2.2+States+of+matter

Matter: States of Matter
//by Anthony Carpi, Ph.D. // The different states of matter have long confused people. The ancient Greeks were the first to identify three classes (what we now call states) of matter based on their observations of water. But these same Greeks, in particular the philosopher [|Thales] (624 - 545 BCE), incorrectly suggested that since water could exist as a solid, liquid, or even a gas under natural conditions, it must be the single principal [|element] in the universe from which all other substances are made. We now know that water is not the fundamental substance of the universe; in fact, it is not even an element. To understand the different states in which matter can exist, we need to understand something called the //Kinetic Molecular [|Theory] of Matter//. Kinetic Molecular Theory has many parts, but we will introduce just a few here. One of the basic concepts of the theory states that [|atoms] and [|molecules] possess an [|energy] of motion that we perceive as temperature. In other words, atoms and molecules are constantly moving, and we measure the energy of these movements as the temperature of the substance. The more energy a substance has, the more molecular movement there will be, and the higher the perceived temperature will be. An important point that follows this is that the amount of energy that atoms and molecules have (and thus the amount of movement) influences their interaction with each other. Unlike simple billiard balls, many atoms and molecules are attracted to each other as a result of various intermolecular [|forces] such as [|hydrogen bonds], [|van der Waals forces], and others. Atoms and molecules that have relatively small amounts of energy (and movement) will interact strongly with each other, while those that have relatively high energy will interact only slightly, if even at all, with others. How does this produce different states of matter? Atoms that have low [|energy] interact strongly and tend to “lock” in place with respect to other atoms. Thus, collectively, these atoms form a hard substance, what we call a solid. [|Atoms] that possess high energy will move past each other freely, flying about a room, and forming what we call a gas. As it turns out, there are several known states of matter; a few of them are detailed below. Solid matter - ice A simulation of the molecular movement within an ice crystal. (Flash required)
 * [[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-3114-041228031208.jpg width="118" height="165" align="right" caption="water-boiling"]] As a young boy, I remember staring in wonder at a pot of boiling water. Searching for an explanation for the bubbles that formed, I believed for a time that the motion of the hot water drew air down into the pot, which then bubbled back to the surface. Little did I know that what was happening was even more magical than I imagined - the bubbles were not air, but actually water in the form of a gas. ||
 * [[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-3102-041228021211.jpg width="170" height="104" align="right" caption="ice-cubes"]] **Solids** are formed when the attractive [|forces] between individual [|molecules] are greater than the [|energy] causing them to move apart. Individual molecules are locked in position near each other, and cannot move past one another. The atoms or molecules of solids remain in motion. However, that motion is limited to vibrational energy; individual molecules stay fixed in place and vibrate next to each other. As the temperature of a solid is increased, the amount of vibration increases, but the solid retains its shape and volume because the molecules are locked in place relative to each other. To view an example of this, click on the animation below which shows the molecular structure of ice [|crystals]. ||

Liquid matter - water A simulation of molecular movement within liquid water. (Flash required)
 * [[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-3104-041228021213.jpg width="170" height="106" align="right" caption="water-liquid"]] **Liquids** are formed when the energy (usually in the form of heat) of a [|system] is increased and the rigid structure of the solid state is broken down. In liquids, molecules can move past one another and bump into other molecules; however, they remain relatively close to each other like solids. Often in liquids, intermolecular forces (such as the [|hydrogen bonds] shown in the animation below) pull molecules together and are quickly broken. As the temperature of a liquid is increased, the amount of movement of individual molecules increases. As a result, liquids can “flow” to take the shape of their container but they cannot be easily compressed because the molecules are already close together. Thus liquids have an undefined shape, but a defined volume. In the example animation below we see that liquid water is made up of molecules that can freely move past one another, yet remain relatively close in distance to each other. ||

Gaseous matter - steam A simulation of the behaviour of water molecules entering the gas state. (Flash required)
 * [[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-3103-041228021211.jpg width="170" height="101" align="right" caption="clouds"]] **Gases** are formed when the [|energy] in the [|system] exceeds all of the attractive [|forces] between [|molecules]. Thus gas molecules have little interaction with each other beyond occasionally bumping into one another. In the gas state, molecules move quickly and are free to move in any direction, spreading out long distances. As the temperature of a gas increases, the amount of movement of individual molecules increases. Gases expand to fill their containers and have low [|density]. Because individual molecules are widely separated and can move around easily in the gas state, gases can be compressed easily and they have an undefined shape. ||

Solids, liquids, and gases are the most common states of matter that exist on our planet. If you would like to compare the three states to one another, click on the comparison animation below. Note the differences in molecular motion of water [|molecules] in these three states. Solid-Liquid-Gas Comparison

As we have seen, increasing [|energy] leads to more molecular motion. Conversely, decreasing energy results in less molecular motion. As a result, one prediction of Kinetic Molecular [|Theory] is that if we continue to decrease the energy (measured as temperature) of a substance, we will reach a point at which all molecular motion stops. The temperature at which molecular motion stops is called **absolute zero** and has been calculated to be -273.15 degrees Celsius. While scientists have cooled substances to temperatures close to [|absolute zero], they have never actually reached absolute zero. The difficulty with observing a substance at absolute zero is that to “see” the substance, [|light] is needed, and light itself transfers energy to the substance, thus raising the temperature. Despite these challenges, scientists have recently observed a fifth state of matter that only exists at temperatures very close to absolute zero. Several other less common states of matter have also either been described or actually seen. Some of these states include liquid [|crystals], fermionic condensates, superfluids, supersolids and the aptly named strange matter. To read more about these phases, visit the Phase page of Wikipedia, linked to below in the //Further Exploration// section. The transformation of one state of matter into another state is called a phase transition. The more common phase transitions even have names; for example, the terms //melting// and //freezing// describe phase transitions between the solid and liquid state, and the terms //evaporation// and //condensation// describe transitions between the liquid and gas state. Phase transitions occur at very precise points, when the [|energy] (measured as temperature) of a substance in a given state exceeds that allowed in the state. For example, liquid water can exist at a range of temperatures. Cold drinking water may be around 4ºC. Hot shower water has more energy and thus may be around 40ºC. However, at 100°C under normal conditions, water will begin to undergo a phase transition into the gas phase. At this point, energy introduced into the liquid will not go into increasing the temperature; it will be used to send [|molecules] of water into the gas state. Thus, no matter how high the flame is on the stove, a pot of boiling water will remain at 100ºC until all of the water has undergone transition to the gas phase. The excess energy introduced by a high flame will accelerate the liquid-to-gas transition; it will not change the temperature. The [|heat] curve below illustrates the corresponding changes in energy (shown in calories) and temperature of water as it undergoes a phase transition between the liquid and gas states. ||
 * [[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-820-030605010650.jpg width="150" height="141" align="right" caption="sun"]] **Plasmas** are hot, ionized gases. Plasmas are formed under conditions of extremely high [|energy], so high, in fact, that [|molecules] are ripped apart and only free [|atoms] exist. More astounding, plasmas have so much energy that the outer [|electrons] are actually ripped off of individual atoms, thus forming a gas of highly energetic, charged [|ions]. Because the atoms in plasma exist as charged ions, plasmas behave differently than gases, thus representing a fourth state of matter. Plasmas can be commonly seen simply by looking upward; the high energy conditions that exist in stars such as our sun force individual atoms into the plasma state. ||
 * Bose-Einstein Condensates** represent a fifth state of matter only seen for the first time in 1995. The state is named after [|Satyendra Nath Bose] and [|Albert Einstein] who predicted its existence in the 1920’s. B-E condensates are gaseous superfluids cooled to temperatures very near [|absolute zero]. In this weird state, all the [|atoms] of the condensate attain the same quantum-mechanical state and can flow past one another without friction. Even more strangely, B-E condensates can actually “trap” [|light], releasing it when the state breaks down.
 * [[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-3114-041228031208.jpg width="118" height="165" align="right" caption="water"]] **Phase Transitions**


 * || <span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: left;">[[image:http://www.visionlearning.com/library/modules/mid120/Image/VLObject-3127-041229031208.jpg caption="graph2 - heat curve"]] ||  ||

<span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: left;">As can be seen in the graph above, as we move from left to right, the temperature of liquid water increases as [|energy] (heat) is introduced. At 100ºC, water begins to undergo a phase transition and the temperature remains constant even as energy is added (the flat part of the graph). The energy that is introduced during this period goes toward breaking intermolecular [|forces] so that individual water [|molecules] can “escape” into the gas state. Finally, once the transition is complete, if further energy is added to the [|system], the [|heat] of the gaseous water, or steam, will increase. This same process can be seen in reverse if we simply look at the graph above starting on the right side and moving left. As steam is cooled, the movement of gaseous water [|molecules] and thus temperature will decrease. When the gas reaches 100ºC, more [|energy] will be lost from the [|system] as the attractive [|forces] between molecules reform; however the temperature remains constant during the transition (the flat part of the graph). Finally, when condensation is complete, the temperature of the liquid will begin to fall as energy is withdrawn. Phase transitions are an important part of the world around us. For example, the [|energy] withdrawn when perspiration evaporates from the surface of your skin allows your body to correctly regulate its temperature during hot days. Phase transitions play an important part in geology, influencing [|mineral] formation and possibly even [|earthquakes]. And who can ignore the phase transition that occurs at about -3ºC, when cream, perhaps with a few strawberries or chocolate chunks, begins to form solid ice cream. Now we understand what is happening in a pot of boiling water. The [|energy] (heat) introduced at the bottom of the pot causes a localized phase transition of liquid water to the gaseous state. Because gases are less dense than liquids, these localized phase transitions form pockets (or bubbles) of gas, which rise to the surface of the pot and burst. But nature is often more magical than our imaginations. Despite all that we know about the states of matter and phase transitions, we still cannot predict where the individual bubbles will form in a pot of boiling water.

<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">taked from:[]

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<span style="color: #ff0000; display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//taken from:// <span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">// http://www.youtube.com/watch?v=ezCwUzFsoDY <span style="color: #ff0000; font-family: Arial,Helvetica,sans-serif;"> taken from: http://www.youtube.com/watch?v=pKvo0XWZtjo// <span style="font-family: Arial,Helvetica,sans-serif;"> . <span style="font-family: 'Arial Black',Gadget,sans-serif; font-size: 130%;">properties of matter <span style="font-family: Arial,Helvetica,sans-serif;"> Matter is everything around you. **Matter** is anything made of [|atoms] and molecules. Matter is anything that has a **mass**. Matter is also related to light and electromagnetic radiation. Even though matter can be found all over the universe, you usually find it in just a few forms. As of 1995, scientists have identified five **states of matter**. They may discover one more by the time you get old.

You should know about solids, liquids, gases, plasmas, and a new one called Bose-Einstein condensates. The first four have been around a long time. The scientists who worked with the [|Bose-Einstein condensate] received a Nobel Prize for their work in 1995. But what makes a state of matter? It's about the physical state of molecules and atoms.

=<span style="font-family: Arial,Helvetica,sans-serif;">Changing States of Matter =

[|Elements] and compounds can move from one [|physical state] to another and not change. Oxygen (O2) as a gas still has the same properties as liquid oxygen. The [|liquid] state is colder and denser but the molecules are still the same. Water is another example. The **compound** water is made up of two hydrogen (H) atoms and one oxygen (O) atom. It has the same molecular structure whether it is a [|gas], liquid, or [|solid]. Although its physical state may change, its chemical state remains the same.

So you ask, "What is a chemical state?" If the formula of water were to change, that would be a **chemical change**. If you added another oxygen atom, you would make hydrogen peroxide (H2O2). Its molecules would not be water anymore. Changing states of matter is about changing densities, pressures, temperatures, and other physical properties. The basic chemical structure does not change. //<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;"> <span style="font-family: Arial,Helvetica,sans-serif;">taken from:http://www.chem4kids.com/files/matter_intro.html //

There are five main states of matter. Solids, liquids, gases, plasmas, and Bose-Einstein condensates are all different states of matter. Each of these states is also known as a phase. Elements and compounds can move from one phase to another phase when special **physical forces** are present. One example of those forces is temperature. The phase or state of matter can change when the temperature changes. Generally, as the temperature rises, matter moves to a more active state.



Phase describes a physical state of matter. The key word to notice is physical. Things only move from one phase to another by physical means. If energy is added (like increasing the temperature or increasing pressure) or if energy is taken away (like freezing something or decreasing pressure) you have created a physical change.



One compound or element can move from phase to phase, but still be the same substance. You can see water **vapor** over a boiling pot of water. That vapor (or gas) can **condense** and become a drop of water. If you put that drop in the freezer, it would become a solid. No matter what phase it was in, it was always water. It always had the same chemical properties. On the other hand, a chemical change would change the way the water acted, eventually making it not water, but something completely new

<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//[|taken from: http://www.chem4kids.com/files/matter_states.html]//

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 150%; text-align: center;">video:states of matter
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<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//taked from: [|http://video.google.com/videoplay?docid=6483201716001918124&ei=FwK]

[|HS-vOMqC6rAK48Z2PCw&q=states+of+matter&hl=es&view=3#]//

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 169%; text-align: center;">States of Matter All the material on earth is in three states-**solid, liquid, and gas**. The "state" of the matter refers to the group of matter with the same properties.

<span style="font-family: Arial,Helvetica,sans-serif;"> <span style="color: #ff0000; font-family: Arial,Helvetica,sans-serif;">Solid <span style="font-family: Arial,Helvetica,sans-serif;">A solid has a certain size and shape. The wood block does not change size or shape. Examples of solids are the computer, the desk, and the floor.

<span style="color: #ff0000; font-family: Arial,Helvetica,sans-serif;">Liquids <span style="font-family: Arial,Helvetica,sans-serif;">It has a size or volume. Volume means it takes up space. It takes the shape of its container. Liquids can flow, be poured, and spilled.

<span style="color: #ff0000; font-family: Arial,Helvetica,sans-serif;">Gases Gases do not have volume,the volume of gases are the volume of their containers.Gases are all around you. You can feel gas when the wind blows. The wind is moving air. Air is many gases mixed together.

<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">**// taken from //**:[]

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<span style="color: #000000; display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">**// 0Taken from: http://www.youtube.com/watch?v=s-KvoVzukHo//** media type="youtube" key="V9WYweBA6vA" width="425" height="350" align="center"

<span style="color: #000000; display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">**// Taken From: //**http://www.youtube.com/watch?v=V9WYweBA6vA

<span style="color: #ff0000; font-family: Tahoma,Geneva,sans-serif; font-size: 16pt;">The Matter <span style="font-family: Arial,Helvetica,sans-serif;">**<span style="font-family: 'Tahoma','sans-serif'; font-size: 16pt; line-height: 115%;">States of matter **<span style="font-family: 'Tahoma','sans-serif'; font-size: 16pt; line-height: 115%;"> are the distinct forms that different phases of matter take on.

<span style="color: #ffc000; font-family: 'Tahoma','sans-serif'; font-size: 16pt; line-height: 115%;">-solid -liquid -gas -plasma

<span style="font-family: 'Tahoma','sans-serif'; font-size: 16pt;"> More recently, distinctions between states have been based on differences in molecular interrelationships. Solid is the state in which intermolecular attractions keep the molecules in fixed spatial relationships. Liquid is the state in which intermolecular attractions keep molecules in proximity, but do not keep the molecules in fixed relationships. Gas is that state in which the molecules are comparatively separated and intermolecular attractions have relatively little effect on their respective motions. Plasma is a highly ionized gas that occurs at high temperatures. The intermolecular forces created by ionic attractions and repulsions give these compositions distinct properties, for which reason plasma is described as a fourth state of matter. <span style="color: #000000; font-family: 'Tahoma','sans-serif'; font-size: 16pt;">Forms of matter that are not composed of molecules and are organized by different forces can also be considered different states of matter.
 * <span style="color: #000000; font-family: 'Tahoma','sans-serif'; font-size: 16pt;">States of matter **<span style="font-family: 'Tahoma','sans-serif'; font-size: 16pt;"> may also be defined in terms of phase transitions. A phase transition indicates a change in structure and can be recognized by an abrupt change in properties. By this definition, a distinct state of matter is any set of states distinguished from any other set of states by a phase transition. Water can be said to have several distinct solid states.



<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//taked from: [] []//

Solid
The particles (ions, atoms or molecules) are packed closely together. The forces between particles are strong enough so that the particles cannot move freely but can only vibrate. As a result, a solid has a stable, definite shape, and a definite volume. Solids can only change their shape by force, as when broken or cut.

A crystalline solid: atomic resolution image of strontium titanate. Brighter atoms are Sr and darker ones are Ti. In crystalline solids, the particles (atoms, molecules, or ions) are arranged in an ordered three-dimensional structure. There are many different crystal structures, and the same substance can have more than one structure (or solid phase). For example, iron has a body-centred cubic structure at temperatures below 912 °C, and a face-centred cubic structure between 912 and 1394 °C. Ice has fifteen known crystal structures, or fifteen solid phases which exist at various temperatures and pressures.

Solids can be transformed into liquids by melting, and liquids can be transformed into solids by freezing. Solids can also change directly into gases through the process of sublimation.

Liquid
Structure of a classical monatomic liquid. Atoms have many nearest neighbors in contact, yet no long-range order is present. ‎ The volume is definite if the temperature and pressure are constant. When a solid is heated above its melting point, it becomes liquid. Intermolecular (or interatomic or interionic) forces are still important, but the molecules have enough energy to move relative to each other and the structure is mobile. This means that the shape of a liquid is not definite but is determined by its container. The volume is usually greater than that of the corresponding solid, the most well known exception being water, H2O. The highest temperature at which a given liquid can exist is its critical temperature.

Gas
In a gas, the molecules have enough kinetic energy so that the effect of intermolecular forces is small (or zero for an ideal gas), and the typical distance between neighboring molecules is much greater than the molecular size. A gas has no definite shape or volume, but occupies the entire container in which it is confined. A liquid may be converted to a gas by heating at constant pressure to the boiling point, or else by reducing the pressure at constant temperature. At temperatures below its critical temperature, a gas is also called a vapor, and can be liquefied by compression alone without cooling. A vapor can exist in equilibrium with a liquid (or solid), in which case the gas pressure equals the vapor pressure of the liquid (or solid). A supercritical fluid (SCF) is a gas whose temperature and pressure are above the critical temperature and critical pressure respectively. It has the physical properties of a gas, but its high density confers solvent properties in some cases which lead to useful applications. For example, supercritical carbon dioxide is used to extract caffeine in the manufacture of decaffeinated coffee.

<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//<span style="font-family: Arial,Helvetica,sans-serif; font-size: 12pt;">Taken from: http://en.wikipedia.org/wiki/State_of_matter //

<span style="font-family: Arial,Helvetica,sans-serif;">Gases, liquids and solids are all made up of microscopic particles, but the behaviors of these particles differ in the three phases. The following figure illustrates the microscopic differences. || <span style="font-family: Arial,Helvetica,sans-serif;"> || <span style="font-family: Arial,Helvetica,sans-serif;"> || <span style="font-family: Arial,Helvetica,sans-serif;">Note that: <span style="font-family: Arial,Helvetica,sans-serif;">Liquids and solids are often referred to as **//condensed phases//** because the particles are very close together. The following table summarizes properties of gases, liquids, and solids and identifies the microscopic behavior responsible for each property.
 * <span style="display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: center;">States of Matter **
 * <span style="font-family: Arial,Helvetica,sans-serif;">[[image:../liquids/gas.gif width="104" height="104" align="center" caption="Microscopic view of a gas"]]
 * <span style="font-family: Arial,Helvetica,sans-serif;">Microscopic view of a gas. || <span style="font-family: Arial,Helvetica,sans-serif;">Microscopic view of a liquid. || <span style="font-family: Arial,Helvetica,sans-serif;">Microscopic view of a solid. ||
 * <span style="font-family: Arial,Helvetica,sans-serif;">Particles in a:
 * <span style="font-family: Arial,Helvetica,sans-serif;">gas are well separated with no regular arrangement.
 * <span style="font-family: Arial,Helvetica,sans-serif;">liquid are close together with no regular arrangement.
 * <span style="font-family: Arial,Helvetica,sans-serif;">solid are tightly packed, usually in a regular pattern.
 * <span style="font-family: Arial,Helvetica,sans-serif;">Particles in a:
 * <span style="font-family: Arial,Helvetica,sans-serif;">gas vibrate and move freely at high speeds.
 * <span style="font-family: Arial,Helvetica,sans-serif;">liquid vibrate, move about, and slide past each other.
 * <span style="font-family: Arial,Helvetica,sans-serif;">solid vibrate (jiggle) but generally do not move from place to place.

particles can move past one another || <span style="font-family: Arial,Helvetica,sans-serif;">assumes the shape of the part of the container which it occupies particles can move/slide past one another || <span style="font-family: Arial,Helvetica,sans-serif;">retains a fixed volume and shape rigid - particles locked into place || lots of free space between particles || <span style="font-family: Arial,Helvetica,sans-serif;">not easily compressible little free space between particles || <span style="font-family: Arial,Helvetica,sans-serif;">not easily compressible little free space between particles || particles can move past one another || <span style="font-family: Arial,Helvetica,sans-serif;">flows easily particles can move/slide past one another || <span style="font-family: Arial,Helvetica,sans-serif;">does not flow easily rigid - particles cannot move/slide past one another ||
 * ~ **<span style="font-family: Arial,Helvetica,sans-serif;">Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior ** ||
 * **<span style="font-family: Arial,Helvetica,sans-serif;">gas ** || **<span style="font-family: Arial,Helvetica,sans-serif;">liquid ** || **<span style="font-family: Arial,Helvetica,sans-serif;">solid ** ||
 * <span style="font-family: Arial,Helvetica,sans-serif;">assumes the shape and volume of its container
 * <span style="font-family: Arial,Helvetica,sans-serif;">compressible
 * <span style="font-family: Arial,Helvetica,sans-serif;">flows easily

<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//taken from:[|//http://www.chem.purdue.edu/gchelp/atoms/states.html//]//

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<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">taken fromwww.youtube.com/watch?v=HAPc6JH85pM: http:

//<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: center; vertical-align: sub;">video of states of the matter (mario bros version) //

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<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//<span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: left;">taken from:http://www.youtube.com/watch?v=hfb5bA_A6zE&feature=related //



<span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: right;">taken from:http://www.chem.ufl.edu/~itl/2045/lectures/FG11_001.GIF

=<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: center;">The States of Matter = <span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: left;">Now that you know a bit about chemical bonding, let’s talk about the different forms that groups of molecules can take. In other words, let’s talk about the states of matter. The states of matter that you’ll need to know for the SAT II Chemistry test are solid, liquid, and gas. You might wonder why there are different states of matter at all. After all, molecules only bond together in one way, right? The answer lies in the type of intramolecular and intermolecular forces that exist both within and between molecules of substances.



<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//taken from: []//

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: center;">Solids
As we mentioned above, the molecules that make up solids are generally held together by ionic or strong covalent bonding, and the attractive forces between the atoms, ions, or molecules in solids are very strong. In fact, these forces are so strong that particles in a solid are held in fixed positions and have very little freedom of movement. Solids have definite shapes and definite volumes and are not compressible to any extent. There are a few types of solids that you should be familiar with for the SAT II Chemistry test, and we’ve listed them below. However, we will start by saying that there are two main categories of solids—crystalline solids and amorphous solids. **Crystalline solids** are those in which the atoms, ions, or molecules that make up the solid exist in a regular, well-defined arrangement. The smallest repeating pattern of crystalline solids is known as the **unit cell**, and unit cells are like bricks in a wall—they are all identical and repeating. The other main type of solids are called the amorphous solids. **Amorphous solids** do not have much order in their structures. Though their molecules are close together and have little freedom to move, they are not arranged in a regular order as are those in crystalline solids. Common examples of this type of solid are glass and plastics.There are four types of crystalline solids, all of which you should be familiar with for the exam.**Ionic solids—**Made up of positive and negative ions and held together by electrostatic attractions. They’re characterized by very high melting points and brittleness and are poor conductors in the solid state. An example of an ionic solid is table salt, NaCl.**Molecular solids—**Made up of atoms or molecules held together by London dispersion forces, dipole-dipole forces, or hydrogen bonds. Characterized by low melting points and flexibility and are poor conductors. An example of a molecular solid is sucrose.**Covalent-network (also called atomic) solids—**Made up of atoms connected by covalent bonds; the intermolecular forces are covalent bonds as well. Characterized as being very hard with very high melting points and being poor conductors. Examples of this type of solid are diamond and graphite, and the fullerenes. As you can see below, graphite has only 2-D hexagonal structure and therefore is not hard like diamond. The sheets of graphite are held together by only weak London forces!**Metallic solids—**Made up of metal atoms that are held together by metallic bonds. Characterize d by high melting points, can range from soft and malleable to very hard, and are good conductors of electricity. <span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">take from: []

= State of matter =

From Wikipedia, the free encyclopedia
Jump to: [|navigation], [|search] This diagram shows the nomenclature for the different phase transitions. More recently, distinctions between states have been based on differences in molecular interrelationships. [|Solid] is the state in which intermolecular attractions keep the molecules in fixed spatial relationships. [|Liquid] is the state in which intermolecular attractions keep molecules in proximity, but do not keep the molecules in fixed relationships. [|Gas] is that state in which the molecules are comparatively separated and intermolecular attractions have relatively little effect on their respective motions. [|Plasma] is a highly ionized gas that occurs at high temperatures. The intermolecular forces created by ionic attractions and repulsions give these compositions distinct properties, for which reason plasma is described as a fourth state of matter.[|[1]][|[2]] Forms of matter that are not composed of molecules and are organized by different forces can also be considered different states of matter. [|Fermionic condensate] and the [|quark–gluon plasma] are examples.
 * States of matter** are the distinct forms that different [|phases] of matter take on. Historically, the distinction is made based on qualitative differences in bulk properties. [|Solid] is the state in which matter maintains a fixed volume and shape; [|liquid] is the state in which matter maintains a fixed volume but adapts to the shape of its container; and [|gas] is the state in which matter expands to occupy whatever volume is available.
 * States of matter** may also be defined in terms of [|phase transitions]. A phase transition indicates a change in structure and can be recognized by an abrupt change in properties. By this definition, a distinct state of matter is any set of [|states] distinguished from any other set of states by a [|phase transition]. Water can be said to have several distinct solid states.[|[3]] The appearance of superconductivity is associated with a phase transition, so there are [|superconductive] states. Likewise, [|liquid crystal] states and [|ferromagnetic] states are demarcated by phase transitions and have distinctive propertie

<span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//take from:http://en.wikipedia.org/wiki/State_of_matter// solid state **Solids have a fixed volume and shape.**
 * In a solid, the atoms (or molecules) are in fixed positions relative to one another. They vibrate but stay in relative position. When the solid is heated the atoms vibrate faster. This causes the solid to grow slightly in size. It is said to expand. All solids expand on heating, especially metals. Railway lines can buckle on hot days because of this expansion.**
 * If the atoms are arranged in a regular sequence, a crystal results. Metals and salts are crystaline. If the atoms are arranged haphazardly the solid is said to be amorphous (Greek for without shape). Glass is a good example of an amorphous solid.**
 * Metals are crystaline because their atoms are arranged in a regular sequence. However, atoms of metals tend to have loose outer electrons. These electrons can move between atoms. This causes the metal to conduct electricity.**
 * taken from:http://www.krysstal.com/states.html **

<span style="font-family: Arial,Helvetica,sans-serif;">All matter is made from atoms with the configuration of the atom, the number of protons, neutrons, and electrons, determining the kind of matter present (oxygen, lead, silver, neon ...). Every substance has a unique number of protons, neutrons, and electrons. Oxygen, for example, has 8 protons, 8 neutrons, and 8 electrons. Individual atoms can combine with other atoms to form molecules. Water molecules contain two atoms of hydrogen **H** and one atom of oxygen **O** and is chemically called **H2O**. Oxygen and nitrogen, which are the major components of air, occur in nature as **diatomic** (two atom) molecules. Regardless of the type of molecule, matter normally exists as either a **solid, a liquid, or a gas**. We call this property of matter the **state** of the matter. The three normal states of matter have unique characteristics which are listed on the slide. In a **solid** the molecules are closely bound to one another by molecular forces. A solid holds its shape and the [|volume] of a solid is fixed by the shape of the solid. In a **liquid** the molecular forces are weaker than in a solid. A liquid will take the shape of its container with a free surface in a gravitational field. In microgravity, a liquid forms a ball inside a free surface. Regardless of gravity, a liquid has a fixed volume. In a **gas** the molecular forces are very weak. A gas fills its container, taking both the shape and the volume of the container. Liquids and gases are called **fluids** because they can be made to flow, or move. In any fluid, the molecules themselves are in constant, random motion, colliding with each other and with the walls of any container. The motion of fluids and the reaction to external forces are described by the [|Navier-Stokes Equations], which express a conservation of [|mass], [|momentum], and [|energy]. The motion of solids and the reaction to external forces are described by [|Newton's Laws of Motion]. Any substance can occur in any state. Under [|standard atmospheric conditions], water exists as a liquid. But if we lower the [|temperature]below 0 degrees Celsius, or 32 degrees Fahrenheit, water changes its state into a solid called ice. Similarly, if we [|heat]a volume of water above 100 degrees Celsius, or 212 degrees Fahrenheit, water changes its state into a gas called water vapor. Changes in the state of matter are **physical changes**, not chemical changes. A molecule of water vapor has the same chemical composition, **H2O**, as a molecule of liquid water or a molecule of ice. When studying [|gases], we can investigate the motions and interactions of individual molecules, or we can investigate the large scale action of the gas as a whole. Scientists refer to the large scale motion of the gas as the **macro scale** and the individual molecular motions as the **micro scale**. Some phenomenon are easier to understand and explain based on the macro scale, while other phenomenon are more easily explained on the micro scale. Macro scale investigations are based on things that we can easily [|observe and measure]. But micro scale investigations are based on rather simple [|theories]because we cannot actually observe an individual gas molecule in motion. Macro scale and micro scale investigations are just two views of the same thing. The three normal states of matter listed on the slide have been known for many years and studied in physics and chemistry classes. In recent times, we have begun to study matter at the very high temperatures and pressures which typically occur on the Sun, or during re-entry from space. Under these conditions, the atoms themselves begin to break down; electrons are stripped from their orbit around the nucleus leaving a positively charged **ion** behind. The resulting mixture of neutral atoms, free electrons, and charged ions is called a **plasma**. A plasma has some unique qualities that causes scientists to label it a "fourth state" of matter. A plasma is a fluid, like a liquid or gas, but because of the charged particles present in a plasma, it responds to and generates electro-magnetic forces. There are fluid dynamic equations, called the Boltzman equations, which include the electro-magnetic forces with the normal fluid forces of the Navier-Stokes equations. NASA is currently doing research into the use of plasmas for an ion propulsion system.
 * <span style="font-family: Arial,Helvetica,sans-serif;">
 * Solid**
 * Liquid**
 * Gas**
 * Fluids (Liquids and Gases)**
 * Plasma - the "fourth state"**

|| <span style="display: block; font-family: Arial,Helvetica,sans-serif; font-size: 80%; text-align: right;">//taken from:http://www.grc.nasa.gov/WWW/K-12/airplane/state.html//

Solid
Main article: [|Solid] The particles (ions, atoms or molecules) are packed closely together. The forces between particles are strong enough so that the particles cannot move freely but can only vibrate. As a result, a solid has a stable, definite shape, and a definite volume. Solids can only change their shape by force, as when broken or cut. A crystalline solid: atomic resolution image of [|strontium titanate]. Brighter atoms are [|Sr] and darker ones are [|Ti]. In [|crystalline solids], the particles (atoms, molecules, or ions) are arranged in an ordered three-dimensional structure. There are many different [|crystal structures], and the same substance can have more than one structure (or solid phase). For example, [|iron] has a [|body-centred cubic] structure at temperatures below 912 °C, and a [|face-centred cubic] structure between 912 and 1394 °C. [|Ice] has fifteen known crystal structures, or fifteen solid phases which exist at various temperatures and pressures. Solids can be transformed into liquids by melting, and liquids can be transformed into solids by freezing. Solids can also change directly into gases through the process of [|sublimation].

Liquid
Main article: [|Liquid] Structure of a classical monatomic liquid. Atoms have many nearest neighbors in contact, yet no long-range order is present. ‎ The volume is definite if the [|temperature] and [|pressure] are constant. When a solid is heated above its [|melting point], it becomes liquid. Intermolecular (or interatomic or interionic) forces are still important, but the molecules have enough energy to move relative to each other and the structure is mobile. This means that the shape of a liquid is not definite but is determined by its container. The volume is usually greater than that of the corresponding solid, the most well known exception being water, H2O. The highest temperature at which a given liquid can exist is its [|critical temperature].

Gas
Main article: [|Gas] In a gas, the molecules have enough [|kinetic energy] so that the effect of intermolecular forces is small (or zero for an [|ideal gas]), and the typical distance between neighboring molecules is much greater than the molecular size. A gas has no definite shape or volume, but occupies the entire container in which it is confined. A liquid may be converted to a gas by heating at constant pressure to the [|boiling point], or else by reducing the pressure at constant temperature. At temperatures below its [|critical temperature], a gas is also called a [|vapor], and can be liquefied by compression alone without cooling. A vapor can exist in equilibrium with a liquid (or solid), in which case the gas pressure equals the [|vapor pressure] of the liquid (or solid). A [|supercritical fluid] (SCF) is a gas whose temperature and pressure are above the critical temperature and [|critical pressure] respectively. It has the physical properties of a gas, but its high density confers solvent properties in some cases which lead to useful applications. For example, [|supercritical carbon dioxide] is used to [|extract] [|caffeine] in the manufacture of [|decaffeinated] coffee.



Taken From:

- http://en.wikipedia.org/wiki/State_of_matter - http://www.ilpi.com/msds/ref/gifs/statesofmatter.gif - http://www.suntrek.org/images/states.gif

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">states of matter <span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">introducction: Troise Carmine] Gases, liquids, and especially solids surround us and give form to our world. Chemistry at its most fundamental level is about atoms and the forces that act between them to form larger structural units. But the matter that we experience with our senses is far removed from this level. This unit will help you see how these //macroscopic// properties of matter depend on the //microscopic// particles of which it is composed.

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">clasification: 1 Solids, liquids and gases What distinguishes solids, liquids, and gases– the three major //states of matter//— from each other? Let us begin at the microscopic level, by reviewing what we know about gases, the simplest state in which matter can exist. At ordinary pressures, the molecules of a gas are so far apart that intermolecular forces have an insignificant effect on the random thermal motions of the individual particles. As the temperature decreases and the pressure increases, intermolecular attractions become more important, and there will be an increasing tendency for molecules to form temporary clusters. These are so short-lived, however, that even under extreme conditions, gases cannot be said to possess “structure” in the usual sense. The contrast at the level between solids, liquids and gases is most clearly seen in the simplified schematic views above. The molecular units of crystalline solids tend to be highly ordered, with each unit occupying a fixed position with respect to the others. In liquids, the molecules are able to slip around each other, introducing an element of disorder and creating some void spaces that decrease the density. Gases present a picture of almost total disorder, with practically no restrictions on where any one molecule can be.

Solids, liquids and gases: how to tell them apart
Having lived our lives in a world composed of solids, liquids, and gases, few of us ever have any difficulty deciding into which of these categories a given sample of matter falls. Our decision is most commonly based on purely visual cues: Our experience also tells us that these categories are quite distinct; a phase, which you will recall is a region of matter having uniform intensive properties, is either a gas, a liquid, or a solid. Thus the three states of matter are not simply three points on a continuum; when an ordinary solid melts, it usually does so at a definite temperature, without apparently passing through any states that are intermediate between a solid and a liquid.
 * a gas is transparent and has no definite boundaries other than those that might be imposed by the walls of a confining vessel.
 * Liquids and solids possess a clearly delineated that gives solids their definite shapes and whose light-reflecting properties enable us to distinguish one phase from another.
 * Solids can have any conceivable shape, and their surfaces are usually too irregular to show //specular// (mirror-like) reflection of light. Liquids, on the other hand, are //mobile//; except when in the form of tiny droplets, liquids have no inherent shape of their own, but assume the shape of their container and show an approximately flat upper surface.

Although these common-sense perceptions are usually correct, they are not infallible, and in fact there are gases that are not transparent, there are solids such as and many plastics that do not have sharp melting points, but instead undergo a gradual transition from solid to liquid known as //softening//, and when subject to enough pressure, solids can exhibit something of the flow properties of liquids (glacial ice, for example).

Some solids can flow — slowly![[image:http://www.chem1.com/acad/webtext/states/state-images/YukonG.jpg width="350" height="233" caption="glacial flow"]]
A more scientific approach would be to compare the physical properties of the three states of matter, but even here we run into difficulty. It is true, for example, that the density of a gas is usually about a thousandth of that of the liquid or solid at the same temperature and pressure; thus one gram of water vapor at 100°C and 1 atm pressure occupies a volume of 1671 mL; when it condenses to liquid water at the same temperature, it occupies only 1.043 mL.

(42 cm3/mol excluded volume) || The table at the left compares the molar volumes of **neon** in its three states. For the gaseous state, P = 1 atm and T = 0°C. The excluded volume is the volume actually taken up by the neon atoms according to the van der Waals model.
 * **gas** || 22,400 cm3/mol total volume
 * **liquid** || 16.8 cm3/mol ||
 * **solid** || 13.9 cm3/mol ||

It is this extreme contrast with the gaseous states that leads to the appellation “condensed states of matter” for liquids and solids. However, gases at very high pressures can have densities that exceed those of other solid and liquid substances, so density alone is not a sufficiently comprehensive criterion for distinguishing between the gaseous and condensed states of matter. Similarly, the density of a //solid// is usually greater than that of the corresponding liquid at the same temperature and pressure, but not always: you have certainly seen ice floating on water!

Compare the density of gaseous xenon (molar mass 131 g) at 100 atm and 0°C with that of a hydrocarbon liquid for which ρ = 0.104 g/mL at the same temperature. //Solution:// For simplicity, we will pretend that xenon approximates an ideal gas under these conditions, which it really does not. The ideal molar volume at 0° C and 1 atm is 22.4 L; at 100 atm, this would be reduced to .22 L or 220 mL, giving a density ρ = (131 g) / (224 mL) = 0.58 g/mL. In his autobiographical //Uncle Tungsten//, the physician/author Oliver Sacks describes his experience with xenon-filled balloons of //"astonishing density — as near to 'lead balloons" as could be// [imagined]. //If one twirled these xenon balloons in one's hand, then stopped, the heavy gas, by its own momentum, would continue rotating for a minute, almost as if it were a liquid."// Other physical properties, such as the, surface tension, and viscosity, are somewhat more useful for distinguishing between the different states of matter. Even these, however, provide no well-defined dividing lines between the various states. Rather than try to develop a strict scheme for classifying the three states of matter, it will be more useful to simply present a few generalizations.
 * Problem example 1**

If you don't know the meaning of one of the terms in the "property" column in the table, just move your mouse over it to bring up a brief definition. Some of these deal with macroscopic properties (that is, properties such as the //density// that relate to //bulk matter//), and others with microscopic properties that refer to the individual molecular units.
 * Relative magnitudes of some properties of the three states of matter** || **property** || **gas** || **liquid** || **solid** ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Density || very small || large || large ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Thermal expansion coefficient || large (= //R/P//) || small || small ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Cohesiveness || nil || small || large ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Surface tension || nil || medium || very large ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Viscosity || small || medium || very large ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Kinetic energy per molecule || large || small || smaller ||
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/tinyinfo.gif width="10" height="10"]] Disorder || random || medium || small ||

Condensed states of matter
Even the most casual inspection of the above table shows that solids and liquids possess an important commonality that distinguishes them from gases: in solids and liquids, the molecules are in contact with their neighbors. As a consequence, these condensed states of matter generally possess much higher //densities// than gases.

Equations of state
In our study of gases, we showed that the macroscopic properties of a gas (the pressure, volume, and temperature) are related through an equation of state, and that for the limiting case of an ideal gas, this equation of state can be derived from the relatively small set of assumptions of the kinetic molecular theory. To the extent that a volume of gas consists mostly of empty space, all gases have very similar properties. Equations of state work for gases because gases consist mostly of empty space, so intermolecular interactions can be largely neglected. In condensed matter, these interactions dominate, and they tend to be unique to each particular substance, so there is no such thing as a genrally useful equation of state of liquids and solids. Is there a somewhat more elaborate theory that can predict the behavior of the other two principal states of matter, liquids and solids? Very simply, the answer is "no"; despite much effort, no one has yet been able to derive a general equation of state for condensed states of matter. The best one can do is to construct models based on the imagined interplay of attractive and repulsive forces, and then test these models by computer simulation. Nevertheless, the very factors that would seem to make an equation of state for liquids and solids impossibly complicated also give rise to new effects that are easily observed, and which ultimately define the distinguishing characteristics of the gaseous, liquid, and solid states of matter. In this unit, we will try to learn something about these distinctions, and how they are affected by the chemical constitution of a substance.

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">2. Liquids: Crystalline solids and gases stand at the two extremes of the spectrum of perfect order and complete chaos. Liquids display elements of both qualities, and both in limited and imperfect ways. Liquids and solids share most of the properties of having their molecular units in direct contact as discussed in the previous section on condensed states of matter. At they same time, liquids, like gases, are fluids, meaning that their molecular units can move more or less independently of each other. But whereas the //volume// of a gas depends entirely on the pressure (and thus generally on the volume within which it is confined), the volume of a liquid is largely independent of the pressure. We discuss the properties of liquids in some detail in [|another lesson]. Here we offer just enough to help you see how they relate to the other major states of matter. <span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">3 SolidsOf the four ancient elements of "fire, air, earth and water", it is the many forms of solids ("earths") that we encounter in daily life and which give form, color and variety to our visual world: The solid state, being the form of any substance that prevails at lower temperatures, is one in which thermal motion plays an even smaller role than in liquids. The thermal kinetic energy that the individual molecular units do have at temperatures below their melting points allows them to oscillate around a fixed center whose location is determined by the balance between local forces of attraction and repulsion due to neighboring units, but only very rarely will a molecule jump out of the fixed space alloted to it in the lattice. Thus solids, unlike liquids, exhibit //long-range order//, //cohesiveness// and //rigidity//, and possess definite //shapes//.

Classification of solids
Most people who have lived in the world long enough to read this have already developed a rough way of categorizing sollds on the basis of macroscopic properties they can easily observe; everyone knows that a piece of metal is fundamentally different from a rock or a chunk of wood. Unfortunately, nature's ingenuity is far too versatile to fit into any simple system of classifying solids, especially those composed of more than a single chemical substance.

Classification according to bond type
The most commonly used classification is based on the kinds of forces that join the molecular units of a solid together. We can usually distinguish four major categories on the basis of properties such as general appearance, hardness, and melting point.
 * **type of solid** || **molecular units** || **dominant forces** || **typical properties** ||
 * ionic || ions || coulombic || high-melting, hard, brittle ||
 * covalent || atoms of electronegative elements || chemical bonds || non-melting (decompose), extremely hard ||
 * metallic || atoms of electropositive elements || mobile electrons || moderate-to-high melting, deformable, conductive, metallic lustre ||
 * molecular || molecules || van der Waals || low-to-moderate mp, low hardness ||

Classifications should not be taken too seriously!
It's important to understand that these four categories are in a sense idealizations that fail to reflect the diversity found in nature. The triangular diagram shown here illustrates this very nicely by displaying examples of binary compounds whose properties suggest that they fall somewhere other than at a vertex of the triangle. For an excellent discussion of diagrams of this kind with many examples, see [|this page] from Mark Leach's //Chemogenesis// site. The triangle shown above excludes what is probably the largest category: molecular solids that are bound by //van der Waals forces// (which are described in the [|next lesson]). One way of including these is to expand the triangle to a tetrahedron (the so-called //Laing tetrahedron//). Although this illustrates the concept, it is visually awkward to include many examples of the intermediate cases.

Classification by type of molecular unit
Solids, like the other states of matter, can be classified according to whether their fundamental molecular units are atoms, electrically-neutral molecules, or ions. But solids possess an additional property that gases and liquids do not: an enduring structural arrangement of their molecular units. Over-simplifying only a bit, we can draw up a rough classification of solids according to the following scheme:
 * **structure** || **atoms** || **molecules** ||
 * array of discrete units || noble gas solids, metals || molecular solids ||
 * array of linked units || metals and covalent solids || "extended molecule" compounds ||
 * disordered arrangement || alternative forms of some elements (e.g. S, Se) || [|polymers], glasses ||

Classification by dominant attractive force
Notice how the boiling points in the following selected examples reflect the major type of attractive force that binds the molecular units together. Bear in mind, however, that more than one type of attractive force can be operative in many substances. This topic is discussed in more detail [|here].
 * **substance** || **bp °C** || **molecular units** || **dominant attractive force** || **separation distance (pm)** || **attraction energy (kJ/mol)** ||
 * sodium fluoride || 990 || Na+ F– || coulombic || 18.8 || 657 ||
 * sodium hydroxide || 318 || Na+ OH– || ion-dipole || 21.4 || 90.4 ||
 * water || 100 || H2O || dipole-dipole || 23.7 || 20.2 ||
 * neon || –249 || Ne || dispersion || 33.0 || 0.26 ||

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">4 Crystalline solids In a solid comprised of identical molecular units, the most favored (lowest potential energy) locations occur at regular intervals in space. If each of these locations is actually occupied, the solid is known as a //perfect// crystal. Much more on crystals and their structures can be found in these units that deal with [|lattices and external shapes], [|cubic crystals and packing of spheres], [|ionic solids], and determination of crystal structures. What really defines a crystalline solid is that its structure is composed of repeating //unit cells// each containing a small number of molecular units bearing a fixed geometric relation to one another. The resulting long-range order defines a three-dimensional geometric framework known as a lattice. image] Geometric theory shows that only fourteen different types of lattices are possible in three dimensions, and that just six different unit cell arrangements can generate these lattices. The regularity of the external faces of crystals, which in fact correspond to lattice planes, reflects the long-range order inherent in the underlying structure. Perfection is no more attainable in a crystal than in anything else; real crystals contain //defects// of various kinds, such as lattice positions that are either vacant or occupied by impurities, or by abrupt displacements or dislocations of the lattice structure. Most pure substances, including the metallic elements, form crystalline solids. But there are some important exceptions.

Metallic solids
The details of metallic bonding are covered in [|this Chemical Bonding unit]. In metals the valence electrons are free to wander throughout the solid, instead of being localized on one atom and shared with a neighboring one. The valence electrons behave very much like a mobile fluid in which the fixed lattice of atoms is immersed. This provides the ultimate in electron sharing, and creates a very strong binding effect in solids composed of elements that have the requisite number of electrons in their valence shells. The characteristic physical properties of metals such as their ability to bend and deform without breaking, their high thermal and electrical conductivities and their metallic sheen are all due to the fluid-like behavior of the valence electrons.

Molecular solids
Recall that a "molecule" is defined as a discrete aggregate of atoms bound together sufficiently tightly (that is, by //directed covalent// forces) to allow it to retain its individuality when the substance is dissolved, melted, or vaporized. The two words italicized in the preceding sentence are important; //covalent bonding// implies that the forces acting between atoms //within// the molecule are much stronger than those acting //between// molecules, and the //directional property// of covalent bonding confers on each molecule a distinctive shape which affects a number of its properties. Most compounds of carbon — and therefore, most chemical substances, fall into this category. Many simpler compounds also form molecules; H2O, NH3, CO2, and PCl5 are familiar examples. Some of the elements, such as H2, O2, O3, P4 and S8 also occur as discrete molecules. Liquids and solids that are composed of molecules are held together by //van der Waals forces//, and many of their properties reflect this weak binding. Thus molecular solids tend to be soft or deformable, have low melting points, and are often sufficiently volatile to evaporate (sublime) directly into the gas phase; the latter property often gives such solids a distinctive odor.

**Iodine** is a good example of a volatile molecular crystal. The solid (mp 114° C, bp 184°) consists of I2 molecules bound together only by dispersion forces. If you have ever worked with solid iodine in the laboratory, you will probably recall the smell and sight of its purple vapor which is easily seen in a closed container.

Because dispersion forces and the other van der Waals forces increase with the number of atoms, larger molecules are generally less volatile, and have higher melting points, than do the smaller ones. Also, as one moves down a column in the periodic table, the outer electrons are more loosely bound to the nucleus, increasing the polarisability of the atom and thus its susceptibility to van der Waals-type interactions. This effect is particularly apparent in the progression of the boiling points of the successively heavier noble gas elements.

Covalent solids
These are a class of extended-lattice compounds (see Section 6 below) in which each atom is covalently bonded to its nearest neighbors. This means that the entire crystal is in effect one super-giant “molecule”. The extraordinarily strong binding forces that join all adjacent atoms account for the extreme hardness of such substances; these solids cannot be broken or abraded without cleaving a large number of covalent chemical bonds. Similarly, a covalent solid cannot “melt” in the usual sense, since the entire crystal is its own giant molecule. When heated to very high temperatures, these solids usually decompose into their elements. ↑ diamond lattice structure source] The Hard Materials Web site has a lot of interesting information on some of the substances described in this section.

Diamond
Diamond is the hardest material known, defining the upper end of the 1-10 scale known as Moh's hardness. Diamond cannot be melted; above 1700°C it is converted to graphite, the more stable form of carbon. The diamond unit cell is face-centered cubic and contains 8 carbon atoms. The four darkly shaded ones are contained within the cell and are completely bonded to other members of the cell. The other carbon atoms (6 in faces and 4 at corners) have some bonds that extend to atoms in other cells. (Two of the carbons nearest the viewer are shown as open circles in order to more clearly reveal the bonding arrangement.) [|Wikipedia] has a very good article on the the occurrence, properties, and uses of diamonds.

Other covalent solids
Boron nitride BN is similar to carbon in that it exists as a diamond-like cubic polymorph as well as in a hexagonal form analogous to graphite. Cubic BN is the second hardest material after diamond, and finds use in industrial abrasives and cutting tools. Recent interest in BN has centered on its carbon-like ability to form nanotubes and related nanostructures. ([|A-Zmaterials article], [|Nanotechnology article], [|Wikipedia article]) Silicon carbide is an extremely rare mineral on the earth, and comes mostly from meteorites which are believed to have their origins in carbonaceous stars. The first synthetic SiC was made accidently by E.G. Acheson in 1891 who immediately recognized its industrial prospects and founded the Carborundum Co. Silicon carbide SiC is also known as //carborundum//. Its structure is very much like that of diamond with every other carbon replaced by silicon. On heating at atmospheric pressure, it decomposes at 2700°C, but has never been observed to melt. Structurally, it is very complex; at least 70 crystalline forms have been identified. Its extreme hardness and ease of synthesis have led to a diversity of applications — in cutting tools and abrasives, high-temperature semiconductors, and other high-temperature applications, manufacture of specialty steels, jewelry, and many more. For much more about the history, properties and uses of this versatile material, see this [|Wikipedia article]. Tungsten carbide WC is probably the most widely-encountered covalent solid owing to its use in "carbide" cutting tools and as the material used to make the rotating balls in [|ball-point pens]. It's high-melting (2870°C) form has a structure similar to that of diamond and is only slightly less hard. In many of its applications it is embedded in a softer matrix of cobalt or coated with titanium compounds. ([|Wikipedia article])

 <span style="color: #ff0000; font-family: Tahoma,Geneva,sans-serif; font-size: 130%;">5 Amorphous solids In some solids there is so little long-range order that the substance cannot be considered crystalline at all; such a solid is said to be //amorphous//. Amorphous solids possess short-range order but are devoid of any organized structure over longer distances; in this respect they resemble liquids. However, their rigidity and cohesiveness allow them to retain a definite shape, so for most practical purposes they can be considered to be solids.

Glasses
This term refers generally to solids formed from their melts that do not return to their crystalline forms on cooling, but instead form hard, and often transparent amorphous solids. Although some organic substances such as sugar can form glasses ("rock candy"), the term more commonly describes inorganic compounds, especially those based on //silica//, SiO2. Natural silica-based glasses, known as obsidian, are formed when certain volcanic magmas cool rapidly. The difference between crystalline silica and silica glass is shown in these simplified two-dimensional projections. It is readily apparent that much of the ordered arrangement of crystalline silica is lost in the glassy form. But the glass retains enough Si–O bonds to form a hard, rigid material. Ordinary glass is composed mostly of SiO2, which usually exists in nature in a crystalline form known as //quartz//. If quartz (in the form of sand) is melted and allowed to cool, it becomes so viscous that the molecules are unable to move to the low potential energy positions they would occupy in the crystal lattice, so that the disorder present in the liquid gets “frozen into” the solid. In a sense, glass can be regarded as a supercooled liquid. Glasses are transparent because the distances over which disorder appears are small compared to the wavelength of visible light, so there is nothing to scatter the light and produce cloudiness. Some excellent references on glass and its composition: [|Wikipedia article] [|Answers.com article] Ordinary glass is made by melting silica sand to which has been added some calcium and sodium carbonates. These additives reduce the melting point and make it more difficult for the SiO2 molecules to arrange themselves into crystalline order as the mixture cools. [|See here] for a brief discussion of this important topic. [image from Wikimedia Commons]

Glass is believed to have first been made in the Middle East at least as early as 3000 BCE. Its workability and ease of coloring has made it one of mankind's most important and versatile materials. But even after all these years, "The Nature of Glass Remains Anything but Clear", according to this interesting 2008 //NY Times// article. [images: left, right]

5 Types of molecular units

Molecules
Molecules, not surprisingly, are the most common building blocks of pure substances. Most of the 15-million-plus chemical substances presently known exist as distinct molecules. Chemists commonly divide molecular compounds into "small" and "large-molecule" types, the latter usually falling into the class of //polymers// (see below.) The dividing line between the two categories is not very well defined, and tends to be based more on the properties of the substance and how it is isolated and purified.

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">Atoms
We usually think of atoms as the building blocks of molecules, so the only pure substances that consist of plain atoms are those of some of the elements — mostly the metallic elements, and also the noble-gas elements. The latter do form liquids and crystalline solids, but only at very low temperatures. Although the **metallic elements** form crystalline solids that are essentially atomic in nature, the special properties that give rise to their "metallic" nature puts them into a category of their own. Most of the non-metallic elements exist under ordinary conditions as small molecules such as O2 or S6, or as extended structures that can have a somewhat polymeric nature. Many of these elements can form more than one kind of structure, each one stable under different ranges of temperature and pressure. Multiple structures of the same element are known as //allotropes//, although the more general term //polymorph// is now preferred.

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">Ions
Ions, you will recall, are atoms or molecules that have one or more electrons missing (positive ions) or in excess (negative ions), and therefor possess an electric charge. A basic law of nature, the //electroneutrality principle//, states that bulk matter cannot acquire more than a trifling (and chemically insignificant) net electric charge. So one important thing to know about ions is that in ordinary matter, whether in the solid, liquid, or gaseous state, any positive ions must be accompanied by a compensating number of negative ions.

Ionic solids
Ionic solids are covered in much more detail in [|this section]. Ionic substances such as sodium chloride form crystalline solids that can be regarded as made of ions. These solids tend to be quite hard and have high melting points, reflecting the strong forces between oppositely-charged ions. Solid metal oxides, such as CaO and MgO which are composed of doubly-charged ions don't melt at all, but simply dissociate into the elements at very high temperatures.

Please see [|this section] for more on these unusual states of matter.

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">Polymers
[|See here] for a more detailed discussion of polymers. Plastics and natural materials such as rubber or cellulose are composed of very large molecules called //polymers//; many important biomolecules are also polymeric in nature. Owing to their great length, these molecules tend to become entangled in the liquid state, and are unable to separate to form a crystal lattice on cooling. In general, it is very difficult to get such substances to form anything other than amorphous solids. 6 Extended solids Many compounds for which we write very simple formulas such as NiBr2 actually exist in their solid forms as linked assemblies of these basic units arranged in chains or layers that extend indefinitely in one, two, or three dimensions. Thus the very simple models of chemical bonding that apply to the isolated molecules in gaseous form must be modified to account for bonding in some of these solids. The terms "one-dimensional" and "two-dimensional", commonly employed in this context, should more accurately be prefixed by "//quasi-//"; after all, even a single atom occupies three-dimensional space!

One-dimensional solids
Atoms of some elements such as sulfur and selenium can bond together in long chains of indefinite length, thus forming polymeric, amorphous solids. The most well known of these is the amorphous "plastic sulfur" formed when molten sulfur is cooled rapidly by pouring it into water.These are never the most common (or stable) forms of these elements, which prefer to form discrete molecules. Rubber-like strands of plastic sulfur formed by pouring hot molten sulfur into cold water. After a few days, it will revert to ordinary crystalline sulfur. || image || But small molecules can also form extended chains. Sulfur trioxide is a gas above room temperature, but when it freezes at 17°C the solid forms long chains in which each S atom is coordinated to four oxygen atoms.
 * [[image:http://www.chem1.com/acad/webtext/states/state-images/sulfur_chain.png width="296" height="71" caption="plastic sulfur chain"]]

<span style="color: #ff0000; display: block; font-family: Tahoma,Geneva,sans-serif; font-size: 130%; text-align: left;">Multi-dimensional solids
C//ovalent solids//, which also fall into this category,are discussed in Section 4 above. Many inorganic substances form crystalline solids which are built up from parallel chains in which the basic formula units are linked by weak bonds involving dipole-dipole and dipole-induced dipole interactions. Neighborisg chains are bound mainly by dispersion forces. Solid copper (II) chloride consists of multiple CuCl2 units joined into long chains by covalent bonding. The chains are held together laterally by weaker van der Waals forces (mainly dispersion and ion-induced dipole. Each copper atom resides at the center of an imaginary octahedron whose vertices are defined by chlorine atoms in adjacent chains.

Layer or sheet-like structures
Solid cadmium chloride is a good example of a //layer// structure. The Cd and Cl atoms occupy separate layers; each of these layers extends out in a third dimension to form a //sheet//. The CdCl2 crystal is built up from stacks of these layers held together by van der Waals forces. Each Cd atom is covalently bonded to two chlorine atoms in adjacent layers; all other "bonds" are weak van der Waals attractions. It's worth pointing out that although salts such as CuCl2 and CdCl2 are dissociated into ions when in aqueous solution, the solids themselves should not be regarded as "ionic solids". See also this section of the lesson on ionic solids.

Graphite
Graphite is a //polymorph// of carbon and its most stable form. It consists of sheets of fused benzene rings stacked in layers. The spacing between layers is sufficient to admit molecules of water vapor and other atmospheric gases which become absorbed in the interlamellar spaces and act as lubricants, allowing the layers to slip along each other. Thus graphite itself often has a flake-like character and is commonly used as a solid lubricant, although it loses this property in a vacuum. As would be expected from its //anisotropic// structure, the electric and thermal conductivity of graphite are much greater in directions parallel to the layers than across the layers. The melting point of 4700-5000°C makes graphite useful as a high-temperature refractory material. Graphite is the most common form of relatively pure carbon found in nature. Its name comes from the same root as the Greek word for "write" or "draw", reflecting its use as pencil "lead" since the 16th century. (The misnomer, which survives in common use, is due to its mis-identification as an ore of the metallic element of the same name at a time long before modern chemistry had developed.) Some interesting Wikipedia pages: [|Graphite]; [|Pencils, their history and manufacture]

Graphene
//Graphene// is a two-dimensional material consisting of a single layer of graphite — essentially "chicken wire made of carbon" that was discovered in 2004. Small fragments of graphene can be obtained by several methods; one is to attach a piece of Scotch Tape™ to a piece of graphite and then carefully pull it off (a process known as //exfoliation//.) Fragments of graphene are probably produced whenever one writes with a pencil. Graphene has properties that are uniquely different from all other solids. It is the strongest known material, and it exhibits extremely high electrical conductivity due to its massless electrons which are apparently able to travel at relativistic velocities through the layer. Some excellent articles on graphene: //[|Scientific American]// (April 2007) - [|Wikipedia] - [|Lawrence Berkeley Laboratory] What you should be able to do Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the green-highlighted terms in the context of this topic. concept map taken from:http://www.chem1.com/acad/webtext/states/states.html
 * State the major feature that characterizes a condensed state of matter.
 * Describe some of the **major observable properties** that distinguish gases, liquid and solids, and state their relative magnitudes in these three states of matter.
 * Describe the **dominant forces** and the resulting **physical properties** that distinguish ionic, covalent, metallic, and molecular solids.
 * Explain the difference between crystalline and amorphous solids, and cite some examples of each.
 * Name some of the basic **molecular units** from which solids of different type can be composed.
 * What is meant by an "extended" or "infinite-molecule solid"?
 * Describe some of the special properties of //graphite// and their structural basis.