نقطه جوش

جوشیدن آب

نقطهٔ جوش (به انگلیسی: Boiling point) دمایی است که در آن فشار بخار مایع کاملاً با فشار خارجی برابر می‌شود. در این دما جنبش ملکول‌های مایع به بیشینه مقدار ممکن می‌رسد.

وابستگی نقطه جوش

نقطهٔ جوش یک عنصر، حالتی است که در آن فشار بخار مایع با فشار بخار اطراف مایع یکسان می‌شود. با افزایش فشار، نقطه جوش مایع افزایش می‌یابد و با کاهش آن نقطه جوش کاهش می‌یابد. هم‌چنین نیروهای بین مولکولی و نوع ایزومرهای هندسی نیز در میزان نقطه جوش یک ماده نقش دارند. نقطه جوش یک مایع به صورت درجه حرارتی تعریف می‌شود که در آن فشار بخار مایع برابر با فشار بیرونی باشد. چنانچه فشار خارج ۷۶۰ میلی‌متر جیوه باشد. یک مایع وقتی خواهد جوشید که فشار بخار آن برابر با این مقدار باشد.

برای مثال آب وقتی از دمای ۱۰۰Cْ خواهد جوشید که فشار خارجی و فشار بخار آن برابر با ۷۶۰ میلیمتر جیوه باشد. در فشار ۵۲۶ میلی‌متر جیوه آب در ۹۰Cْ می‌جوشد و اگر فشار را به ۹٫۲ میلی‌متر برسانیم نقطه جوش آب در ۱۰Cْ خواهد بود به این علت است که نقطه جوش آب در ارتفاعات کمتر از مناطق نزدیک به سطح دریا می‌باشد.

چگونگی جوشیدن یک مایع

وقتی که فشار بخار یک مایع با فشار جو برابر می‌شود، مایع شروع به جوشیدن می‌کند. در این دما، بخار حاصل در داخل مایع سبب ایجاد حباب و غلیان خاص جوشش می‌شود. تشکیل حباب در دمای پایینتر از نقطه جوش غیرممکن است، زیرا فشار جو بر سطح مایع که بیش از فشار داخل آن است، مانع از تشکیل حباب می‌شود. دمای مایع در حال جوش تا هنگامی که تمام مایع بخار نشده‌است، ثابت می‌ماند در یک ظرف بدون درپوش حداکثر فشار بخاری که هر مایع می‌تواند داشته باشد برابر با فشار جو می‌باشد.

فشار بخار هر مایع تنها از روی دما معین می‌شود؛ بنابراین اگر فشار بخار ثابت باشد دما نیز ثابت است. برای ثابت ماندن دمای یک مایع در حال جوش باید به آن گرما داده شود. زیرا در فرایند جوش مولکول‌ها با انرژی زیاد از مایع خارج می‌شوند. اگر سرعت افزایش گرما بیش از حداقل لازم برای ثابت نگهداشتن دمای مایع در حال جوش باشد، سرعت جوشش زیاد می‌شود ولی دمای مایع بالا نمی‌رود.

دما و فشار اشباع

مقدار انرژی گرمایی مایع اشباع به میزانی است که اضافه شدن مقدار بسیار بسیار کمی انرژی باعث جوشیدن مایع می‌شود.

دمای اشباع همان دمای جوش است. دمای اشباع وابسته به یک فشاری است که فشار اشباع خوانده می‌شود و آن دمایی است که در آن مایع با شروع جوشش تغییر فاز داده و تبدیل به بخار می‌شود. دمای اشباع و فشار اشباع به همدیگر کاملاً وابسته اند به این معنی که هر فشاری، دمای اشباع مخصوص به خود را دارد یا برعکس هر دمایی فشار اشباع مخصوص به خود دارد و با تغییر هر کدام دیگری نیز تغییر می‌کند.

تأثیر فشار در نقطه جوش

نقطه جوش یک مایع با تغییر فشار خارجی تغییر می‌کند. نقطه جوش نرمال یک مایع، دمایی است که در آن فشار بخار مایع برابر با یک اتمسفر باشد. نقطه جوش داده شده در کتابهای مرجع، نقاط جوش نرمال می‌باشند. نقطه جوش یک مایع را می‌توان از منحنی فشار بخار آن بدست آورد و آن دمایی است که در آن فشار بخار مایع با فشار وارد بر سطح آن برابری می‌کند.

نوسانات فشار جو در یک موقعیت جغرافیایی، نقطه جوش آب را حداکثر تا ۲Cْ تغییر می‌دهد؛ ولی تغییر محل ممکن است باعث تغییرات بیشتر شود، متوسط فشاری که هواسنج در سطح دریا نشان می‌دهد یک اتمسفر، ولی در ارتفاعات بالاتر کمتر از این مقدار است. مثلاً در ارتفاع ۵۰۰۰ پایی (۱۵۲۴ متر) از سطح دریا متوسط فشاری که فشارسنج نشان می‌دهد ۰٫۸۳۶atm است و نقطه جوش آب در این فشار ۹۵٫۱Cْ می‌باشد.

پایین آوردن نقطه جوش یک مایع

اگر نقطه جوش نرمال مایعی بالا باشد یا مایع در اثر گرما تجزیه شود، می‌توان با کاهش فشار آن را در دماهای پایین به جوش آورد از این روش برای تقطیر مایعات در خلأ استفاده می‌شود. مثلاً با کاهش فشار تا ۰٫۰۱۲۱atm می‌توان نقطه جوش آب را به ۱۰Cْ رساند که به‌طور قابل ملاحظه پایینتر از دمای معمولی است، رساند. با کاهش فشار می‌توان آب غیرضروری بسیاری از فراورده‌های غذایی را خارج کرده و آن‌ها را تغلیظ کرد. در این روش دمای فراورده مورد نظر به دمایی که ممکن است سبب تجزیه یا تغییر رنگ آن شود، نمی‌رسد.

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خواص کولیگاتیو

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

Boiling water

The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid 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 high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water boils at 100 °C (212 °F) at sea level, but at 93.4 °C (200.1 °F) at 1,905 metres (6,250 ft) 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. 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.

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

Play media

Demonstration of the lower boiling point of water at lower pressure, achieved by using a vacuum pump.

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:

$T_{\text{B}}=\left({\frac {1}{T_{0}}}-{\frac {R\,\ln {\frac {P}{P_{0}}}}{\Delta H_{\text{vap}}}}\right)^{-1},$

where:

$T_{B}$ is the boiling point at the pressure of interest,

$R$ is the ideal gas constant,

$P$ is the vapour pressure of the liquid at the pressure of interest,

$P_{0}$ is some pressure where the corresponding $T_{0}$ is known (usually data available at 1 atm or 100 kPa),

$\Delta H_{\text{vap}}$ is the heat of vaporization of the liquid,

$T_{0}$ is the boiling temperature,

$\ln$ is the natural logarithm.

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) is 99.61 °C (211.3 °F). For comparison, on top of Mount Everest, at 8,848 m (29,029 ft) elevation, the pressure is about 34 kPa (255 Torr) 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

A log-lin vapor pressure chart for various 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. 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.

Properties of the elements

Further information: List of elements by boiling point

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.

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.

Boiling points of alkanes, alkenes, ethers, halogenoalkanes, aldehydes, ketones, alcohols and carboxylic acids as a function of molar mass

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.

Comparison of butane (C₄H₁₀) isomer boiling points

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

Comparison of pentane isomer boiling points

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

Binary boiling point diagram of two hypothetical only weakly interacting components without an azeotrope

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.

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
• Joback method (Estimation of normal boiling points from molecular structure)
• List of gases including boiling points
• Subcooling
• Superheating
• Trouton's constant relating latent heat to boiling point
• Triple point

External links

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