degree Kelvin. Water has a heat capacity of ~4200 J/K/kg meaning that 4,200 Joules (4.2kiloJoules) of energy is required to increase the temperature of 1kilogram (which is 1 litre) of water by 1 degree Kelvin.

A note about Kelvin: a change in temperature of 1K is the same as a change in temperature of 1 degree C. However, 0K = -273.13 degrees C. Room temperature is around 298K, ice melts at approximately 273K, and water boils at approx 373K.

The following expression is useful in calculations:

dE = m.Cp.dT

Where dE is the change in energy of the water. A positive value corresponds to an increase in energy (ie heat in) and a negative value to a decrease in energy (ie heat given out).

m is the mass of water in Kilograms. 1 litre of water weighs 1 kilogram.

Cp is the heat capacity, defined above.

dT is the change in temperature.


Latent heats: What and Why?

If you were to plot a graph of energy put into a mass of ice, vs the temperature of the matter, you would think that the temperature would increase linearly with energy put in, right?

Wrong. Around 273K and 373K you would find that, for a while, even though you are still putting energy in, the temperature does not change! This is due to latent heats of fusion, and vapourisation.

If the ice is heated to a temperature of ~273K it begins to melt and form water. However, this change of state requires energy, this energy is required to break the lattice of H2O that is ice. This is called the Latent Heat of Fusion. This is determined by how easily (or not) the solid structure is broken and converted to a liquid. The higher this is, the more kinetic energy the molecules need to become a liquid.

A similar case is the Latent Heat of Vapourisation. When water is heated to 373K it begins to turn into steam. Extra energy is needed to spread out the molecules, so that the water enters the vapour state. Again, the higher this is, the more kinetic energy the molecules need for the substance to turn into a gas.

When a substance turns from a liquid to a solid, or a gas to a liquid, these latent heats are given out as heat. This is why steam at 100C scalds more than water at 100C. When you are scalded by water at 100C you are burned by the heat stored in the water. With steam you get the latent heat of vapourisation PLUS the stored heat.

The third latent heat is the Latent Heat of Sublimation. This is required to convert a Solid straight to a gas (usually under low pressure and high temperature). This is given out when a gas is converted straight back to a solid. Carbon Dioxide sublimes at atmospheric pressure - "dry ice" sublimes straight to a gas.
 
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Thermodynamics, Entropy and Enthalpy.

The Laws of Thermodynamics

The Laws of Thermodynamics are in bold, with an explaination following.


  • 0. When two systems are in thermal equilibrium with a third body (like a thermometer), they are also in thermal equilibrium with each other. Further, these bodies are all at the same temperature.

    This is our basic reasoning behind temperature scales. Thermal equilibrium meaning once temperature is constant.

  • 1. The change in internal energy of a system(dU) is equal to the heat put into the system(Q), minus the work done by the system(W).

    dU = Q - W


    This is essentially conservation of energy as applied to thermodynamics. If you consider any system, for example a gas in a flexible container, the internal energy (ie kinetic energy of the gas molecules) is equal to the heat put into the system (heating a gas makes the molecules move faster) minus the work done by the system (the work done by the gas molecules stretching the container).


  • 2. During any spontaneous process, the entropy of the Universe increases.

    Basically, everything that happens without being "pushed" will increase the disorganisation of the Universe. Gas expanding, ice melting, etc.


  • 3. The entropy of a pure, perfect crystal at 0K is 0.

    Entropy is explained below. Basically, a perfect, pure crystal at 0K is perfectly organised, and ordered. However, for quantum mechanical reasons, 0K cannot be reached.


Entropy

So what is Entropy? Simply put, it is a measure of disorder. As defined by the third law of thermodynamics, the entropy, the disorder of a substance, is 0 when the substance is in perfect crystalline form (perfectly organised) at 0K (atoms have no kinetic energy whatsoever). Entropy is represented by the letter S. The exact definition of S can be explained using Boltzmanns equation (which is on his tombstone): S = kB loge W, where kb is Boltzmanns constant and W is defined as "... the number of possible microstates corresponding to the macroscopic state of a system — the number of (unobservable) "ways" the (observable) thermodynamic state of a system can be realized by assigning different positions and momenta to the various molecules ..." (Wikipedia). W is a pain to calculate and define and derive althought it can be done.


Substances have an absolute entropy value at each temperature. An entropy change is referred to as dS.

Entropy is measured in Joules per Kelvin per mole.

Enthalpy

Enthalpy is defined as:
Oxford Dictionary of Chemistry said:
A thermodynamic property of a system defined by H= U + pV where H is the enthalpy, U is the internal energy of the system, p is pressure and V is volume. In a chemical reaction carried out in the atmosphere the pressure remains constant and the enthalpy of reaction dH = dU + p(dV). For an exothermic reaction dH < 0.

Thus enthalpy is a measure of the internal energy of a substance. It is impossible to define an "absolute" enthalpy change so thus we deal with changes in enthalpy, denoted dH.

A negative dH implies heat is lost from the system, and is given out into its surroundings. This is otherwise known as an exothermic reaction.

A postive dH implies heat is gained by the system, and is lost from the surroundings. This is known as an endothermic reaction.

There are many different types of enthalpy change, they occur for all reactions. Some examples:
  • Enthalpy of Combustion is the enthalpy change (always -ve as far as I know) when one mole of a substance is burned completely in oxygen.
  • Enthalpy of Dissolution is the enthalpy change when one mole of a substance is dissociated in a solvent.
  • Enthalpy of Formation is the enthalpy change when one mole of a substance is formed from the relevant quantities of its constituent elements.

Enthalpy is measured in Joules per mole.


Gibbs Free Energy and reaction feasability

Gibbs Free Energy is a measure of a systems ability to do work. If in a reaction, the change in Gibbs Free Energy (denoted dG) is negative, it will proceed spontaneously to equilibrium and it said to be feasible.

dG = dH - TdS

Thus the change in Gibbs Free Energy is equal to Enthalpy change minus the temperature (in Kelvin!) times Entropy change.

  • Positive dH, negative dS means the reaction is not feasible.
  • Postive dH, positive dS means the reaction is feasible but only over a certain temperature.
  • Negative dH, postive dS means the reaction is feasible at all temperatures.
  • Negative dH, negative dS means the reaction is feasible but only below a certain temperature.
 
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Properties of Common Materials


Heat Capacities Conductivities

Common Metals
Silver 419 W/m-k
Copper 385 W/m-k
Aluminum 210 W/m-k

Metals used for plating
Nickel 60.7 W/m-k
Chromium 69.1 W/m-k

Used in Thermal Compounds
Zinc Oxide 23.4 W/m-k
Boron Nitride 2-60 W/m-k depending on direction - Not sure about this one.
Aluminum oxide <30 W/m-k

Solders
63% Sn 37% Pb (eutectoid) 50.9 w/m-k melt point: 183 deg C
52% In 48% Sn 34 W/m-k melt point: 118 deg C

Above contributed by redwraith94

Symbols/Forms of the Elements
1 - Hydrogen, H. Exists usually as H2 {gas}
2 - Helium, He. Inert {gas}
3 - Lithium, Li. [metal] {solid}
4 - Beryllium, Be. [metal] {solid}
5 - Boron, B. {solid}
6 - Carbon, C. Exists as graphite or diamond crystals, or as buckminsterfullerenes. {solid}
7 - Nitrogen, N. Exists usually as N2 {gas}
8 - Oxygen, O. Exists usually as O2 {gas}
9 - Fluorine, F. Exists usually as F2. Very reactive. {gas}
10 - Neon, Ne. Inert {gas}
11 - Sodium (Natrium), Na. [metal] {solid}
12 - Magnesium, Mg. [metal] {solid}
13 - Aluminium, Al. [metal] {solid}
14 - Silicon, Si. {solid}
15 - Phosphorous, P. Tends to exit as P4 structures.Will react quickly with oxygen. {solid}
16 - Sulphur, S. Tends to exist as S8 rings {solid}
17 - Chlorine, Cl. Exists as Cl2 {gas}
18 - Argon, Ar. Inert {gas}
19 - Potassium, K. [metal] {solid}
20 - Calcium, Ca. [metal] {solid}
21 - Scandium, Sc [metal] {solid}
22 - Titanium, Ti [metal] {solid}
23 - Vanadium, V [metal] {solid}
24 - Chromium, Cr [metal] {solid}
25 - Manganese, Mn [metal] {solid}
26 - Iron (Ferrate), Fe [metal] {solid}
27 - Cobalt, Co [metal] {solid}
28 - Nickel, Ni [metal] {solid}
29 - Copper (Cuprate), Cu [metal] {solid}
30 - Zinc, Zn [metal] {solid}
31 - Gallium, Ga [metal] {solid}
32 - Germanium, Ge [metal] {solid}
33 - Arsenic, As {solid}
34 - Selenium, Se {solid}
35 - Bromine, Br. Exists as Br2 {liquid}
36 - Krypton, Kr {gas}
37 - Rubidium, Rb [metal] {solid}
38 - Strontium, Sr [metal] {solid}
39 - Yttrium, Y [metal] {solid}
40 - Zirconium, Zr [metal] {solid}
41 - Niobium, Nb [metal] {solid}
42 - Molybdenum, Mo [metal] {solid}
43 - Technetium, Tc [metal] {solid}
44 - Ruthenium, Ru [metal] {solid}
45 - Rhodium, Rh [metal] {solid}
46 - Palladium, Pd [metal] {solid}
47 - Silver (Argentium), Ag [metal] {solid}
48 - Cadmium, Cd [metal] {solid}
49 - Indium, In [metal] {solid}
50 - Tin (Stannum), Sn [metal] {solid}
51 - Antimony (Stibium), Sb [metal] {solid}
52 - Tellurium, Te {solid}
53 - Iodine, I. Exists as I2 {solid}
54 - Xenon, Xe {gas}
55 - Caesium, Cs [metal] {solid}
56 - Barium, Ba [metal] {solid}
57 - Lathanum, La [metal] {solid}
... [lathanides omitted as probably not applicable here]
72 - Hafnium, Hf [metal] {solid}
73 - Tantalum, Ta [metal] {solid}
74 - Tungsten (Wolframium), W [metal] {solid}
75 - Rhenium, Re [metal] {solid}
76 - Osmium, Os [metal] {solid}
77 - Iridium, Ir [metal] {solid}
78 - Platinum, Pt [metal] {solid}
79 - Gold (Aurium), Au [metal] {solid}
80 - Mercury (Hydrargyrum), Hg [metal] {liquid}
81 - Thallium, Tl [metal] {solid}
82 - Lead (Plumbus), Pb [metal] {solid}
83 - Bismuth, Bi [metal] {solid}
84 - Polonium, Po [metal] {solid}
85 - Astatine, At {solid}
86 - Radon, Rn {gas}
87 - Francium, Fr [metal] {solid}
88 - Radium, Ra [metal] {solid}
89 - Actinium, Ac [metal] {solid}
...
92 - Uranium, U [metal] {solid}
...
94 - Plutonium, Pu [metal] {solid}

This site is useful also: http://www.webelements.com/webelements/elements/text/periodic-table/key.html

The Peltier Effect -- by mumrah

A typical TEC (thermoelectric cooler) is comprised of an array of semiconductor-metal junctions that has been cleverly engineered in such a way that there is a side that absorbs heat and a side that releases heat. This has led to the application of TECs in cpu cooling b/c they do a good job of absorbing heat from the cpu and displacing it somewhere else (usually a waterblock).

Any ways, back the the S-N junction (semiconductor-metal). First come some restrictions. In order for an S-N junction to experience the Peltier Effect, the conduction electrons in the semiconductor must exist at a higher energy level than the conduction electrons in the metal. Ok, so now the metal electrons are just chilling. They are in the lowest energy state available and cannot flow into the semiconductor. The reason they cannot flow into the semiconductor is that they encounter a potential "hill" at the junction. This hill comes from the fact that the semiconductor electrons exist at a higher energy level than the metal electrons (recall this is not true in general, just in this case). In order for the metal electrons to flow into the semiconductor, they need to increase their energy.

-enter Peltier-

If you apply a current across a S-N junction, these electrons somehow must overcome this potential barrier. How, you ask? They absorb energy from the phonons, of course! In other words, they absorb thermal energy from the lattice. Now that the electrons have gained this energy from the lattice, they can easily flow across the junction. Current flowing from metal to semiconductor results in heat absorption, and current flowing from semiconductor to metal results in heat radiation (remember, the current flows in the opposite direction of electrons).

So now, you take your clever little TEC, apply a current, and get the hot and cold sides.

Physics Disclaimer: the previous only applies to n-type semiconductors, for p-type simply reverse everything.


Changelog
14/05/05: Edited to add Moto7451's comment about the gas constant, edited Cp to Heat Capacity at constant pressure (more correct). Also added properties of compounds into this post. Added list of major elements and Peltier effect also.

04/03/06: Edited to add delta bonds, definition of entropy expanded, types of atomic orbital, magnetic properties added.
 
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:thup:!

Wow, talk about a super sticky! I have just skimmed it so far, but I will do some more in depth reading soon. I am sure I will learn quite a bit.

I say David deserves a nice round of :clap: .
 
I got myself a revision page on matter for my next chemistry/physics exam. Nice work.
You wrote it up yourself?
 
musawi said:
I got myself a revision page on matter for my next chemistry/physics exam. Nice work.
You wrote it up yourself?

Written up mostly from knowledge (studying for a Chemistry degree), although I admit I had to google to get the laws of thermodynamics in the right order :D.

If anyone has any questions or would like me to explore any area in more detail let me know.
 
A table with the thermal properties of common metals would be useful. And you could add a picture of the periodic table for completeness.
 
RoadWarrior said:
I see no latent heat of solution. :D

You could use it in a kind of liquid convection heatpipe having a supersaturated salt solution.....

Isnt that the same as:
Enthalpy of Dissolution is the enthalpy change when one mole of a substance is dissociated in a solvent.


SuperNade - Ill see what I can come up with :D
 
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