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Boiling point

The boiling point of a substance is the temperature at which the vapor pressure of its liquid phase equals the pressure surrounding it, resulting in the formation of bubbles throughout the liquid and transition to the gaseous phase.^[1]^[2] The normal boiling point specifically refers to this temperature at standard atmospheric pressure of 1 atmosphere (760 mmHg or 101.3 kPa).^[1] This property serves as a key physical characteristic for identifying and characterizing pure substances in chemistry.^[1] The boiling point is highly sensitive to external pressure; as pressure decreases, the boiling point lowers because less thermal energy is required for the vapor pressure to match the surroundings, which is why water boils below 100°C at high altitudes.^[2] For pure liquids, boiling points are influenced by intermolecular forces, with stronger attractions—such as those from increased molecular weight, polarity, or hydrogen bonding—leading to higher boiling temperatures.^[3] For instance, branched hydrocarbons have lower boiling points than their linear isomers due to reduced surface area for van der Waals interactions, while polar molecules like water exhibit anomalously high boiling points from hydrogen bonding.^[3] In solutions, the boiling point typically elevates compared to the pure solvent, a colligative property dependent on the concentration and number of solute particles rather than their identity.^[4] This elevation, quantified by the formula ΔT_b = K_b × m × i (where K_b is the solvent's boiling point elevation constant, m is molality, and i is the van 't Hoff factor for particle dissociation), explains phenomena like the higher boiling point of saltwater.^[4] Boiling points are practically measured via techniques like the Thiele tube or distillation and play essential roles in processes such as purification, industrial separations, and phase diagrams.^[1]

Fundamentals of Boiling

Definition and Process

The boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the surrounding pressure, typically atmospheric pressure at standard conditions, leading to a phase transition from liquid to vapor throughout the bulk of the liquid. This equilibrium condition allows vapor bubbles to form, grow, and detach from nucleation sites—such as microscopic crevices on the heating surface, impurities, or gas pockets trapped within the liquid—initiating the boiling process.^[5] Once nucleated, these bubbles expand due to the heat input, rise through the liquid due to buoyancy, and release vapor at the surface, facilitating efficient heat transfer.^[6] Boiling differs fundamentally from evaporation, as the latter is a slower, surface-limited process where individual molecules gain sufficient kinetic energy to escape the liquid-air interface without bubble formation, occurring at temperatures below the boiling point.^[7] In contrast, boiling involves vigorous bubble generation and detachment across the liquid volume, driven by the rapid phase change once the saturation temperature is reached.^[8] The first systematic investigations into the influence of pressure on boiling emerged in the 17th century through experiments by Robert Boyle, who used an air pump to demonstrate that reducing atmospheric pressure lowers the boiling temperature of liquids like water.^[9] Illustrations of the boiling process commonly depict vapor bubbles originating from nucleation sites at the bottom of a container, expanding as they ascend through the denser liquid, and rupturing at the free surface to emit steam, highlighting the dynamic convective currents induced by the rising bubbles.

Saturation Temperature and Pressure

The saturation temperature, also known as the boiling point at a given pressure, is defined as the temperature at which the vapor pressure of a liquid equals the pressure of the surrounding system, allowing the liquid and vapor phases to exist in thermodynamic equilibrium.^[10] At this point, the liquid can vaporize without further temperature increase, as the rates of evaporation and condensation balance.^[11] This equilibrium condition is fundamental to phase changes and is observed across various substances under controlled pressures.^[12] The boiling point of a liquid varies inversely with external pressure: higher pressures elevate the saturation temperature by requiring greater molecular energy to overcome the increased resistance to vapor formation, while lower pressures reduce it.^[13] For example, pressure cookers exploit this principle by sealing in steam to build internal pressure, thereby raising the saturation temperature and enabling faster cooking at higher temperatures.^[14] In contrast, at high altitudes where atmospheric pressure drops, the saturation temperature decreases, prolonging cooking times; specifically, it falls by approximately 1°C for every 300 meters of elevation increase due to the reduced ambient pressure.^[15] In a typical pressure-temperature phase diagram for a pure substance, the saturation line—also called the vapor-liquid equilibrium curve—separates the liquid and vapor regions, illustrating how saturation temperature changes with pressure along this boundary.^[16] This curve begins at the triple point, the unique condition where solid, liquid, and vapor phases coexist in equilibrium, and ends at the critical point, beyond which distinct liquid and vapor phases merge into a supercritical fluid. The normal boiling point corresponds to the saturation temperature at standard atmospheric pressure of 1 atm.^[12]

Theoretical Relations

Normal Boiling Point

The normal boiling point of a liquid is defined as the temperature at which its vapor pressure equals 101.325 kPa (1 atm), the standard atmospheric pressure, allowing the liquid to transition to vapor throughout the bulk.^[17] Note that since 1982, IUPAC has recommended the standard boiling point at 1 bar (100 kPa) for standard state conditions, which for water is approximately 99.61 °C, differing slightly from the normal boiling point. This condition, denoted as T_b, represents the saturation temperature specifically at this benchmark pressure and serves as a fundamental reference for comparing the volatility of substances under standardized conditions.^[17] The concept of the normal boiling point emerged in the 19th century as chemists and physicists sought consistent metrics for thermophysical properties, but it was formally standardized by the International Union of Pure and Applied Chemistry (IUPAC) in the 20th century to ensure uniformity in scientific data reporting. This adoption, detailed in IUPAC recommendations from 1994, emphasized the use of 101.325 kPa to align with historical atmospheric pressure conventions while facilitating reproducible measurements in chemical thermodynamics. Prior to broader IUPAC codification, variations in pressure definitions had led to inconsistencies in reported values, prompting the need for this precise benchmark. Measurement of the normal boiling point typically involves ebulliometric or dynamic distillation techniques under controlled conditions to maintain exactly 101.325 kPa. In ebulliometry, the liquid is heated in a specialized apparatus like a Beckmann thermometer-equipped ebulliometer, where the steady-state reflux temperature is recorded as vapor recondenses, ensuring equilibrium at the target pressure. Distillation methods, such as those using a simple or fractional column apparatus, observe the plateau temperature during vaporization while barometric pressure is monitored and adjusted if necessary to match 1 atm. These approaches prioritize purity and pressure control to achieve accuracy within 0.5–1 K for most organic liquids. Values are conventionally reported in degrees Celsius (°C) or Kelvin (K), with the latter absolute scale preferred in thermodynamic calculations; for instance, the normal boiling point of water is 100 °C, equivalent to 373.15 K.^[18] This unit choice reflects practical laboratory conventions, where °C aligns with historical scales, while K ensures additivity in equations without negative values.

Vapor Pressure Connection

The boiling point of a liquid is defined as the temperature at which its vapor pressure equals the surrounding external pressure, marking the onset of boiling where the liquid and vapor phases are in equilibrium.^[2] This equilibrium condition arises because the vapor pressure represents the pressure exerted by the escaping molecules, and when it matches the external pressure, bubbles of vapor can form throughout the liquid without restriction.^[13] The normal boiling point specifically refers to this temperature when the external pressure is 1 atm (101.325 kPa), serving as a standard reference for comparing substances.^[19] The temperature dependence of vapor pressure, which directly governs boiling behavior, is described by the Clausius-Clapeyron equation, derived from thermodynamic principles.^[20] The equation takes the form:

\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_\text{vap}}{R} \left(\frac{1}{T_2} - \frac{1}{T_1}\right)

where P_1 and P_2 are vapor pressures at absolute temperatures T_1 and T_2, \Delta H_\text{vap} is the enthalpy of vaporization, and R is the gas constant.^[21] This relation quantifies how vapor pressure increases exponentially with temperature, explaining why boiling points rise with increasing external pressure.^[19] The derivation begins with the Clapeyron equation, \frac{dP}{dT} = \frac{\Delta H}{T \Delta V}, which relates the slope of the phase boundary in the pressure-temperature diagram to the enthalpy change \Delta H and volume change \Delta V across the phase transition.^[20] For vaporization, \Delta H = \Delta H_\text{vap} and \Delta V \approx V_\text{vapor} = RT/P from the ideal gas law, assuming the liquid volume is negligible compared to the vapor volume.^[21] Integrating this form, with the assumption of constant \Delta H_\text{vap}, yields the Clausius-Clapeyron equation.^[20] This approximation holds reasonably well for many substances over moderate temperature ranges but may deviate at high pressures or near critical points where ideality fails.^[19] In practice, the Clausius-Clapeyron equation enables predictions of boiling point shifts with pressure changes; for instance, it can estimate how the boiling point of water decreases at high altitudes due to lower atmospheric pressure.^[21] For more accurate modeling over wider ranges, empirical correlations like the Antoine equation are often employed, given by:

\log_{10} P = A - \frac{B}{T + C}

where P is vapor pressure in mmHg, T is temperature in °C, and A, B, C are substance-specific constants fitted to experimental data.^[22] This form provides a practical tool for engineering applications, such as distillation processes, by offering a simple way to interpolate vapor pressure curves without relying solely on theoretical assumptions.^[23]

Boiling Points of Pure Substances

Chemical Elements

The normal boiling points of chemical elements, defined as the temperature at which their vapor pressure reaches 1 atm (101.325 kPa), vary dramatically across the periodic table, from cryogenic temperatures for light gases to over 5000 °C for refractory metals.^[24] These values serve as key physical properties for identifying elements and understanding their behavior in chemical processes. Boiling points exhibit distinct periodic trends influenced by bonding types and atomic structure. In groups of non-metals and metalloids, boiling points generally increase down the group due to larger atomic sizes leading to stronger van der Waals forces; for instance, noble gases show low values rising from helium at -268.9 °C to xenon at -108.1 °C.^[25] Metallic elements display higher boiling points overall, with transition metals like tungsten reaching 5555 °C owing to robust delocalized metallic bonding involving d-electrons.^[26] Anomalies occur, such as mercury's relatively low boiling point of 356.7 °C compared to neighboring transition metals, resulting from relativistic effects that contract the 6s orbital, reduce s-p orbital mixing, and weaken interatomic bonds.^[27] Measuring boiling points for reactive or volatile elements poses significant challenges. Alkali metals, highly reactive with oxygen and moisture, require inert atmospheres like argon or nitrogen, often in sealed ampoules or gloveboxes to avoid oxidation during vaporization.^[28] Cryogenic elements and gases, such as hydrogen or helium, demand specialized low-temperature setups like cryostats or dilution refrigerators to achieve and maintain sub-ambient conditions precisely, preventing contamination from atmospheric gases.^[25] The table below provides a comprehensive list of normal boiling points for all 118 elements, compiled from authoritative references including the CRC Handbook of Chemistry and Physics. Values are in °C; some for superheavy elements (atomic numbers 104–118) are theoretical estimates based on empirical trends and quantum calculations, as these elements have not been produced in sufficient quantities for direct measurement. For elements that sublime at 1 atm, the sublimation temperature is provided with a note.^[26]^[25]^[29]

Atomic Number	Element	Symbol	Boiling Point (°C)
1	Hydrogen	H	-252.9
2	Helium	He	-268.9
3	Lithium	Li	1342
4	Beryllium	Be	2470
5	Boron	B	3927
6	Carbon	C	3642 (sublimes)
7	Nitrogen	N	-195.8
8	Oxygen	O	-183.0
9	Fluorine	F	-188.1
10	Neon	Ne	-246.1
11	Sodium	Na	883
12	Magnesium	Mg	1090
13	Aluminum	Al	2467
14	Silicon	Si	3265
15	Phosphorus	P	280
16	Sulfur	S	444.6
17	Chlorine	Cl	-34.0
18	Argon	Ar	-185.8
19	Potassium	K	759
20	Calcium	Ca	1484
21	Scandium	Sc	2836
22	Titanium	Ti	3287
23	Vanadium	V	3913
24	Chromium	Cr	2671
25	Manganese	Mn	2061
26	Iron	Fe	2862
27	Cobalt	Co	2927
28	Nickel	Ni	2913
29	Copper	Cu	2562
30	Zinc	Zn	907
31	Gallium	Ga	2204
32	Germanium	Ge	2830
33	Arsenic	As	614 (sublimes)
34	Selenium	Se	685
35	Bromine	Br	59
36	Krypton	Kr	-153.4
37	Rubidium	Rb	688
38	Strontium	Sr	1382
39	Yttrium	Y	3338
40	Zirconium	Zr	4409
41	Niobium	Nb	4742
42	Molybdenum	Mo	4639
43	Technetium	Tc	4538
44	Ruthenium	Ru	3900
45	Rhodium	Rh	3695
46	Palladium	Pd	2963
47	Silver	Ag	2162
48	Cadmium	Cd	767
49	Indium	In	2072
50	Tin	Sn	2602
51	Antimony	Sb	1587
52	Tellurium	Te	988
53	Iodine	I	184
54	Xenon	Xe	-108.1
55	Cesium	Cs	671
56	Barium	Ba	1897
57	Lanthanum	La	3464
58	Cerium	Ce	3443
59	Praseodymium	Pr	3520
60	Neodymium	Nd	3074
61	Promethium	Pm	3000 (est.)
62	Samarium	Sm	2076
63	Europium	Eu	1597
64	Gadolinium	Gd	3273
65	Terbium	Tb	3232
66	Dysprosium	Dy	2567
67	Holmium	Ho	2700
68	Erbium	Er	2868
69	Thulium	Tm	2540 (est.)
70	Ytterbium	Yb	1196
71	Lutetium	Lu	3402
72	Hafnium	Hf	4602
73	Tantalum	Ta	5458
74	Tungsten	W	5555
75	Rhenium	Re	5596
76	Osmium	Os	5012
77	Iridium	Ir	4130
78	Platinum	Pt	3825
79	Gold	Au	2856
80	Mercury	Hg	356.7
81	Thallium	Tl	1473
82	Lead	Pb	1749
83	Bismuth	Bi	1564
84	Polonium	Po	962
85	Astatine	At	337 (est.)
86	Radon	Rn	-62
87	Francium	Fr	677 (est.)
88	Radium	Ra	1737 (est.)
89	Actinium	Ac	3198
90	Thorium	Th	4788
91	Protactinium	Pa	4171
92	Uranium	U	4131
93	Neptunium	Np	4175
94	Plutonium	Pu	3228
95	Americium	Am	2607
96	Curium	Cm	3100 (est.)
97	Berkelium	Bk	2597 (est.)
98	Californium	Cf	1470 (est.)
99	Einsteinium	Es	1087 (est.)
100	Fermium	Fm	~1000 (est.)
101	Mendelevium	Md	~1100 (est.)
102	Nobelium	No	~1800 (est.)
103	Lawrencium	Lr	~1630 (est.)
104	Rutherfordium	Rf	~2100 (est.)
105	Dubnium	Db	~2200 (est.)
106	Seaborgium	Sg	~2200 (est.)
107	Bohrium	Bh	~2200 (est.)
108	Hassium	Hs	~480 (est.)
109	Meitnerium	Mt	~1800 (est.)
110	Darmstadtium	Ds	~1500 (est.)
111	Roentgenium	Rg	~1570 (est.)
112	Copernicium	Cn	357 (est.)
113	Nihonium	Nh	~1400 (est.)
114	Flerovium	Fl	107 (est.)
115	Moscovium	Mc	~1100 (est.)
116	Livermorium	Lv	404 (est.)
117	Tennessine	Ts	574 (est.)
118	Oganesson	Og	177 (est.)

Recent relativistic quantum calculations for superheavy elements like oganesson predict deviations from group trends; its estimated boiling point of 177 °C arises from enhanced electron stability and reduced volatility compared to lighter noble gases, based on molecular dynamics simulations.^[30]

Reference Property for Compounds

The boiling point serves as a fundamental physical property for identifying and characterizing pure chemical compounds, particularly in processes like distillation and chromatography, where it helps confirm purity and distinguish between substances. For instance, the normal boiling point of ethanol at 78.37 °C allows it to be separated from water, which boils at 100 °C, during simple distillation, enabling verification of the compound's identity based on the observed vaporization temperature.^[31] In gas chromatography, the retention time of a compound correlates with its boiling point, providing a means to compare experimental data against known values for purity assessment and identification. Boiling points of pure organic compounds are extensively documented in authoritative databases, often with high precision to support accurate characterization. PubChem compiles experimental and predicted boiling points for thousands of compounds, drawing from peer-reviewed sources to ensure reliability.^[32] Similarly, the CRC Handbook of Chemistry and Physics lists boiling points for organic compounds to two decimal places (e.g., 0.01 °C precision) under standard conditions of 101.325 kPa, serving as a standard reference for laboratory and industrial applications.^[33] Despite its utility, the boiling point is not always unique for compound identification, as structural isomers can exhibit similar values, requiring combination with other properties like melting point or refractive index for unambiguous characterization. For example, n-butanol and isobutanol have boiling points of 117.7 °C and 107.9 °C, respectively, which are close enough to necessitate additional techniques for differentiation.^[34] Historically, boiling point measurements via fractional distillation played a pivotal role in 19th-century organic chemistry for classifying hydrocarbons, as demonstrated by Auguste Laurent's isolation of naphthalene from coal tar fractions in the 1830s, which helped establish systematic categorization based on volatility.^[35] This approach facilitated the early structural elucidation of complex mixtures, laying groundwork for modern synthetic organic chemistry.^[36]

Influences on Boiling

Impurities and Mixtures

The presence of non-volatile impurities in a liquid solvent acts as a colligative property, lowering the vapor pressure of the solvent and thereby elevating the boiling point of the solution compared to the pure solvent.^[37] This elevation depends solely on the number of solute particles, not their identity, and is particularly relevant in processes like desalination where salts increase the energy required for vaporization.^[38] In multi-component mixtures of volatile liquids, the boiling behavior deviates from that of pure substances, often requiring fractional distillation to separate components as the composition and temperature vary during the process.^[39] For ideal mixtures, Raoult's law governs the total vapor pressure P_{\text{total}}, given by

P_{\text{total}} = x_A P_A^* + x_B P_B^*

where x_A and x_B are the mole fractions of components A and B, and P_A^* and P_B^* are their pure-component vapor pressures at the given temperature.^[39] The boiling point of the mixture occurs when P_{\text{total}} equals the external pressure P_{\text{ext}}, resulting in a temperature intermediate between those of the pure components, weighted by their mole fractions.^[39] Non-ideal mixtures can form azeotropes, which are constant-boiling compositions where the vapor phase has the same composition as the liquid, limiting separation by simple distillation.^[40] Minimum-boiling azeotropes, such as the ethanol-water system at 95.6% ethanol by mass boiling at 78.2°C (compared to 78.4°C for pure ethanol and 100°C for water), exhibit positive deviations from Raoult's law and lower boiling points than either component.^[40] Maximum-boiling azeotropes, conversely, show negative deviations and higher boiling points, as seen in systems like nitric acid-water.^[41] Recent studies from the 2020s have advanced azeotrope prediction and separation using molecular dynamics simulations integrated with experimental data, improving accuracy for binary mixtures beyond traditional thermodynamic models.^[42] In biofuel production, overcoming ethanol-water azeotropes remains critical; innovations like advanced distillation hybrids and membrane pervaporation have reduced energy demands by up to 30% in ethanol dehydration processes since 2020.^[43] These developments, including ionic liquid extractive distillation, enable higher-purity biofuels while addressing sustainability challenges in azeotropic separations.^[44]

Boiling Point Elevation in Solutions

Boiling point elevation refers to the increase in the boiling point of a solvent when a non-volatile solute is added, a colligative property dependent on the number of solute particles rather than their identity.^[45] For aqueous solutions, this elevation is quantified by the formula

\Delta T_b = i \cdot K_b \cdot m

where \Delta T_b is the boiling point elevation in °C, i is the van't Hoff factor representing the effective number of particles per solute molecule, K_b is the ebullioscopic constant of the solvent (0.512 °C/kg/mol for water), and m is the molality of the solution in mol/kg.^[45] This equation applies primarily to dilute solutions where solute-solvent interactions are minimal. In practical examples, seawater with approximately 0.6 molal NaCl equivalent (from 3.5% salinity) exhibits a boiling point of about 100.5°C at 1 atm, due to i \approx 2 for NaCl dissociation into Na⁺ and Cl⁻ ions, yielding \Delta T_b \approx 0.5 °C.^[46] Similarly, a 1 molal NaCl aqueous solution has a boiling point of approximately 101.0°C, as the elevation is \Delta T_b = 2 \cdot 0.512 \cdot 1 = 1.024 °C.^[47] This phenomenon has applications in everyday cooking and industrial processes. Adding salt to water for boiling pasta raises the temperature slightly (e.g., by approximately 0.2–0.4°C for typical concentrations), which can enhance cooking efficiency once boiling begins, though it marginally delays reaching the boil due to the higher required temperature.^[45] In desalination, boiling point elevation in concentrated brines reduces thermal efficiency in evaporation systems by lowering the temperature driving force for vaporization, impacting energy costs in multi-effect distillation.^[48] Several factors influence the accuracy of the elevation formula, particularly the van't Hoff factor i. For electrolytes like NaCl, i approaches the ideal value (e.g., 2) in dilute solutions but decreases in concentrated ones due to ion pairing, where oppositely charged ions associate and behave as fewer particles.^[49] Non-electrolytes, such as sugar in aqueous solutions, have i = 1 since they do not dissociate, resulting in elevation solely proportional to molality.^[45]

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Sep 27, 2024 · This paper seeks to examine the current trends employed in the separation procedures for azeotropic mixtures in industry, particularly the techniques and ...Missing: biofuel | Show results with:biofuel
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Boiling Point Elevation Problems #1-10 - chemteam.info
Problem #3: Seawater is about 3.5% (by weight) dissolved solids, almost all of which is NaCl. Calculate the normal boiling point of seawater. Solution. Some ...
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Does Adding Salt to Water Help it Boil Faster? and Why? - Thermtest
May 25, 2021 · Celsius: Water starts boiling at 100°C under normal atmospheric conditions. Fahrenheit: Water starts boiling at 212°F under normal atmospheric ...
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High salinity seawater boiling point elevation: Experimental verification
May 15, 2021 · Boiling point elevation (BPE) for high salinity solutions is a crucial parameter that significantly affects the overall thermal efficiency ...