What is the temperature just before boiling?

Boiling is the process by which a liquid turns into a vapor when it is heated to its boiling point. The change from a liquid phase to a gaseous phase occurs when the vapor pressure of the liquid is equal to the atmospheric pressure exerted on the liquid. Boiling is a physical change and molecules are not chemically altered during the process.

Table of Contents Show

• Contributors and Attributions
• Saturation temperature and pressure
• Relation between the normal boiling point and the vapor pressure of liquids
• Boiling point of chemical elements
• Boiling point as a reference property of a pure compound
• Impurities and mixtures
• External links

When atoms or molecules of a liquid are able to spread out enough to change from a liquid phase to a gaseous phase, bubbles form and boiling occurs.

Figure 1: Formation of bubbles in boiling water.

Video: Boiling basics (https://www.youtube.com/embed/Py0GEByCke4).

The boiling point is the temperature at which boiling occurs for a specific liquid. For example, for water, the boiling point is 100ºC at a pressure of 1 atm. The boiling point of a liquid depends on temperature, atmospheric pressure, and the vapor pressure of the liquid. When the atmospheric pressure is equal to the vapor pressure of the liquid, boiling will begin.

A liquid will begin to boil when Atmospheric Pressure = Vapor Pressure of Liquid

Exercise 1: Boiling Basics

When a liquid boils, what is inside the bubbles?

Answer

The bubbles in a boiling liquid are made up of molecules of the liquid which have gained enough energy to change to the gaseous phase.

Exercise 2

Describe the formation of bubbles in a boiling liquid (see video for answer).

When boiling occurs, the more energetic molecules change to a gas, spread out, and form bubbles. These rise to the surface and enter the atmosphere. It requires energy to change from a liquid to a gas (see enthalpy of vaporization). In addition, gas molecules leaving the liquid remove thermal energy from the liquid. Therefore the temperature of the liquid remains constant during boiling. For example, water will remain at 100ºC (at a pressure of 1 atm or 101.3 kPa) while boiling. A graph of temperature vs. time for water changing from a liquid to a gas, called a heating curve, shows a constant temperature as long as water is boiling.

Exercise 3: Heating Curve for Water

Based on the heating curve below, when will the temperature of \(H_2O\) exceed 100ºC (in an open system)?

Answer

The temperature of \(H_2O\) will only exceed 100 ºC once it has entirely changed to the gaseous phase. As long as there is liquid the temperature will remain constant.

The pressure of gas above a liquid affects the boiling point. In an open system this is called atmospheric pressure. The greater the pressure, the more energy required for liquids to boil, and the higher the boiling point.

Higher Atmospheric Pressure = More Energy Required to Boil = Higher Boiling Point

In an open system this can be visualized as air molecules colliding with the surface of the liquid and creating pressure. This pressure is transmitted throughout the liquid and makes it more difficult for bubbles to form and for boiling to take place. If the pressure is reduced, the liquid requires less energy to change to a gaseous phase, and boiling occurs at a lower temperature.

Video: Atmospheric Pressure and Boiling (www.youtube.com/watch?v=aiwy...ature=youtu.be).

Exercise 4

Based on the atmospheric pressure, predict the boiling point for water at the following locations. Remember that water boils at 100ºC at sea level on earth. Assume constant temperature.

• Earth at Sea Level: 101.3 kPa
• Mount Everest Summit: 33.7 kPa
• Mars (average): 0.6 kPa
• Venus (surface): 9200 kPa

Answer

Since water boils at 100ºC, water would boil quickly on Mars (actual value us about 10ºC). The boiling point on Mt. Everest would be closer to water (actual value about 70ºC). On Venus water would boil well over 100ºC.

The molecules leaving a liquid through evaporation create an upward pressure as they collide with air molecules. This upward push is called the vapor pressure. Different substances have different vapor pressures and therefore different boiling points. This is due to differing intermolecular forces between molecules.

Video: Vapor Pressure and Bioling (youtu.be/ffBusZO-TO0)

The vapor pressure of a liquid lowers the amount of pressure exerted on the liquid by the atmosphere. As a result, liquids with high vapor pressures have lower boiling points. Vapor pressure can be increased by heating a liquid and causing more molecules to enter the atmosphere. At the point where the vapor pressure is equal to the atmospheric pressure boiling will begin. In effect, without any external pressure the liquid molecules will be able to spread out and change from a liquid to a gaseous phase. The gas, as bubbles in the liquid, will rise to the surface and be released into the atmosphere.

Contributors and Attributions

• Wayne Breslyn, NBCT, Ph.D. (Gaithersburg High School)

• Chadwick Wyler

Boiling is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

The simple answer to this question is that the boiling point of water is 100 °C or 212 °F at 1 atmosphere of pressure (sea level).

However, the value is not a constant. The boiling point of water depends on the atmospheric pressure, which changes according to elevation. Water boils at a lower temperature as you gain altitude (e.g., going higher on a mountain), and boils at a higher temperature if you increase atmospheric pressure (coming back down to sea level or going below it).

The boiling point of water also depends on the purity of the water. Water that contains impurities (such as salted water) boils at a higher temperature than pure water. This phenomenon is called boiling point elevation, which is one of the colligative properties of matter.

If you want to know more about the properties of water, you can explore the freezing point of water and the melting point of water. You can also contrast the boiling point of water to the boiling point of milk.

• Goldberg, David E. (1988). 3,000 Solved Problems in Chemistry (1st ed.). McGraw-Hill. section 17.43, p. 321. ISBN 0-07-023684-4.
• West, J. B. (1999). "Barometric pressures on Mt. Everest: New data and physiological significance." Journal of Applied Physiology. 86 (3): 1062–6. doi:10.1152/jappl.1999.86.3.1062

Temperature at which a substance changes from liquid into vapor

This article is about the boiling point of liquids. For other uses, see Boiling point (disambiguation).

The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid[1][2] and the liquid changes into a vapor.

The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at low pressure has a lower boiling point than when that liquid is at atmospheric pressure. Because of this, water boils at 99.97 °C (211.95 °F) under standard pressure at sea level, but at 93.4 °C (200.1 °F) at 1,905 metres (6,250 ft)[3] altitude. For a given pressure, different liquids will boil at different temperatures.

The normal boiling point (also called the atmospheric boiling point or the atmospheric pressure boiling point) of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, one atmosphere.[4][5] At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid. The standard boiling point has been defined by IUPAC since 1982 as the temperature at which boiling occurs under a pressure of one bar.[6]

The heat of vaporization is the energy required to transform a given quantity (a mol, kg, pound, etc.) of a substance from a liquid into a gas at a given pressure (often atmospheric pressure).

Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid.

Saturation temperature and pressure

Main article: Vapor–liquid equilibrium

A saturated liquid contains as much thermal energy as it can without boiling (or conversely a saturated vapor contains as little thermal energy as it can without condensing).

Saturation temperature means boiling point. The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition.

If the pressure in a system remains constant (isobaric), a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy (heat) is removed. Similarly, a liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied.

The boiling point corresponds to the temperature at which the vapor pressure of the liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Boiling points may be published with respect to the NIST, USA standard pressure of 101.325 kPa (or 1 atm), or the IUPAC standard pressure of 100.000 kPa. At higher elevations, where the atmospheric pressure is much lower, the boiling point is also lower. The boiling point increases with increased pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point. Likewise, the boiling point decreases with decreasing pressure until the triple point is reached. The boiling point cannot be reduced below the triple point.

If the heat of vaporization and the vapor pressure of a liquid at a certain temperature are known, the boiling point can be calculated by using the Clausius–Clapeyron equation, thus:

T B = ( 1 T 0 − R ln ⁡ P P 0 Δ H vap ) − 1 {\displaystyle T_{\text{B}}=\left({\frac {1}{T_{0}}}-{\frac {R\,\ln {\frac {P}{P_{0}}}}{\Delta H_{\text{vap}}}}\right)^{-1}}

where:

T B {\displaystyle T_{B}}

R {\displaystyle R}

P {\displaystyle P}

P 0 {\displaystyle P_{0}}

T 0 {\displaystyle T_{0}}

Δ H vap {\displaystyle \Delta H_{\text{vap}}}

ln {\displaystyle \ln }

Saturation pressure is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased, so is saturation temperature.

If the temperature in a system remains constant (an isothermal system), vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. Similarly, a liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased.

There are two conventions regarding the standard boiling point of water: The normal boiling point is 99.97 °C (211.9 °F) at a pressure of 1 atm (i.e., 101.325 kPa). The IUPAC-recommended standard boiling point of water at a standard pressure of 100 kPa (1 bar)[7] is 99.61 °C (211.3 °F).[6][8] For comparison, on top of Mount Everest, at 8,848 m (29,029 ft) elevation, the pressure is about 34 kPa (255 Torr)[9] and the boiling point of water is 71 °C (160 °F). The Celsius temperature scale was defined until 1954 by two points: 0 °C being defined by the water freezing point and 100 °C being defined by the water boiling point at standard atmospheric pressure.

Relation between the normal boiling point and the vapor pressure of liquids

The higher the vapor pressure of a liquid at a given temperature, the lower the normal boiling point (i.e., the boiling point at atmospheric pressure) of the liquid.

The vapor pressure chart to the right has graphs of the vapor pressures versus temperatures for a variety of liquids.[10] As can be seen in the chart, the liquids with the highest vapor pressures have the lowest normal boiling points.

For example, at any given temperature, methyl chloride has the highest vapor pressure of any of the liquids in the chart. It also has the lowest normal boiling point (−24.2 °C), which is where the vapor pressure curve of methyl chloride (the blue line) intersects the horizontal pressure line of one atmosphere (atm) of absolute vapor pressure.

The critical point of a liquid is the highest temperature (and pressure) it will actually boil at.

See also Vapour pressure of water.

Boiling point of chemical elements

Further information: List of elements by boiling point and Boiling points of the elements (data page)

The element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5000 K at standard pressure; because it is difficult to measure extreme temperatures precisely without bias, both have been cited in the literature as having the higher boiling point.[11]

Boiling point as a reference property of a pure compound

As can be seen from the above plot of the logarithm of the vapor pressure vs. the temperature for any given pure chemical compound, its normal boiling point can serve as an indication of that compound's overall volatility. A given pure compound has only one normal boiling point, if any, and a compound's normal boiling point and melting point can serve as characteristic physical properties for that compound, listed in reference books. The higher a compound's normal boiling point, the less volatile that compound is overall, and conversely, the lower a compound's normal boiling point, the more volatile that compound is overall. Some compounds decompose at higher temperatures before reaching their normal boiling point, or sometimes even their melting point. For a stable compound, the boiling point ranges from its triple point to its critical point, depending on the external pressure. Beyond its triple point, a compound's normal boiling point, if any, is higher than its melting point. Beyond the critical point, a compound's liquid and vapor phases merge into one phase, which may be called a superheated gas. At any given temperature, if a compound's normal boiling point is lower, then that compound will generally exist as a gas at atmospheric external pressure. If the compound's normal boiling point is higher, then that compound can exist as a liquid or solid at that given temperature at atmospheric external pressure, and will so exist in equilibrium with its vapor (if volatile) if its vapors are contained. If a compound's vapors are not contained, then some volatile compounds can eventually evaporate away in spite of their higher boiling points.

In general, compounds with ionic bonds have high normal boiling points, if they do not decompose before reaching such high temperatures. Many metals have high boiling points, but not all. Very generally—with other factors being equal—in compounds with covalently bonded molecules, as the size of the molecule (or molecular mass) increases, the normal boiling point increases. When the molecular size becomes that of a macromolecule, polymer, or otherwise very large, the compound often decomposes at high temperature before the boiling point is reached. Another factor that affects the normal boiling point of a compound is the polarity of its molecules. As the polarity of a compound's molecules increases, its normal boiling point increases, other factors being equal. Closely related is the ability of a molecule to form hydrogen bonds (in the liquid state), which makes it harder for molecules to leave the liquid state and thus increases the normal boiling point of the compound. Simple carboxylic acids dimerize by forming hydrogen bonds between molecules. A minor factor affecting boiling points is the shape of a molecule. Making the shape of a molecule more compact tends to lower the normal boiling point slightly compared to an equivalent molecule with more surface area.

Common name	n-butane	isobutane
IUPAC name	butane	2-methylpropane
Molecular form
Boiling point (°C)	−0.5	−11.7

Common name	n-pentane	isopentane	neopentane
IUPAC name	pentane	2-methylbutane	2,2-dimethylpropane
Molecular form
Boiling point (°C)	36.0	27.7	9.5

Most volatile compounds (anywhere near ambient temperatures) go through an intermediate liquid phase while warming up from a solid phase to eventually transform to a vapor phase. By comparison to boiling, a sublimation is a physical transformation in which a solid turns directly into vapor, which happens in a few select cases such as with carbon dioxide at atmospheric pressure. For such compounds, a sublimation point is a temperature at which a solid turning directly into vapor has a vapor pressure equal to the external pressure.

Impurities and mixtures

In the preceding section, boiling points of pure compounds were covered. Vapor pressures and boiling points of substances can be affected by the presence of dissolved impurities (solutes) or other miscible compounds, the degree of effect depending on the concentration of the impurities or other compounds. The presence of non-volatile impurities such as salts or compounds of a volatility far lower than the main component compound decreases its mole fraction and the solution's volatility, and thus raises the normal boiling point in proportion to the concentration of the solutes. This effect is called boiling point elevation. As a common example, salt water boils at a higher temperature than pure water.

In other mixtures of miscible compounds (components), there may be two or more components of varying volatility, each having its own pure component boiling point at any given pressure. The presence of other volatile components in a mixture affects the vapor pressures and thus boiling points and dew points of all the components in the mixture. The dew point is a temperature at which a vapor condenses into a liquid. Furthermore, at any given temperature, the composition of the vapor is different from the composition of the liquid in most such cases. In order to illustrate these effects between the volatile components in a mixture, a boiling point diagram is commonly used. Distillation is a process of boiling and [usually] condensation which takes advantage of these differences in composition between liquid and vapor phases.

Table

v t e Boiling point of the elements in the periodic table
Group →	1	2		3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
↓ Period
1	H2 20.271 K (−252.879 °C)		He4.222 K (−268.928 °C)
2	Li1603 K (1330 °C)	Be2742 K (2469 °C)		B 4200 K (3927 °C)	C 3915 K (subl.) (3642 °C)	N2 77.355 K (−195.795 °C)	O2 90.188 K (−182.962 °C)	F2 85.03 K (−188.11 °C)	Ne27.104 K (−246.046 °C)
3	Na1156.090 K (882.940 °C)	Mg1363 K (1091 °C)		Al2743 K (2470 °C)	Si3538 K (3265 °C)	P 553.7 K (280.5 °C)	S 717.8 K (444.6 °C)	Cl2239.11 K (−34.04 °C)	Ar87.302 K (−185.848 °C)
4	K 1032 K (759 °C)	Ca1757 K (1484 °C)		Sc3109 K (2836 °C)	Ti3560 K (3287 °C)	V 3680 K (3407 °C)	Cr2945.15 K (2672.0 °C)	Mn2334 K (2061 °C)	Fe3134 K (2861 °C)	Co3200 K (2927 °C)	Ni3003 K (2730 °C)	Cu2835 K (2562 °C)	Zn1180 K (907 °C)	Ga2673 K (2400 °C)	Ge3106 K (2833 °C)	As887 K (subl.) (615 °C)	Se958 K (685 °C)	Br2332.0 K (58.8 °C)	Kr119.735 K (−153.415 °C)
5	Rb961 K (688 °C)	Sr1650 K (1377 °C)		Y 3203 K (2930 °C)	Zr4650 K (4377 °C)	Nb5017 K (4744 °C)	Mo4912 K (4639 °C)	Tc4538 K (4265 °C)	Ru4423 K (4150 °C)	Rh3968 K (3695 °C)	Pd3236 K (2963 °C)	Ag2483 K (2210 °C)	Cd1040 K (767 °C)	In2345 K (2072 °C)	Sn2875 K (2602 °C)	Sb1908 K (1635 °C)	Te1261 K (988 °C)	I2 457.4 K (184.3 °C)	Xe165.051 K (−108.099 °C)
6	Cs944 K (671 °C)	Ba2118 K (1845 °C)		Lu3675 K (3402 °C)	Hf4876 K (4603 °C)	Ta5731 K (5458 °C)	W 6203 K (5930 °C)	Re5900.15 K (5627.0 °C)	Os5285 K (5012 °C)	Ir4403 K (4130 °C)	Pt4098 K (3825 °C)	Au3243 K (2970 °C)	Hg629.88 K (356.73 °C)	Tl1746 K (1473 °C)	Pb2022 K (1749 °C)	Bi1837 K (1564 °C)	Po1235 K (962 °C)	At2503±3 K (230±3 °C)	Rn211.5 K (−61.7 °C)
7	Fr950 K (677 °C)	Ra2010 K (1737 °C)		Lr	Rf5800 K (5500 °C)	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn340±10 K (67±10 °C)	Nh1430 K (1130 °C)	Fl380 K (107 °C)	Mc~1400 K (~1100 °C)	Lv1035–1135 K (762–862 °C)	Ts883 K (610 °C)	Og450±10 K (177±10 °C)
	La3737 K (3464 °C)	Ce3716 K (3443 °C)6	Pr3403 K (3130 °C)	Nd3347 K (3074 °C)	Pm3273 K (3000 °C)	Sm2173 K (1900 °C)	Eu1802 K (1529 °C)	Gd3273 K (3000 °C)	Tb3396 K (3123 °C)	Dy2840 K (2567 °C)	Ho2873 K (2600 °C)	Er3141 K (2868 °C)	Tm2223 K (1950 °C)	Yb1469 K (1196 °C)
	Ac3471 K (3198 °C)	Th5061 K (4788 °C)	Pa4300? K (4027 °C)	U 4404 K (4131 °C)	Np4175.15 K (3902.0 °C)	Pu3508.15 K (3235.0 °C)	Am2880 K (2607 °C)	Cm3383 K (3110 °C)	Bk2900 K (2627 °C)	Cf1743 K (1470 °C)	Es1269  (996 °C)	Fm	Md	No
Legend
Values are in Kelvin K and Celsius °C, rounded
For the equivalent in Fahrenheit °F, see: Boiling points of the elements (data page)
Some values are predictions
Primordial  From decay  Synthetic Border shows natural occurrence of the element s-block g-block f-block d-block p-block

See also

• Boiling points of the elements (data page)
• Boiling-point elevation
• Critical point (thermodynamics)
• Ebulliometer, a device to accurately measure the boiling point of liquids
• Hagedorn temperature
• Joback method (Estimation of normal boiling points from molecular structure)
• List of gases including boiling points
• Melting point
• Subcooling
• Superheating
• Trouton's constant relating latent heat to boiling point
• Triple point

References

1. ^ Goldberg, David E. (1988). 3,000 Solved Problems in Chemistry (1st ed.). McGraw-Hill. section 17.43, p. 321. ISBN 0-07-023684-4.
2. ^ Theodore, Louis; Dupont, R. Ryan; Ganesan, Kumar, eds. (1999). Pollution Prevention: The Waste Management Approach to the 21st Century. CRC Press. section 27, p. 15. ISBN 1-56670-495-2.
3. ^ "Boiling Point of Water and Altitude". www.engineeringtoolbox.com.
4. ^ General Chemistry Glossary Purdue University website page
5. ^ Reel, Kevin R.; Fikar, R. M.; Dumas, P. E.; Templin, Jay M. & Van Arnum, Patricia (2006). AP Chemistry (REA) – The Best Test Prep for the Advanced Placement Exam (9th ed.). Research & Education Association. section 71, p. 224. ISBN 0-7386-0221-3.
6. ^ a b Cox, J. D. (1982). "Notation for states and processes, significance of the word standard in chemical thermodynamics, and remarks on commonly tabulated forms of thermodynamic functions". Pure and Applied Chemistry. 54 (6): 1239–1250. doi:10.1351/pac198254061239.
7. ^ Standard Pressure IUPAC defines the "standard pressure" as being 105 Pa (which amounts to 1 bar).
8. ^ Appendix 1: Property Tables and Charts (SI Units), Scroll down to Table A-5 and read the temperature value of 99.61 °C at a pressure of 100 kPa (1 bar). Obtained from McGraw-Hill's Higher Education website.
9. ^ West, J. B. (1999). "Barometric pressures on Mt. Everest: New data and physiological significance". Journal of Applied Physiology. 86 (3): 1062–6. doi:10.1152/jappl.1999.86.3.1062. PMID 10066724. S2CID 27875962.
10. ^ Perry, R.H.; Green, D.W., eds. (1997). Perry's Chemical Engineers' Handbook (7th ed.). McGraw-Hill. ISBN 0-07-049841-5.
11. ^ DeVoe, Howard (2000). Thermodynamics and Chemistry (1st ed.). Prentice-Hall. ISBN 0-02-328741-1.

External links

• "Boiling-Point" . The New Student's Reference Work . 1914.

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