Intermolecular++Forces

=Intermolecular Forces=

Definition
Intermolecular forces are the forces of attraction between molecules. They vary in strength but are generally wearker than the bonds that join atoms in molecules or metal atoms in solid metals. Since intermolecular forces are weaker than these other bonds, molecules tend to have lower boiling points than metals. There are four distinct types of intermolecular forces, all with distinct strengths.

Ion-Ion Forces:
The strongest intermolecular force exists between two ionic compounds. The positive end of an ionic compound will be attracted to the negative end in another, sticking together and forming a crystaline structure. Because these bonds are strong and always come together in a predictable pattern, these substances tend to be rather hard and have a high boiling point.

Dipole-Dipole Forces:
A very strong intermolecular force exists between two polar molecules. Because of their uneven charge distribution, polar molecules can act as dipoles. All of the electrons in a polar molecule will tend to flow in a certain direction, creating the dipole. The negative partial charge in one molecule will be attracted to the positive partial charge in another, much like in ions. The molecular polarity of a molecule depends on both the polarity and orientation of the atoms that make up the molecule.

Hydrogen Bonding
Also quite strong are hydrogen bonds, which are made very polar by the large difference in electronegativity between hydrogen and highly electronegative atoms like flourine, oxygen, or nitrogen. The hydrogen atom bonds to an unshared pair of electrons in the electronegative atom to form a hydrogen bond. Water's relatively high boiling point can be attributed to its hydrogen bonds.

London Dispersion Forces
The weakest intermolecular forces are between nonpolar molecules. In an atom or molecule, the electrons are in continuous motion. As a result, at any instant, the electron distribution may be slightly uneven, creating very weak dipoles. The bonds created between these dipoles are known as London dispersion forces (named after Fritz London), and exist between any and every molecule, but are the only force acting upon noble-gas atoms, nonpolar molecules, and slightly polar molecules. These substances all have low boiling points, reflecting the weakness of their intermolecular forces.

=**Physical Properties of Liquids & Solids**=

The phase of a substance is determined by the intermolecular forces that attract one molecule to another. At standard temperature and pressure, a substance that is a gas will have relatively weak attraction between each molecule in the substance, a liquid has a stronger attraction between each molecule in the substance, and a solid has the strongest attraction between each molecule in the substance. The varying intermolecular forces have strong impacts on the properties of liquids solids and gases.

Liquids
Liquids are fluid substances that take the shape of their containers. All liquids have a certain resistance to flowing. This resistance is known as viscosity. A liquid's viscosity is dependent on a number of factors, such as temperature, pressure, and for some liquids, time. In general, the stronger the intermolecular force, the higher the viscosity of the liquid. Liquids are nearly incompressible. Two more physical properties of liquids are surface tension and capillary action.

__Surface Tension__:
Surface tension is the attraction of molecules of a liquid towards the surface of that liquid. This reduces the surface area to the smallest possible size. It is caused by the equal strength of the intermolecular force in a liquid. In the liquid, all molecules are being pulled from all directions by all other molecules. However, at the surface, the molecules are being pulled by the rest of the liquid more than the air. This causes an inward force. Since liquids have a high resistance to compression, an almost solid surface forms. Surface tension is the reason why jumping off a high bridge is just as bad as jumping off a building. If you don't consider the cleanup, that is. __Capillary action__: Capillary action is the tendency of the surface of a liquid to "climb up" its solid container. The surface of a liquid is attracted to the surface of the solid and thus the liquid clings to the side of the walls. For the same reason surface tension occurs, the attractive force pulls more molecules of the liquid from the surface and up the walls of the container. Capillary action is what causes a meniscus to occur. It is also explains why if you submerge part of a piece of paper, the water seems to climb up the rest.

Solids
Solids are rigid substances with definite shape. They have a comparatively high density and are generally incompressible. There are many different structures a solid can take, determined by the intermolecular forces of the substance.

__Ionic Crystals__:
Ionic crystals are crystals of ions bound together by electrostatic or ion-ion attraction. Examples would be NaCl or Ki

__Covalent Network Crystals__:
Covalent network crystals are giant molecular lattice structures held together by covalent bonding, NOT BY INTERMOLECULAR FORCES. Examples are diamond and silicon oxide. All covalent network crystals form giant networks because no matter how many molecules are used, there will always be unused bonding sites on the outside of the lattice.

__Metallic Crystals:__ Metallic crystals are pure metals such as Hg or Fe which form into crystalline structures with free electrons shared throughout the structure as a whole. This creates a surplus of electrons. These free electrons are the reason metals are such good conductors of heat and electricity. The molecules are attracted to each other by a force know as metallic bonding, which is vaguely similar to ionic bonding.

__Covalent Molecular Crystal:__
Covalent molecular crystals are covalently bond molecules held together through any one of the intermolecular forces depending on the substance. For example, H2O is held together with hydrogen bonding, C6H6 is held together with dipole-dipole attraction,

CH4 is held together with London dispersion force, and they are all covalent molecular crystals.
=Heating Curves=

There are two diagrams that are important. Phase change diagrams and phase diagrams.

Phase Change diagrams
Phase change diagrams present a graphical representation of the phase of a substance based on temperature over time, with heat being constantly added. Here is a phase change diagram of water.



As you can see it takes heat to raise the temperature of ice to the melting point. At the melting point it takes heat for the H2O to change state. The heat required to change ice at 0 degree Celsius into water at 0 degree Celsius is called the heat of fusion and is 79.7 calories per gram. Then it takes heat to raise the temperature of water to boiling point. At boiling point it takes heat for the water to change state. The heat required to change 100 degree Celsius water into 100 degree Celsius steam is called the heat of vaporization and is 539 calories per gram. Then it takes heat to raise the temperature of steam.

Phase Diagrams:
Phase diagrams are graphical representations of a substance's state based on pressure over temperature. Water has a unique phase diagram. All other phase diagrams are generally the same, but with different values. By reading a phase diagram you can find several distinct qualities of a substance. There are a few crucial points on every phase diagram. These are the boiling point of the substance at 1 ATM, the triple point, and the critical point. The boiling point at 1 ATM shows you the substance's boiling point in normal circumstances. The triple point shows you the necessary pressure and temperature where the substance will exist in a constant trade off of all three states of matter. The Critical point shows you the temperature where at any temperature greater than or equal to that temperature, no matter how much the pressure is increased, the substance will condensate.

__Phase Diagram (of CO2):__
Below is the phase diagram of carbon dioxide.



As you can see, the normal boiling point is -78.5 degrees Celsius, the triple point is at 5.11 ATM and -56.4 degrees Celsius, and the critical point is at 73 ATM and 31.1 degrees Celsius. Also, you can plainly see that there are two ways that a change of state can occur: through increase or decrease in temperature, or by increase or decrease in pressure. All phase diagrams, except for water, have the same basic shape as above.

__Phase Diagram Of H2O:__
Below is the phase diagram of water.



As you can see water has one striking difference from carbon dioxide or any other substance. It can change from liquid to solid to gas by decreasing the pressure at an appropriate temperature. this quality is unique to water.

The Kinetic-Molecular Theory of Gases

The kinetic-molecular theory of gases is a branch of the kinetic-molecular theory, which, in short, states that all particles of matter are in motion at all times. The kinetic-molecular theory of gases accounts for this principle, but also more specifically outlines the theory as it pertains to gases and provides the requirements for an “ideal gas.” An ideal gas is one which fits each requirement of the kinetic-molecular theory of gases, but can technically never exist. The kinetic-molecular theory of gases is based off five assumptions, which are listed below.

1. Gases are made up of minute particles that are far apart from one another. -Molecules in gases are farther apart than those in liquids or solids, which is why gases are so much less dense than liquids and solids, and are much more easily compressed.

2. Collisions between gas particles and between particles and container walls are elastic collisions (an elastic collision has no net loss of kinetic energy). -Kinetic energy may be transferred, but as long as the temperature is constant the net kinetic energy will remain the same.

3. Gas particles are constantly in motion, so they constantly possess the energy of motion, kinetic energy. -The immense amount of kinetic energy overcomes any attractive forces in molecules.

4. There are no forces of attraction or repulsion in gas molecules. -Because kinetic energy overrides attraction and repulsion in gas molecules, molecules collide often but immediately bounce apart.

5. The average kinetic energy of gas molecules depends on the temperature of the gas. -All gases at the same temperature have the same kinetic energy. -The kinetic energy of particles increases and decreases as temperature increases and decreases. -The Kinetic energy of any moving object or particle is given by the equation KE = (1/2) mv^2, where KE is kinetic energy, m is the mass of the object/particle, and v is its velocity.

Many gases with low pressure and high temperature come close to fulfilling the requirements of becoming an Ideal Gas, and thus the necessities of the kinetic-molecular theory of gases. The physical properties of such gases are made possible by kinetic-molecular theory of gases.

-Points three and four allow expansion to occur when gas is transferred from one container to another. The gas will expand or compress to fit the shape of any container because under the kinetic-molecular theory of gases, gas molecules are constantly in motion and moving in every direction, and have no forces of attraction towards one another. Thus, they will move to fill up the space and will not be drawn together in the middle.

-Gas possesses Fluidity, just as liquid does, because due to condition four, gas molecules are not attracted to or repelled from one and other, and thus move with ease.

-Gases are also low in density, because condition one states that gas molecules are far apart from each other.

-Because condition one states that gas molecules are far apart from one and other, gases can be easily compressed since there is much room for molecules to condense into.

-Gases are easily combined, or “diffused,” because gas particles are constantly in motion and will move every which way (condition three).

Phase Diagrams A phase diagram is a way of determining the phase of a substance (ie: gas, liquid, or solid) at any given pressure and temperature.



Above is the phase diagram for a typical pure substance. As is shown, the temperature of the substance is on the x axis while its pressure is on the y axis. If the point that gives the temperature and pressure of the substance land in the solid region, then, understandably, the substance is in its solid phase (this is true for the liquid and gas regions as well).

If the substance falls on the line between the solid and gas areas it is simply in equilibrium between the solid and gas states. In addition, if the substance that falls on the line between the solid and liquid areas it is in equilibrium between the solid and liquid stages, and if the substance falls on the line between the liquid and gas areas, it is in equilibrium between the liquid and gas stages. If the substance were to fall on the “triple point” shown in the above diagram, it would be in equilibrium between all three states. Phase Diagrams can be used to find the normal melting and boiling points of a substance. Since the boiling and melting points both occur when the pressure is equal to 1atm, and they both indicate a change in phase, they can be found by calculating the intersection of the line P=1atm and the lines indicating phase changes (see figure below).



= =

=PRESSURE= = = Pressure is defined as the force per unit area on a surface. The same amount of force will cause varying pressure depending on the area of the surface upon which it is exerted. For example, if you blow air into a glass of liquid through a small diameter straw, more pressure will be created than if you blow through a larger diameter straw. This is because the molecules of air that are being blown into the smaller straw have less space to travel, thereby creating more collisions of air molecules within the straw; this causes the molecules to push outwards against the sides of the straw, creating greater force, and thereby more pressure than in the larger straw, where the air molecules have more room to move, resulting in less collisions, less force, and therefore less pressure.

Pressure is determined by the exertion created on a surface by the flow of mass from an area of high pressure to one of low pressure. An example of this is that, when you blow air into a balloon, the balloon expands, because the pressure of air molecules that are being blown to the inside of the balloon is greater than the air pressure outside the balloon. If the neck of the balloon is released, the air flows out of the balloon, rather than more air flowing in, because the air moves from a region of high pressure (inside the balloon) to a region of low pressure (outside the balloon).



 When pressure is referred to in conjunction with temperature and intermolecular forces, we are usually referring to atmospheric pressure, or the force of air on another substance. The normal boiling point of an element refers to its boiling point at the atmospheric pressure at sea level (one atm); this is equal to about 14.7 lbs per square inch. In comparison, the air pressure in a car’s tire should be between 26 to thirty lbs per square inch, and bicycle tires should be 40 to 60 lbs per square inch.

 There are several laws governing pressure. One of these is that pressure is affected by temperature; as the temperature of a gas rises, the pressure of the gas will also rise because the molecules will gain more energy and create more collisions. This is known as Le Chatelier’s Principle. Gas pressure is also affected by atmosphere. As one travels to higher atmospheres, pressure decreases because there are less air molecules at higher elevations. Boyles Law, discovered in 1662, states that, at a constant temperature, the volume of a gas varies inversely with the pressure exerted on it.

 Atmospheric pressure is measured with a barometer; this is also referred to as barometric pressure. Specifically, a Torricelli barometer is used; it consists of a long tube closed at one end, filled with mercury and inverted in a container of mercury. At sea level, one atm will cause the mercury in the tube to reach a standard height of 760 mm and different atmospheric pressures can be measured against this standard.



Fun Pressure experiments that you can do on your own:

[|Antigravity[[http://www.elmhurst.edu/~chm/demos/][|Battle of Two Balloons]

[|Magic Leaky Bottle]

**BOILING POINT and VAPOR PRESSURE**

Technically speaking, the //boiling point// of a substance is the temperature at which the liquid and the vapor (gas) phases of a substance can exist in equilibrium.Put more simply, //boiling point// is the temperature at which a substance changes its state from liquid to a vapor (gas), both within the liquid and at its surface.This occurs because, at a constant pressure, a certain amount of heat applied to a liquid will increase the energy of the liquid’s molecules enough to break the intermolecular forces that bind them together in their liquid state.As the liquid reaches the necessary temperature and the molecules obtain enough additional energy to break these bonds, the process known as //vaporization// occurs, allowing the molecules to escape the liquid form as individual molecules of vapor.//Vaporization// at the boiling point is simply another term for //boiling//.


 * [[image:http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch14/graphics/14_7fig.gif width="349" height="356" align="center" caption="graph"]] ||
 * Liquids boil when their vapor pressure is equal to the pressure exerted on the liquid by its surroundings. ||

When vaporization occurs, the pressure of the vapor above the liquid from which molecules are escaping is in equilibrium with the pressure exerted by its liquid. The pressure exerted by the gas at the instant the liquid reaches its boiling point is called the vapor pressure. The stronger the bonds are between the molecules of the liquid, the lower the vapor pressure will be in the vapor/gas that is released at the boiling point. For example, water molecules have strong bonds and relatively low vapor pressure. If there is a weak bond between the liquid molecules, it is easier for the molecules to evaporate, thereby creating a higher vapor pressure.


 * When a solid or a liquid evaporates to a gas in a closed container, the molecules cannot escape.
 * Some of the gas molecules will eventually strike the condensed phase and condense back into it.
 * When the rate of condensation of the gas becomes equal to the rate of evaporation of the liquid or solid, the amount of gas, liquid and/or solid no longer changes.
 * The gas in the container is in **//equilibrium//** with the liquid or solid.
 * [[image:http://www.chem.purdue.edu/gchelp/liquids/liqint.gif height="179" align="center"]] || [[image:http://www.chem.purdue.edu/gchelp/liquids/solidint.gif height="179" align="center"]] ||
 * Microscopic equilibrium between gas and liquid. Note that the rate of evaporation of the liquid is equal to the rate of condensation of the gas. || Microscopic equilibrium between gas and solid. Note that the rate of evaporation of the solid is equal to the rate of condensation of the gas. ||
 * The pressure exerted by the gas in equilibrium with a solid or liquid in a closed container at a given temperature is called the vapor pressure. (http://www.chem.purdue.edu/gchelp/liquids/vpress.html)

Vapor pressure is also affected by the presence of dissolved substances in the liquid. According to Raoult's Law, the vapor pressure of a pure liquid will be reduced by the addition of a solute. Temperature also affects the vapor pressure of a liquid. As the temperature of a liquid increases, its vapor pressure also increases; vapor pressure decreases as the temperature decreases.



Once the temperature of a liquid has reached its boiling point, the temperature does not continue to increase; rather, it stays constant at that boiling point temperature until all of the liquid has been converted to vapor.Since the intermolecular forces bonding the molecules of each element differ, each element has a different boiling point. You can use click on this link for an automated program that determines boiling points of different molecules: http://pirika.com/chem/TCPEE/BP/ourBP.htm.

A comprehensive chart of each element’s normal boiling point follows:

Periodic Table of Elements Sorted by Boiling Point ||< He ||< 2 || ||< H ||< 1 || ||< Ne ||< 10 || ||< N ||< 7 || ||< F ||< 9 || ||< Ar ||< 18 || ||< O ||< 8 || ||< Kr ||< 36 || ||< Xe ||< 54 || ||< Rn ||< 86 || ||< Cl ||< 17 || ||< Br ||< 35 || ||< I ||< 53 || ||< P ||< 15 || ||< At ||< 85 || ||< Hg ||< 80 || ||< S ||< 16 || ||< As ||< 33 || ||< Cs ||< 55 || ||< Fr ||< 87 || ||< Se ||< 34 || ||< Rb ||< 37 || ||< K ||< 19 || ||< Cd ||< 48 || ||< Na ||< 11 || ||< Zn ||< 30 || ||< Po ||< 84 || ||< Te ||< 52 || ||< Mg ||< 12 || ||< Yb ||< 70 || ||< Li ||< 3 || ||< Sr ||< 38 || ||< Tl ||< 81 || ||< Ca ||< 20 || ||< Ra ||< 88 || ||< Bi ||< 83 || ||< Sb ||< 51 || ||< Eu ||< 63 || ||< Pb ||< 82 || ||< Sm ||< 62 || ||< Ba ||< 56 || ||< Tm ||< 69 || ||< Mn ||< 25 || ||< In ||< 49 || ||< Ag ||< 47 || ||< Sn ||< 50 || ||< Si ||< 14 || ||< Ga ||< 31 || ||< Al ||< 13 || ||< Dy ||< 66 || ||< Cu ||< 29 || ||< Am ||< 95 || ||< Cr ||< 24 || ||< Ho ||< 67 || ||< Ni ||< 28 || ||< Fe ||< 26 || ||< Au ||< 79 || ||< Ge ||< 32 || ||< Sc ||< 21 || ||< Er ||< 68 || ||< Co ||< 27 || ||< Pd ||< 46 || ||< Be ||< 4 || ||< Tb ||< 65 || ||< Nd ||< 60 || ||< Cm ||< 96 || ||< Ac ||< 89 || ||< Pu ||< 94 || ||< Gd ||< 64 || ||< Ti ||< 22 || ||< Y ||< 39 || ||< Lu ||< 71 || ||< V ||< 23 || ||< Ce ||< 58 || ||< La ||< 57 || ||< Pm ||< 61 || ||< Pr ||< 59 || ||< Rh ||< 45 || ||< Pt ||< 78 || ||< Ru ||< 44 || ||< Np ||< 93 || ||< B ||< 5 || ||< Pa ||< 91 || ||< U ||< 92 || ||< Zr ||< 40 || ||< Ir ||< 77 || ||< Hf ||< 72 || ||< Mo ||< 42 || ||< Nb ||< 41 || ||< Th ||< 90 || ||< C ||< 6 || ||< Tc ||< 43 || ||< Os ||< 76 || ||< Ta ||< 73 || ||< Rf ||< 104 || ||< Re ||< 75 || ||< W ||< 74 ||
 * ~ Boiling Point ||~ Name ||~ Sym ||~ # ||
 * < 4.365K ||< -268.785°C ||< -451.813°F ||< [|Helium]
 * < 20.418K ||< -252.732°C ||< -422.918°F ||< [|Hydrogen]
 * < 27.246K ||< -245.904°C ||< -410.6°F ||< [|Neon]
 * < 77.5K ||< -195.65°C ||< -320.17°F ||< [|Nitrogen]
 * < 85.1K ||< -188.05°C ||< -306.49°F ||< [|Fluorine]
 * < 87.45K ||< -185.7°C ||< -302.3°F ||< [|Argon]
 * < 90.33K ||< -182.82°C ||< -297.08°F ||< [|Oxygen]
 * < 119.95K ||< -153.2°C ||< -243.8°F ||< [|Krypton]
 * < 165.18K ||< -107.97°C ||< -162°F ||< [|Xenon]
 * < 211K ||< -62°C ||< -80°F ||< [|Radon]
 * < 239.25K ||< -33.9°C ||< -29°F ||< [|Chlorine]
 * < 332.4K ||< 59.25°C ||< 138.65°F ||< [|Bromine]
 * < 458.55K ||< 185.4°C ||< 365.7°F ||< [|Iodine]
 * < 553K ||< 280°C ||< 536°F ||< [|Phosphorus]
 * < 610K ||< 337°C ||< 639°F ||< [|Astatine]
 * < 630K ||< 357°C ||< 675°F ||< [|Mercury]
 * < 717.9K ||< 444.75°C ||< 832.55°F ||< [|Sulfur]
 * < 876K ||< 603°C ||< 1117°F ||< [|Arsenic]
 * < 944K ||< 671°C ||< 1240°F ||< [|Cesium]
 * < 950K ||< 677°C ||< 1251°F ||< [|Francium]
 * < 958K ||< 685°C ||< 1265°F ||< [|Selenium]
 * < 961K ||< 688°C ||< 1270°F ||< [|Rubidium]
 * < 1032K ||< 759°C ||< 1398°F ||< [|Potassium]
 * < 1038K ||< 765°C ||< 1409°F ||< [|Cadmium]
 * < 1156K ||< 883°C ||< 1621°F ||< [|Sodium]
 * < 1180K ||< 907°C ||< 1665°F ||< [|Zinc]
 * < 1235K ||< 962°C ||< 1764°F ||< [|Polonium]
 * < 1261K ||< 988°C ||< 1810°F ||< [|Tellurium]
 * < 1363K ||< 1090°C ||< 1994°F ||< [|Magnesium]
 * < 1467K ||< 1194°C ||< 2181°F ||< [|Ytterbium]
 * < 1615.15K ||< 1342°C ||< 2448°F ||< [|Lithium]
 * < 1657K ||< 1384°C ||< 2523°F ||< [|Strontium]
 * < 1746K ||< 1473°C ||< 2683°F ||< [|Thallium]
 * < 1757K ||< 1484°C ||< 2703°F ||< [|Calcium]
 * < 1809K ||< 1536°C ||< 2797°F ||< [|Radium]
 * < 1837K ||< 1564°C ||< 2847°F ||< [|Bismuth]
 * < 1860K ||< 1587°C ||< 2889°F ||< [|Antimony]
 * < 1870K ||< 1597°C ||< 2907°F ||< [|Europium]
 * < 2013K ||< 1740°C ||< 3164°F ||< [|Lead]
 * < 2064K ||< 1791°C ||< 3256°F ||< [|Samarium]
 * < 2171K ||< 1898°C ||< 3448°F ||< [|Barium]
 * < 2220K ||< 1947°C ||< 3537°F ||< [|Thulium]
 * < 2235K ||< 1962°C ||< 3564°F ||< [|Manganese]
 * < 2346K ||< 2073°C ||< 3763°F ||< [|Indium]
 * < 2436K ||< 2163°C ||< 3925°F ||< [|Silver]
 * < 2543K ||< 2270°C ||< 4118°F ||< [|Tin]
 * < 2628K ||< 2355°C ||< 4271°F ||< [|Silicon]
 * < 2676K ||< 2403°C ||< 4357°F ||< [|Gallium]
 * < 2740K ||< 2467°C ||< 4473°F ||< [|Aluminum]
 * < 2835K ||< 2562°C ||< 4644°F ||< [|Dysprosium]
 * < 2840K ||< 2567°C ||< 4653°F ||< [|Copper]
 * < 2880K ||< 2607°C ||< 4725°F ||< [|Americium]
 * < 2945K ||< 2672°C ||< 4842°F ||< [|Chromium]
 * < 2968K ||< 2695°C ||< 4883°F ||< [|Holmium]
 * < 3005K ||< 2732°C ||< 4950°F ||< [|Nickel]
 * < 3023K ||< 2750°C ||< 4982°F ||< [|Iron]
 * < 3080K ||< 2807°C ||< 5085°F ||< [|Gold]
 * < 3103K ||< 2830°C ||< 5126°F ||< [|Germanium]
 * < 3104K ||< 2831°C ||< 5128°F ||< [|Scandium]
 * < 3136K ||< 2863°C ||< 5185°F ||< [|Erbium]
 * < 3143K ||< 2870°C ||< 5198°F ||< [|Cobalt]
 * < 3237K ||< 2964°C ||< 5367°F ||< [|Palladium]
 * < 3243K ||< 2970°C ||< 5378°F ||< [|Beryllium]
 * < 3296K ||< 3023°C ||< 5473°F ||< [|Terbium]
 * < 3341K ||< 3068°C ||< 5554°F ||< [|Neodymium]
 * < 3383K ||< 3110°C ||< 5630°F ||< [|Curium]
 * < 3473K ||< 3200°C ||< 5792°F ||< [|Actinium]
 * < 3503K ||< 3230°C ||< 5846°F ||< [|Plutonium]
 * < 3539K ||< 3266°C ||< 5911°F ||< [|Gadolinium]
 * < 3560K ||< 3287°C ||< 5949°F ||< [|Titanium]
 * < 3611K ||< 3338°C ||< 6040°F ||< [|Yttrium]
 * < 3668K ||< 3395°C ||< 6143°F ||< [|Lutetium]
 * < 3682K ||< 3409°C ||< 6168°F ||< [|Vanadium]
 * < 3699K ||< 3426°C ||< 6199°F ||< [|Cerium]
 * < 3730K ||< 3457°C ||< 6255°F ||< [|Lanthanum]
 * < 3785K ||< 3512°C ||< 6354°F ||< [|Promethium]
 * < 3785K ||< 3512°C ||< 6354°F ||< [|Praseodymium]
 * < 4000K ||< 3727°C ||< 6741°F ||< [|Rhodium]
 * < 4100K ||< 3827°C ||< 6921°F ||< [|Platinum]
 * < 4173K ||< 3900°C ||< 7052°F ||< [|Ruthenium]
 * < 4175K ||< 3902°C ||< 7056°F ||< [|Neptunium]
 * < 4275K ||< 4002°C ||< 7236°F ||< [|Boron]
 * < 4300K ||< 4027°C ||< 7281°F ||< [|Protactinium]
 * < 4407K ||< 4134°C ||< 7473°F ||< [|Uranium]
 * < 4650K ||< 4377°C ||< 7911°F ||< [|Zirconium]
 * < 4701K ||< 4428°C ||< 8002°F ||< [|Iridium]
 * < 4876K ||< 4603°C ||< 8317°F ||< [|Hafnium]
 * < 4885K ||< 4612°C ||< 8334°F ||< [|Molybdenum]
 * < 5017K ||< 4744°C ||< 8571°F ||< [|Niobium]
 * < 5061K ||< 4788°C ||< 8650°F ||< [|Thorium]
 * < 5100K ||< 4827°C ||< 8721°F ||< [|Carbon]
 * < 5150K ||< 4877°C ||< 8811°F ||< [|Technetium]
 * < 5285K ||< 5012°C ||< 9054°F ||< [|Osmium]
 * < 5698K ||< 5425°C ||< 9797°F ||< [|Tantalum]
 * < 5800K ||< °C ||< °F ||< [|Rutherfordium]
 * < 5900K ||< 5627°C ||< 10161°F ||< [|Rhenium]
 * < 5928K ||< 5655°C ||< 10211°F ||< [|Tungsten]

The temperature needed to reach the boiling point of any liquid is inversely proportional to the pressure at which the liquid is subject to at the time heat is applied. In most situations, pressure refers to atmospheric pressure, and “normal boiling points” refer to the boiling point at one standard atmosphere of pressure (1 atm) on the liquid. Therefore, since boiling points change inversely to the amount of pressure, it would take less heat to vaporize a liquid at a pressure greater than 1 atm, and more heat at a lower pressure. For example, the normal boiling point of water (H2O) is 1000C (2120F) at 1 atm; cooking an egg in this boiling water’s temperature will take a certain amount of time. If, however, you want to cook this egg in water that you have boiled after climbing a mountain and reaching an elevation of 10,000 feet, it will take longer for the egg to cook because water boils at only 900C at this elevation. Since the water boils at a lower temperature and, remember, it will remain at this temperature until the vaporization processed is completed (or until the heat is reduced), the water will not get hotter than 900C and the egg will cook more slowly.