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Henry's law
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In chemistry, Henry's law is a gas law that states that the amount of dissolved gas is proportional to its partial pressure in the gas phase. The proportionality factor is called the Henry's law constant. It was formulated by the English chemist William Henry, who studied the topic in the early 19th century. In his publication about the quantity of gases absorbed by water,[1] he described the results of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed under the common pressure of the atmosphere."
An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater divers that changes during decompression, leading to decompression sickness. An everyday example is given by one's experience with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is almost pure carbon dioxide, at a pressure higher than atmospheric pressure. After the bottle is opened, this gas escapes, moving the partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes out of solution.
Contents
[hide]
1 Fundamental types and variants of Henry's law constants
1.1 Henry's law solubility constants H {\displaystyle H}
1.1.1 Henry solubility defined via concentration ( H c p {\displaystyle H^{cp}} )
1.1.2 The dimensionless Henry solubility H c c {\displaystyle H^{cc}}
1.1.3 Henry solubility defined via aqueous-phase mixing ratio ( H x p {\displaystyle H^{xp}} )
1.1.4 Henry solubility defined via molality ( H b p {\displaystyle H^{bp}} )
1.1.5 The Bunsen coefficient α {\displaystyle \alpha }
1.1.6 The Kuenen coefficient S {\displaystyle S}
1.2 Henry's law volatility constants K H {\displaystyle K_{\text{H}}}
1.2.1 The Henry volatility defined via concentration ( K H p c {\displaystyle K_{\text{H}}^{pc}} )
1.2.2 The Henry volatility defined via aqueous-phase mixing ratio ( K H p x {\displaystyle K_{\text{H}}^{px}} )
1.2.3 The dimensionless Henry volatility K H c c {\displaystyle K_{\text{H}}^{cc}}
2 Values of Henry's law constants
3 Temperature dependence
4 Effective Henry's law constants H e f f {\displaystyle H_{\rm {eff}}}
5 Dependence on ionic strength (Sechenov equation)
6 Non-ideal solutions
6.1 Solvent mixtures
7 Miscellaneous
7.1 In geochemistry
7.2 Comparison to Raoult's law
8 See also
9 References
10 External links
Fundamental types and variants of Henry's law constants[edit]
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There are many ways to define the proportionality constant of Henry's law, which can be subdivided into two fundamental types: One possibility is to put the aqueous phase into the numerator and the gas phase into the denominator ("aq/gas").[2] This results in the Henry's law solubility constant H {\displaystyle H} . Its value increases with increased solubility. Alternatively, numerator and denominator can be switched ("gas/aq"), which results in the Henry's law volatility constant K H {\displaystyle K_{\rm {H}}} . The value of K H {\displaystyle K_{\rm {H}}} decreases with increased solubility. There are several variants of both fundamental types. This results from the multiplicity of quantities that can be chosen to describe the composition of the two phases. Typical choices for the aqueous phase are molar concentration ( c a {\displaystyle c_{\rm {a}}} ), molality ( b {\displaystyle b} ), and molar mixing ratio ( x {\displaystyle x} ). For the gas phase, molar concentration ( c g {\displaystyle c_{\rm {g}}} ) and partial pressure ( p {\displaystyle p} ) are often used. It is not possible to use the gas-phase mixing ratio ( y {\displaystyle y} ) because at a given gas-phase mixing ratio, the aqueous-phase concentration c a {\displaystyle c_{\rm {a}}} depends on the total pressure and thus the ratio y / c a {\displaystyle y/c_{\rm {a}}} is not a constant.[3] To specify the exact variant of the Henry's law constant, two superscripts are used. They refer to the numerator and the denominator of the definition. For example, H c p {\displaystyle H^{cp}} refers to the Henry solubility defined as c / p {\displaystyle c/p} .
Henry's law solubility constants H {\displaystyle H} [edit]
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Henry solubility defined via concentration ( H c p {\displaystyle H^{cp}} )[edit]
Atmospheric chemists often define the Henry solubility as
H c p = c a / p {\displaystyle H^{cp}=c_{\text{a}}/p} .[2]
Here c a {\displaystyle c_{\text{a}}} is the concentration of a species in the aqueous phase, and p {\displaystyle p} is the partial pressure of that species in the gas phase under equilibrium conditions.[citation needed]
The SI unit for H c p {\displaystyle H^{cp}} is mol/(m3 Pa); however, often the unit M/atm is used, since c a {\displaystyle c_{\text{a}}} is usually expressed in M (1 M = 1 mol/dm3) and p {\displaystyle p} in atm (1 atm = 101325 Pa).[citation needed]
The dimensionless Henry solubility H c c {\displaystyle H^{cc}} [edit]
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration c a {\displaystyle c_{\text{a}}} of a species and its gas-phase concentration c g {\displaystyle c_{\text{g}}} :
H c c = c a / c g {\displaystyle H^{cc}=c_{\text{a}}/c_{\text{g}}} .[2]
For an ideal gas, the conversion is:
H c c = H c p × R T {\displaystyle H^{cc}=H^{cp}\times RT} ,[2]
where R {\displaystyle R} is the gas constant and T {\displaystyle T} is the temperature.
Sometimes, this dimensionless constant is called the "water-air partitioning coefficient" K WA {\displaystyle K_{\text{WA}}} .[4] It is closely related to the various, slightly different definitions of the "Ostwald coefficient" L {\displaystyle L} , as discussed by Battino (1984).[5]
Henry solubility defined via aqueous-phase mixing ratio ( H x p {\displaystyle H^{xp}} )[edit]
Another Henry's law solubility constant is
H x p = x / p {\displaystyle H^{xp}=x/p} .[2]
Here x {\displaystyle x} is the molar mixing ratio in the aqueous phase. For a dilute aqueous solution the conversion between x {\displaystyle x} and c a {\displaystyle c_{\text{a}}} is:
c a ≈ x ϱ H 2 O M H 2 O {\displaystyle c_{\text{a}}\approx x{\frac {\varrho _{\mathrm {H_{2}O} }}{M_{\mathrm {H_{2}O} }}}} ,[2]
where ϱ H 2 O {\displaystyle \varrho _{\mathrm {H_{2}O} }} is the density of water and M H 2 O {\displaystyle M_{\mathrm {H_{2}O} }} is the molar mass of water. Thus
H x p ≈ M H 2 O ϱ H 2 O × H c p {\displaystyle H^{xp}\approx {\frac {M_{\mathrm {H_{2}O} }}{\varrho _{\mathrm {H_{2}O} }}}\times H^{cp}} .[2]
The SI unit for H x p {\displaystyle H^{xp}} is Pa−1, although atm−1 is still frequently used.[2]
Henry solubility defined via molality ( H b p {\displaystyle H^{bp}} )[edit]
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It can be advantageous to describe the aqueous phase in terms of molality instead of concentration. The molality of a solution does not change with T {\displaystyle T} , since it refers to the mass of the solvent. In contrast, the concentration c {\displaystyle c} does change with T {\displaystyle T} , since the density of a solution and thus its volume are temperature-dependent. Defining the aqueous-phase composition via molality has the advantage that any temperature dependence of the Henry's law constant is a true solubility phenomenon and not introduced indirectly via a density change of the solution. Using molality, the Henry solubility can be defined as
H b p = b / p . {\displaystyle H^{bp}=b/p.}
Here b {\displaystyle b} is used as the symbol for molality (instead of m {\displaystyle m} ) to avoid confusion with the symbol m {\displaystyle m} for mass. The SI unit for H b p {\displaystyle H^{bp}} is mol/(kg Pa). There is no simple way to calculate H c p {\displaystyle H^{cp}} from H b p {\displaystyle H^{bp}} , since the conversion between concentration c a {\displaystyle c_{\text{a}}} and molality b {\displaystyle b} involves all solutes of a solution. For a solution with a total of n {\displaystyle n} solutes with indices i = 1 , … , n {\displaystyle i=1,\ldots ,n} , the conversion is:
c a = b ϱ 1 + ∑ i = 1 n b i M i , {\displaystyle c_{\text{a}}={\frac {b\varrho }{1+\sum _{i=1}^{n}b_{i}M_{i}}},}
where ϱ {\displaystyle \varrho } is the density of the solution, and M i {\displaystyle M_{i}} are the molar masses. Here b {\displaystyle b} is identical to one of the b i {\displaystyle b_{i}} in the denominator. If there is only one solute, the equation simplifies to
c a = b ϱ 1 + b M . {\displaystyle c_{\text{a}}={\frac {b\varrho }{1+bM}}.}
Henry's law is only valid for dilute solutions where b M ≪ 1 {\displaystyle bM\ll 1} and ϱ ≈ ϱ H 2 O {\displaystyle \varrho \approx \varrho _{\mathrm {H_{2}O} }} . In this case the conversion reduces further to
c a ≈ b ϱ H 2 O , {\displaystyle c_{\text{a}}\approx b\varrho _{\mathrm {H_{2}O} },}
and thus
H b p ≈ H c p / ϱ H 2 O . {\displaystyle H^{bp}\approx H^{cp}/\varrho _{\mathrm {H_{2}O} }.}
The Bunsen coefficient α {\displaystyle \alpha } [edit]
According to Sazonov and Shaw, the dimensionless Bunsen coefficient α {\displaystyle \alpha } is defined as "the volume of saturating gas, V1, reduced to T° = 273.15 K, p° = 1 bar, which is absorbed by unit volume V2* of pure solvent at the temperature of measurement and partial pressure of 1 bar."[6] If the gas is ideal, the pressure cancels out, and the conversion to H c p {\displaystyle H^{cp}} is simply
H c p = α × 1 R T STP {\displaystyle H^{cp}=\alpha \times {\frac {1}{RT^{\text{STP}}}}} ,[citation needed][original research?]
with T STP {\displaystyle T^{\text{STP}}} = 273.15 K. Note, that according to this definition, the conversion factor is not temperature-dependent.[citation needed] Independent of the temperature that the Bunsen coefficient refers to, 273.15 K is always used for the conversion.[citation needed] The Bunsen coefficient, which is named after Robert Bunsen, has been used mainly in the older literature.[citation needed]
The Kuenen coefficient S {\displaystyle S} [edit]
According to Sazonov and Shaw, the Kuenen coefficient S {\displaystyle S} is defined as "the volume of saturating gas V(g), reduced to T° = 273.15 K, p° = bar, which is dissolved by unit mass of pure solvent at the temperature of measurement and partial pressure 1 bar."[6] If the gas is ideal, the relation to H c p {\displaystyle H^{cp}} is
H c p = S × ϱ R T STP {\displaystyle H^{cp}=S\times {\frac {\varrho }{RT^{\text{STP}}}}} ,[citation needed][original research?]
where ϱ {\displaystyle \varrho } is the density of the solvent, and T STP {\displaystyle T^{\text{STP}}} = 273.15 K. The SI unit for S {\displaystyle S} is m3/kg.[6] The Kuenen coefficient, which is named after Johannes Kuenen, has been used mainly in the older literature, and IUPAC considers it to be obsolete.[7]
Henry's law volatility constants K H {\displaystyle K_{\text{H}}} [edit]
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The Henry volatility defined via concentration ( K H p c {\displaystyle K_{\text{H}}^{pc}} )[edit]
A common way to define a Henry volatility is dividing the partial pressure by the aqueous-phase concentration:
K H p c = p / c a = 1 / H c p . {\displaystyle K_{\text{H}}^{pc}=p/c_{\text{a}}=1/H^{cp}.}
The SI unit for K H p c {\displaystyle K_{\text{H}}^{pc}} is Pa m3/mol.
The Henry volatility defined via aqueous-phase mixing ratio ( K H p x {\displaystyle K_{\text{H}}^{px}} )[edit]
Another Henry volatility is
K H p x = p / x = 1 / H x p . {\displaystyle K_{\text{H}}^{px}=p/x=1/H^{xp}.}
The SI unit for K H p x {\displaystyle K_{\text{H}}^{px}} is Pa. However, atm is still frequently used.
The dimensionless Henry volatility K H c c {\displaystyle K_{\text{H}}^{cc}} [edit]
The Henry volatility can also be expressed as the dimensionless ratio between the gas-phase concentration c g {\displaystyle c_{\text{g}}} of a species and its aqueous-phase concentration c a {\displaystyle c_{\text{a}}} :
K H c c = c g / c a = 1 / H c c . {\displaystyle K_{\text{H}}^{cc}=c_{\text{g}}/c_{\text{a}}=1/H^{cc}.}
In chemical engineering and environmental chemistry, this dimensionless constant is often called the air–water partitioning coefficient K AW {\displaystyle K_{\text{AW}}} .
Values of Henry's law constants[edit]
A large compilation of Henry's law constants has been published by Sander (2015).[2] A few selected values are shown in the table below:
Henry's law constants (gases in water at 298.15 K)
equation:
K H p c = p c aq {\displaystyle K_{\text{H}}^{pc}={\frac {p}{c_{\text{aq}}}}}
H c p = c aq p {\displaystyle H^{cp}={\frac {c_{\text{aq}}}{p}}}
K H p x = p x {\displaystyle K_{\text{H}}^{px}={\frac {p}{x}}}
H c c = c aq c gas {\displaystyle H^{cc}={\frac {c_{\text{aq}}}{c_{\text{gas}}}}}
unit:
L ⋅ atm mol {\displaystyle {\frac {{\text{L}}\cdot {\text{atm}}}{\text{mol}}}}
mol L ⋅ atm {\displaystyle {\frac {\text{mol}}{{\text{L}}\cdot {\text{atm}}}}}
atm {\displaystyle {\text{atm}}}
(dimensionless)
O2 770 1.3×10−3 4.3×104 3.2×10−2
H2 1300 7.8×10−4 7.1×104 1.9×10−2
CO2 29 3.4×10−2 1.6×103 8.3×10−1
N2 1600 6.1×10−4 9.1×104 1.5×10−2
He 2700 3.7×10−4 1.5×105 9.1×10−3
Ne 2200 4.5×10−4 1.2×105 1.1×10−2
Ar 710 1.4×10−3 4.0×104 3.4×10−2
CO 1100 9.5×10−4 5.8×104 2.3×10−2
Temperature dependence[edit]
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When the temperature of a system changes, the Henry constant also changes. The temperature dependence of equilibrium constants can generally be described with the van 't Hoff equation, which also applies to Henry's law constants:
d ln  H d ( 1 / T ) = − Δ sol H R , {\displaystyle {\frac {\mathrm {d} \ln H}{\mathrm {d} (1/T)}}={\frac {-\Delta _{\text{sol}}H}{R}},}
where Δ sol H {\displaystyle \Delta _{\text{sol}}H} is the enthalpy of dissolution. Note that the letter H {\displaystyle H} in the symbol Δ sol H {\displaystyle \Delta _{\text{sol}}H} refers to enthalpy and is not related to the letter H {\displaystyle H} for Henry's law constants. Integrating the above equation and creating an expression based on H ∘ {\displaystyle H^{\circ }} at the reference temperature T ∘ {\displaystyle T^{\circ }} = 298.15 K yields:
H ( T ) = H ∘ × exp [ − Δ sol H R ( 1 T − 1 T ∘ ) ] . {\displaystyle H(T)=H^{\circ }\times \exp \displaystyle \left[{\frac {-\Delta _{\text{sol}}H}{R}}\left({\frac {1}{T}}-{\frac {1}{T^{\circ }}}\right)\right].}
The van 't Hoff equation in this form is only valid for a limited temperature range in which Δ sol H {\displaystyle \Delta _{\text{sol}}H} does not change much with temperature.
The following table lists some temperature dependencies:
Values of − Δ sol H / R {\displaystyle -\Delta _{\text{sol}}H/R} (in K)
O2 H2 CO2 N2 He Ne Ar CO
1700 500 2400 1300 230 490 1300 1300
Solubility of permanent gases usually decreases with increasing temperature at around room temperature. However, for aqueous solutions, the Henry's law solubility constant for many species goes through a minimum. For most permanent gases, the minimum is below 120 °C. Often, the smaller the gas molecule (and the lower the gas solubility in water), the lower the temperature of the maximum of the Henry's law constant. Thus, the maximum is at about 30 °C for helium, 92 to 93 °C for argon, nitrogen and oxygen, and 114 °C for xenon.[8]
Effective Henry's law constants H e f f {\displaystyle H_{\rm {eff}}} [edit]
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The Henry's law constants mentioned so far do not consider any chemical equilibria in the aqueous phase. This type is called the "intrinsic" (or "physical") Henry's law constant. For example, the intrinsic Henry's law solubility constant of formaldehyde can be defined as
H c p = c ( H 2 C O ) p ( H 2 C O ) . {\displaystyle H^{cp}={\frac {c(\mathrm {H_{2}CO} )}{p(\mathrm {H_{2}CO} )}}.}
In aqueous solution, methanal is almost completely hydrated:
H 2 C O + H 2 O ⇌ H 2 C ( O H ) 2 . {\displaystyle \mathrm {H_{2}CO} +\mathrm {H_{2}O} \rightleftharpoons \mathrm {H_{2}C(OH)_{2}} .}
The total concentration of dissolved methanal is
c tot = c ( H 2 C O ) + c ( H 2 C ( O H ) 2 ) . {\displaystyle c_{\text{tot}}=c(\mathrm {H_{2}CO} )+c(\mathrm {H_{2}C(OH)_{2}} ).}
Taking this equilibrium into account, an effective Henry's law constant H eff {\displaystyle H_{\text{eff}}} can be defined as
H eff = c tot p ( H 2 C O ) = c ( H 2 C O ) + c ( H 2 C ( O H ) 2 ) p ( H 2 C O ) . {\displaystyle H_{\text{eff}}={\frac {c_{\text{tot}}}{p(\mathrm {H_{2}CO} )}}={\frac {c(\mathrm {H_{2}CO} )+c(\mathrm {H_{2}C(OH)_{2})} }{p(\mathrm {H_{2}CO} )}}.}
For acids and bases, the effective Henry's law constant is not a useful quantity because it depends on the pH of the solution.[verification needed] In order to obtain a pH-independent constant, the product of the intrinsic Henry's law constant H c p {\displaystyle H^{cp}} and the acidity constant K A {\displaystyle K_{\text{A}}} is often used for strong acids like hydrochloric acid (HCl):
H ′ = H c p × K A = c ( H + ) × c ( C l − ) p ( H C l ) . {\displaystyle H'=H^{cp}\times K_{\text{A}}={\frac {c(\mathrm {H^{+}} )\times c(\mathrm {Cl^{-}} )}{p(\mathrm {HCl} )}}.}
Although H ′ {\displaystyle H'} is usually also called a Henry's law constant, it should be noted that it is a different quantity and it has different units than H c p {\displaystyle H^{cp}} .
Dependence on ionic strength (Sechenov equation)[edit]
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Values of Henry's law constants for aqueous solutions depend on the composition of the solution, i.e., on its ionic strength and on dissolved organics. In general, the solubility of a gas decreases with increasing salinity ("salting out"). However, a "salting in" effect has also been observed, for example for the effective Henry's law constant of glyoxal. The effect can be described with the Sechenov equation, named after the Russian physiologist Ivan Sechenov (sometimes the German transliteration "Setschenow" of the Cyrillic name Се́ченов is used). There are many alternative ways to define the Sechenov equation, depending on how the aqueous-phase composition is described (based on concentration, molality, or molar fraction) and which variant of the Henry's law constant is used. Describing the solution in terms of molality is preferred because molality is invariant to temperature and to the addition of dry salt to the solution. Thus, the Sechenov equation can be written as
log  ( H 0 b p H b p ) = k s × b ( salt ) , {\displaystyle \log \left({\frac {H_{0}^{bp}}{H^{bp}}}\right)=k_{\text{s}}\times b({\text{salt}}),}
where H 0 b p {\displaystyle H_{0}^{bp}} is the Henry's law constant in pure water, H b p {\displaystyle H^{bp}} is the Henry's law constant in the salt solution, k s {\displaystyle k_{\text{s}}} is the molality-based Sechenov constant, and b ( salt ) {\displaystyle b({\text{salt}})} is the molality of the salt.
Non-ideal solutions[edit]
Henry's law has been shown to apply to a wide range of solutes in the limit of "infinite dilution" (x → 0), including non-volatile substances such as sucrose. In these cases, it is necessary to state the law in terms of chemical potentials. For a solute in an ideal dilute solution, the chemical potential depends only on the concentration. For non-ideal solutions, the activity coefficients of the components must be taken into account:
μ = μ c ∘ + R T ln  γ c c c ∘ {\displaystyle \mu =\mu _{c}^{\circ }+RT\ln {\frac {\gamma _{c}c}{c^{\circ }}}} ,
where γ c = K H , c p ∗ {\displaystyle \gamma _{c}={\frac {K_{{\text{H}},c}}{p^{*}}}} for a volatile solute; c° = 1 mol/L.
For non-ideal solutions, the activity coefficient γc depends on the concentration and must be determined at the concentration of interest. The activity coefficient can also be obtained for non-volatile solutes, where the vapor pressure of the pure substance is negligible, by using the Gibbs-Duhem relation:
∑ i n i d μ i = 0. {\displaystyle \sum _{i}n_{i}d\mu _{i}=0.}
By measuring the change in vapor pressure (and hence chemical potential) of the solvent, the chemical potential of the solute can be deduced.
The standard state for a dilute solution is also defined in terms of infinite-dilution behavior. Although the standard concentration c° is taken to be 1 mol/l by convention, the standard state is a hypothetical solution of 1 mol/l in which the solute has its limiting infinite-dilution properties. This has the effect that all non-ideal behavior is described by the activity coefficient: the activity coefficient at 1 mol/l is not necessarily unity (and is frequently quite different from unity).
All the relations above can also be expressed in terms of molalities b rather than concentrations, e.g.:
μ = μ b ∘ + R T ln  γ b b b ∘ , {\displaystyle \mu =\mu _{b}^{\circ }+RT\ln {\frac {\gamma _{b}b}{b^{\circ }}},}
where γ b = K H , b p ∗ {\displaystyle \gamma _{b}={\frac {K_{{\text{H}},b}}{p^{*}}}} for a volatile solute; b° = 1 mol/kg.
The standard chemical potential μm°, the activity coefficient γm and the Henry's law constant KH,b all have different numerical values when molalities are used in place of concentrations.
Solvent mixtures[edit]
Henry law constant H2, M for a gas 2 in a mixture of solvents 1 and 3 is related to the constants for individual solvents H21 and H23:
l n H 2 M = x 1 l n H 21 + x 2 l n H 23 − a 13 x 1 x 3 {\displaystyle lnH_{2M}=x_{1}lnH_{21}+x_{2}lnH_{23}-a_{13}x_{1}x_{3}}
where a13 is the interaction parameter of the solvents from Wohl expansion of the excess chemical potential of the ternary mixtures.
Miscellaneous[edit]
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In geochemistry[edit]
In geochemistry, a version of Henry's law applies to the solubility of a noble gas in contact with silicate melt. One equation used is
C melt / C gas = exp  [ − β ( μ melt E − μ gas E ) ] , {\displaystyle C_{\text{melt}}/C_{\text{gas}}=\exp \left[-\beta (\mu _{\text{melt}}^{\text{E}}-\mu _{\text{gas}}^{\text{E}})\right],}
where
C is the number concentrations of the solute gas in the melt and gas phases,
β = 1/kBT, an inverse temperature parameter (kB is the Boltzmann constant),
µE is the excess chemical potentials of the solute gas in the two phases.
Comparison to Raoult's law[edit]
Henry's law is a limiting law that only applies for "sufficiently dilute" solutions. The range of concentrations in which it applies becomes narrower the more the system diverges from ideal behavior. Roughly speaking, that is the more chemically "different" the solute is from the solvent.
For a dilute solution, the concentration of the solute is approximately proportional to its mole fraction x, and Henry's law can be written as
p = K H x . {\displaystyle p=K_{\text{H}}x.}
This can be compared with Raoult's law:
p = p ∗ x , {\displaystyle p=p^{*}x,}
where p* is the vapor pressure of the pure component.
At first sight, Raoult's law appears to be a special case of Henry's law, where KH = p*. This is true for pairs of closely related substances, such as benzene and toluene, which obey Raoult's law over the entire composition range: such mixtures are called "ideal mixtures".
The general case is that both laws are limit laws, and they apply at opposite ends of the composition range. The vapor pressure of the component in large excess, such as the solvent for a dilute solution, is proportional to its mole fraction, and the constant of proportionality is the vapor pressure of the pure substance (Raoult's law). The vapor pressure of the solute is also proportional to the solute's mole fraction, but the constant of proportionality is different and must be determined experimentally (Henry's law). In mathematical terms:
Raoult's law: lim x → 1 ( p x ) = p ∗ . {\displaystyle \lim _{x\to 1}\left({\frac {p}{x}}\right)=p^{*}.}
Henry's law: lim x → 0 ( p x ) = K H . {\displaystyle \lim _{x\to 0}\left({\frac {p}{x}}\right)=K_{\text{H}}.}
Raoult's law can also be related to non-gas solutes.
See also[edit]
Pervaporation
Partition coefficient
Sieverts' law
Graham's law
Henry adsorption constant
References[edit]
Jump up ^ Henry, W. (1803). "Experiments on the quantity of gases absorbed by water, at different temperatures, and under different pressures". Phil. Trans. R. Soc. Lond. 93: 29–274. doi:10.1098/rstl.1803.0004.
^ Jump up to: a b c d e f g h i Sander, R. (2015), "Compilation of Henry's law constants (version 4.0) for water as solvent", Atmos. Chem. Phys., 15: 4399–4981, Bibcode:2015ACP....15.4399S, doi:10.5194/acp-15-4399-2015
Jump up ^ Sander, R. (1999). "Modeling atmospheric chemistry: Interactions between gas-phase species and liquid cloud/aerosol particles" (review). Surv. Geophys. 20: 1–31. Retrieved 21 March 2016.
Jump up ^ P.J. McCall; R.L. Swann; D.A. Laskowski, "6. Partition Models for Equilibrium Distribution of Chemicals in Environmental Compartments", in I.H. Suffet, Fate of Chemicals in the Environment
Jump up ^ Battino, R.; Rettich, T. R.; Tominaga, T. (1984). "The solubility of nitrogen and air in liquids". J. Phys. Chem. Ref. Data. 13: 563–600.
^ Jump up to: a b c Sazonov, V P & Shaw, DG (2006). "Introduction to the Solubility Data Series: 1.5.2. §Physicochemical Quantities and Units, A note on nomenclature, points 10 and 11". Introduction to IUPAC-NIST Solubilities Database. nist.gov. Retrieved 21 March 2016. CS1 maint: Uses authors parameter (link)
Jump up ^ Gamsjäger, H.; Lorimer, J. W.; Salomon, M.; Shaw, D. G.; Tomkins, R. P. T. (2010). "The IUPAC-NIST Solubility Data Series: A guide to preparation and use of compilations and evaluations (IUPAC Technical Report)". Pure Appl. Chem. 82: 1137–1159. doi:10.1351/pac-rep-09-10-33.
Jump up ^ Cohen, P., ed. (1989). The ASME Handbook on Water Technology for Thermal Power Systems. The American Society of Mechanical Engineers. p. 442. ISBN 978-0-7918-0634-0.
External links[edit]
EPA On-line Tools for Site Assessment Calculation – Henry's law conversion
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text:
Summer Holiday (song)
From Wikipedia, the free encyclopedia
Jump to navigation Jump to search
"Summer Holiday"
Single by Cliff Richard and The Shadows
from the album Summer Holiday
B-side
"Dancing Shoes"
Released
2 June 1963
Format
7" single
Recorded
9 May & 19 November 1962, EMI Studios, London
Genre
Pop
Length
2:03
Label
Columbia (EMI) DB 4977[1]
Songwriter(s)
Bruce Welch, Brian Bennett[1]
Producer(s)
Norrie Paramor[1]
Cliff Richard and The Shadows singles chronology
"String Module Error: Match not found"
(String Module Error: Match not found) "Summer Holiday"
(1963) "Lucky Lips"
(1963)
"The Next Time" /
Bachelor Boy (1962) "Summer Holiday"
(1963) "Lucky Lips"
(1963)
"Summer Holiday" is a song recorded by Cliff Richard and The Shadows, written by rhythm guitarist Bruce Welch and drummer Brian Bennett.[1] It is taken from the film of the same name,[1] and was released as the second single from the film in February 1963. It went to number one in the UK Singles Chart for a total of three weeks, as had the first single from the film, "The Next Time".[2] After "Summer Holiday" had spent two weeks at number one, The Shadows' instrumental "Foot Tapper" - also from the same film - took over the top spot for one week, before "Summer Holiday" returned to the top spot for one further week.[3] The track is one of Richard's best known titles and it remains a staple of his live shows. It was one of six hits Richard performed at his spontaneous gig at the 1996 Wimbledon Championships when rain stopped the tennis.
The melody of the song is used in the chorus of the 1986 rap tune "Holiday Rap", by the Dutch duo MC Miker G & DJ Sven.[citation needed]
Contents
1 Chart performance
1.1 Summer Holiday
1.2 Dancing Shoes
2 References
3 External links
Chart performance[edit]
Summer Holiday[edit]
Chart (1963)
Peak
position
Australia (Kent Music Report)[4] 3
Australia (Music Maker, Sydney)[5] 1
Belgium (Ultratop 50 Flanders)[6] 5
Belgium (Ultratop 50 Wallonia)[7] 1
Canada (CHUM)[8] 1
Denmark (Quan Musikbureau)[9][10] 1
Finland (IFPI Finland)[11] 6
Hong Kong[12][13] 2
Ireland (IRMA)[14] 2
Israel (Kol)[5] 2
Netherlands (Single Top 100)[15] 2
New Zealand [12] 2[*]
Norway (VG-lista)[16] 1
Spain (Promusicae)[17] 1
South Africa (SARMDA)[5] 10
UK Singles (Official Charts Company)[18] 1
'Notes:
* ^ Not all New Zealand chart weeks were published in Billboard. The record may have charted higher.
Dancing Shoes[edit]
The B-side "Dancing Shoes" also entered some charts, some of which are listed below.
Chart (1963)
Peak
position
Hong Kong[9] 6
South Africa (SARMDA)[12] 4
UK (NME)[19] 25
References[edit]
^ Jump up to: a b c d e Rice, Jo (1982). The Guinness Book of 500 Number One Hits (1st ed.). Enfield, Middlesex: Guinness Superlatives Ltd. p. 70. ISBN 0-85112-250-7.
Jump up ^ Roberts, David (2006). British Hit Singles & Albums (19th ed.). London: Guinness World Records Limited. pp. 147–8. ISBN 1-904994-10-5.
Jump up ^ "Number 1 Singles – 1960s". The Official Charts Company. Archived from the original on 8 April 2008. Retrieved 27 April 2014.
Jump up ^ Kent, David (2005). Australian Chart Book 1940–1969. Turramurra, New South Wales: Australian Chart Book Pty. ISBN 0-646-44439-5.
^ Jump up to: a b c Nielsen Business Media, Inc. Israel chart courtesy of Kol Israel Broadcasting, South Africa chart courtesy of Southern African Record Manufacturers and Distributors Association (25 May 1963). Billboard. Nielsen Business Media, Inc. p. 45. ISSN 0006-2510. CS1 maint: Multiple names: authors list (link)
Jump up ^ "Ultratop.be – Cliff Richard – Summer Holiday" (in Dutch). Ultratop 50.
Jump up ^ "Ultratop.be – Cliff Richard – Summer Holiday" (in French). Ultratop 50.
Jump up ^ "CHUM Chart". |
^ Jump up to: a b Nielsen Business Media, Inc. (4 May 1963). Billboard. Nielsen Business Media, Inc. p. 28. ISSN 0006-2510.
Jump up ^ "Denmark Singles Chart (Top 20/30) – Summer Holiday". danskehitlister.dk. Retrieved 2018-06-09.
Jump up ^ Lassila, Juha (1990). Mitä Suomi soittaa?: Hittilistat 1954-87 (in Finnish). Jyväskylän yliopisto. ISBN 95-168-0321-0.
^ Jump up to: a b c Nielsen Business Media, Inc. (11 May 1963). Billboard. Nielsen Business Media, Inc. p. 36. ISSN 0006-2510.
Jump up ^ Nielsen Business Media, Inc. (18 May 1963). Billboard. Nielsen Business Media, Inc. pp. 32–. ISSN 0006-2510.
Jump up ^ "Ireland singles charts". Irishcharts.ie. Retrieved 2018-03-17.
Jump up ^ "Dutchcharts.nl – Cliff Richard – Summer Holiday" (in Dutch). Single Top 100.
Jump up ^ "Norwegiancharts.com – Cliff Richard – Summer Holiday". VG-lista.
Jump up ^ Salaverri, Fernando (September 1979). Sólo éxitos: Año a año, 1959–2002. Spain: Fundación Autor-SGAE. ISBN 978-84-8048-639-2.
Jump up ^ "Cliff Richard: Artist Chart History". Official Charts Company.
Jump up ^ Read, Mike; Lewry, Peter; Goodall, Nigel (1993). Cliff Richard - The Complete Chronicle. London: Hamlyn. p. 303. ISBN 0600578976. Retrieved 1 July 2013.
External links[edit]
Song info
Lyrics of this song at MetroLyrics
show
v
t
e
Cliff Richard
Albums (UK)
1950s:
Cliff (1959, live)
Cliff Sings (1959)
1960s:
Me and My Shadows (1960)
21 Today (1961)
Listen to Cliff! (1961)
32 Minutes and 17 Seconds (1962)
The Young Ones (1962, soundtrack)
Summer Holiday (1963, soundtrack)
When in Spain (1963)
Wonderful Life (1964, soundtrack)
Cliff Richard (1965)
When in Rome (1965)
Love is Forever (1965)
Kinda Latin (1966)
Finders Keepers (1966, soundtrack)
Two a Penny (1968, soundtrack)
1970s:
Tracks 'n Grooves (1970)
Take Me High (1974, soundtrack)
I'm Nearly Famous (1976)
Every Face Tells a Story (1977)
Small Corners (1978, gospel)
Green Light (1978)
Thank You Very Much (1979, live)
Rock 'n' Roll Juvenile (1979)
1980s:
I'm No Hero (1980)
Wired for Sound (1981)
Now You See Me, Now You Don't (1982, gospel)
Dressed for the Occasion (1983, live)
Silver (1983)
The Rock Connection (1984)
Always Guaranteed (1987)
Stronger (1989)
1990s:
From a Distance: The Event (1990, live)
Together with Cliff Richard (1991)
The Album (1993)
Songs from Heathcliff (1995, stage show album)
Real as I Wanna Be (1998)
2000s:
Wanted (2001)
Cliff at Christmas (2003)
Something's Goin' On (2004)
Two's Company The Duets (2006)
Love... The Album (2007)
Reunited – Cliff Richard and The Shadows (2009)
2010s:
Bold as Brass (2010)
Soulicious (2011)
The Fabulous Rock 'n' Roll Songbook (2013)
Just... Fabulous Rock 'n' Roll (2016)
EPs (UK)
Serious Charge
Compilations (UK)
Cliff's Hit Album (1963)
40 Golden Greats (1977)
Love Songs (1981)
Walking in the Light (1984, Gospel)
Private Collection: 1979–1988 (1988)
The Hit List (1994)
At the Movies 1959–1974 (1996)
The Whole Story: His Greatest Hits (2000)
The Platinum Collection (2005)
And They Said It Wouldn't Last (My 50 Years in Music) (2008)
The 50th Anniversary Album (2008)
75 at 75 (2015)
Stronger Thru the Years (2017)
Filmography
Serious Charge (1959)
Expresso Bongo (1959)
The Young Ones (Soundtrack album) (1961)
Summer Holiday (Soundtrack album) (1963)
Wonderful Life (Soundtrack album) (1964)
Finders Keepers (1966)
Two a Penny (1967)
Take Me High (1973)
Related articles
Discography
Concert tours
The Shadows
Book:Cliff Richard
show
v
t
e
Cliff Richard singles discography
1950s
"Move It" (1958)
"High Class Baby" (1958)
"Livin' Lovin' Doll" (1959)
"Mean Streak" (1959)
"Living Doll" (1959)
"Travellin' Light" (1959)
1960s
"Please Don't Tease" (1960)
"I Love You" (1960)
"A Girl Like You" (1961)
"When the Girl in Your Arms..." (1961)
"The Young Ones" / "We Say Yeah" (1962)
"I'm Lookin' Out the Window" / "Do You Want to Dance?" (1962)
"It'll Be Me" (1962)
"The Next Time" / "Bachelor Boy" (1962)
"Summer Holiday" (1963)
"Lucky Lips" (1963)
"It's All in the Game" (1963)
"I Only Have Eyes for You" (1964)
"Constantly (L'Edera)" (1964)
"The Twelfth of Never" (1964)
"The Minute You're Gone" (1965)
"Blue Turns to Grey" (1966)
"Visions" (1966)
"The Day I Met Marie" (1967)"
"All My Love" (1967)
"Congratulations" (1968)
"Girl, You'll Be a Woman Soon" (1968)
"Marianne" (1968)
"Early in the Morning" (1969)
1970s
"Sunny Honey Girl" (1971)
"Power to All Our Friends" (1973)
"Honky Tonk Angel" (1975)
"Miss You Nights" (1976)
"Devil Woman" (1976)
"Hey Mr. Dream Maker" (1976)
"My Kinda Life" (1977)
"When Two Worlds Drift Apart" (1977)
"Green Light" (1979)
"We Don't Talk Anymore" (1979)
1980s
"Carrie" (1980)
"Dreamin'" (1980)
"Suddenly" (1980)
"A Little in Love" (1980)
"Wired for Sound" (1981)
"Daddy's Home" (live) (1981)
"The Only Way Out" (1982)
"True Love Ways" (live) (1983)
"Baby You're Dynamite" / "Ocean Deep" (1984)
"Donna" (1984 US)
"Two to the Power of Love" (1984)
"Living Doll" (1986)
"All I Ask of You" (1986)
"My Pretty One" (1987)
"Some People" (1987)
"Remember Me" (1987)
"Two Hearts" (1988)
"Mistletoe and Wine" (1988)
"The Best of Me" (1989)
"I Just Don't Have the Heart" (1989)
"Lean On You" (1989)
"Whenever God Shines His Light" (1989)
1990s
"Stronger Than That" (1990)
"Silhouettes" (live) (1990)
"From a Distance" (live) (1990)
"Saviour's Day" (1990)
"Scarlet Ribbons" (1991)
"We Should Be Together" (1991)
"Peace in Our Time" (1993)
"All I Have to Do Is Dream" (live) / "Miss You Nights" (1994)
"Can't Keep this Feeling In" (1998)
"The Millennium Prayer" (1999)
2000s
"Somewhere Over the Rainbow"/"What a Wonderful World" (medley) (2001)
"Yesterday Once More" (2006)
"Move It" (2006)
"When I Need You" (2007)
"Mistletoe and Wine" (re-entry) (2008)
"Thank You for a Lifetime" (2008)
"Singing the Blues" (2009)
2010s
"Rip It Up" (2013)
"Since I Lost My Baby" (2014)
"Roll Over Beethoven" (2016)
"Blue Suede Shoes" (2017)
Book:Cliff Richard
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Balmer series
From Wikipedia, the free encyclopedia
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The "visible" hydrogen emission spectrum lines in the Balmer series. H-alpha is the red line at the right. Four lines (counting from the right) are formally in the "visible range." Lines five and six are easily seen with the naked eye but considered to be "ultraviolet" as they have wavelengths less than 400 nm.
The Balmer series or Balmer lines in atomic physics, is the designation of one of a set of six named series describing the spectral line emissions of the hydrogen atom. The Balmer series is calculated using the Balmer formula, an empirical equation discovered by Johann Balmer in 1885.
The visible spectrum of light from hydrogen displays four wavelengths, 410 nm, 434 nm, 486 nm, and 656 nm, that correspond to emissions of photons by electrons in excited states transitioning to the quantum level described by the principal quantum number n equals 2.[1] There are several prominent ultraviolet Balmer lines with wavelengths shorter than 400 nm. The number of these lines is an infinite continuum as it approaches a limit of 364.6 nm in the ultraviolet.
Contents
[hide]
1 Overview
2 Balmer's formula
3 Role in astronomy
4 See also
5 Notes
Overview[edit]
In the simplified Rutherford Bohr model of the hydrogen atom, the Balmer lines result from an electron jump between the second energy level closest to the nucleus, and those levels more distant. Shown here is a photon emission. The 3→2 transition depicted here produces H-alpha, the first line of the Balmer series. For hydrogen (Z = 1) this transition results in a photon of wavelength 656 nm (red).
The Balmer series is characterized by the electron transitioning from n ≥ 3 to n = 2, where n refers to the radial quantum number or principal quantum number of the electron. The transitions are named sequentially by Greek letter: n = 3 to n = 2 is called H-α, 4 to 2 is H-β, 5 to 2 is H-γ, and 6 to 2 is H-δ. As the first spectral lines associated with this series are located in the visible part of the electromagnetic spectrum, these lines are historically referred to as "H-alpha", "H-beta", "H-gamma" and so on, where H is the element hydrogen.
Transition of n
3→2 4→2 5→2 6→2 7→2 8→2 9→2 ∞→2
Name
H-α / Ba-α H-β / Ba-β H-γ / Ba-γ H-δ / Ba-δ H-ε / Ba-ε H-ζ / Ba-ζ H-η / Ba-η Balmer break
Wavelength (nm)[2]
656.3 486.1 434.1 410.2 397.0 388.9 383.5 364.6
Energy difference (eV)
1.89 2.55 2.86 3.03 3.13 3.19 3.23 3.40
Color
Red Aqua Blue Violet (Ultraviolet) (Ultraviolet) (Ultraviolet) (Ultraviolet)
Although physicists were aware of atomic emissions before 1885, they lacked a tool to accurately predict where the spectral lines should appear. The Balmer equation predicts the four visible absorption/emission lines of hydrogen with high accuracy. Balmer's equation inspired the Rydberg equation as a generalization of it, and this in turn led physicists to find the Lyman, Paschen, and Brackett series which predicted other absorption/emission lines of hydrogen found outside the visible spectrum.
The familiar red H-alpha spectral line of hydrogen gas, which is the transition from the shell n = 3 to the Balmer series shell n = 2, is one of the conspicuous colours of the universe. It contributes a bright red line to the spectra of emission or ionisation nebula, like the Orion Nebula, which are often H II regions found in star forming regions. In true-colour pictures, these nebula have a distinctly pink colour from the combination of visible Balmer lines that hydrogen emits.
Later, it was discovered that when the spectral lines of the hydrogen spectrum are examined at very high resolution, they are found to be closely spaced doublets. This splitting is called fine structure. It was also found that excited electrons could jump to the Balmer series n = 2 from orbitals where n was greater than 6, emitting shades of violet when doing so.
Two of the Balmer lines (α and β) are clearly visible in this emission spectrum of a deuterium lamp.
Balmer's formula[edit]
Balmer noticed that a single number had a relation to every line in the hydrogen spectrum that was in the visible light region. That number was 6993364506820000000♠364.50682 nm. When any integer higher than 2 was squared and then divided by itself squared minus 4, then that number multiplied by 6993364506820000000♠364.50682 nm (see equation below) gave a wavelength of another line in the hydrogen spectrum. By this formula, he was able to show that some measurements of lines made in his time by spectroscopy were slightly inaccurate and his formula predicted lines that were later found although had not yet been observed. His number also proved to be the limit of the series.
The Balmer equation could be used to find the wavelength of the absorption/emission lines and was originally presented as follows (save for a notation change to give Balmer's constant as B):
λ = B ( n 2 n 2 − m 2 ) = B ( n 2 n 2 − 2 2 ) {\displaystyle \lambda \ =B\left({\frac {n^{2}}{n^{2}-m^{2}}}\right)=B\left({\frac {n^{2}}{n^{2}-2^{2}}}\right)}
Where
λ is the wavelength.
B is a constant with the value of 6993364506820000000♠3.6450682×10−7 m or 6993364506820000000♠364.50682 nm.
m is equal to 2
n is an integer such that n > m.
In 1888 the physicist Johannes Rydberg generalized the Balmer equation for all transitions of hydrogen. The equation commonly used to calculate the Balmer series is a specific example of the Rydberg formula and follows as a simple reciprocal mathematical rearrangement of the formula above (conventionally using a notation of m for n as the single integral constant needed):
1 λ = 4 B ( 1 2 2 − 1 n 2 ) = R H ( 1 2 2 − 1 n 2 ) f o r n = 3 , 4 , 5 , . . . {\displaystyle {\frac {1}{\lambda }}={\frac {4}{B}}\left({\frac {1}{2^{2}}}-{\frac {1}{n^{2}}}\right)=R_{\mathrm {H} }\left({\frac {1}{2^{2}}}-{\frac {1}{n^{2}}}\right)\quad \mathrm {for~} n=3,4,5,...}
where λ is the wavelength of the absorbed/emitted light and RH is the Rydberg constant for hydrogen. The Rydberg constant is seen to be equal to 4/B in Balmer's formula, and this value, for an infinitely heavy nucleus, is 4/6993364506820000000♠3.6450682×10−7 m = 7007109737315700000♠10973731.57 m−1.[3]
Role in astronomy[edit]
The Balmer series is particularly useful in astronomy because the Balmer lines appear in numerous stellar objects due to the abundance of hydrogen in the universe, and therefore are commonly seen and relatively strong compared to lines from other elements.
The spectral classification of stars, which is primarily a determination of surface temperature, is based on the relative strength of spectral lines, and the Balmer series in particular is very important. Other characteristics of a star that can be determined by close analysis of its spectrum include surface gravity (related to physical size) and composition.
Because the Balmer lines are commonly seen in the spectra of various objects, they are often used to determine radial velocities due to doppler shifting of the Balmer lines. This has important uses all over astronomy, from detecting binary stars, exoplanets, compact objects such as neutron stars and black holes (by the motion of hydrogen in accretion disks around them), identifying groups of objects with similar motions and presumably origins (moving groups, star clusters, galaxy clusters, and debris from collisions), determining distances (actually redshifts) of galaxies or quasars, and identifying unfamiliar objects by analysis of their spectrum.
Balmer lines can appear as absorption or emission lines in a spectrum, depending on the nature of the object observed. In stars, the Balmer lines are usually seen in absorption, and they are "strongest" in stars with a surface temperature of about 10,000 kelvin (spectral type A). In the spectra of most spiral and irregular galaxies, AGNs, H II regions and planetary nebulae, the Balmer lines are emission lines.
In stellar spectra, the H-epsilon line (transition 7→2) is often mixed in with another absorption line caused by ionized calcium known by astronomers as "H" (the original designation given by Fraunhofer). That is, H-epsilon's wavelength is quite close to CaH at 396.847 nm, and cannot be resolved in low resolution spectra. The H-zeta line (transition 8→2) is similarly mixed in with a neutral helium line seen in hot stars.
See also[edit]
Astronomical spectroscopy
Bohr model
Hydrogen spectral series
Lyman series
Rydberg formula
Stellar classification
Theoretical and experimental justification for the Schrödinger equation
Notes[edit]
Jump up ^ Nave, C. R. (2006). "Hydrogen Spectrum". HyperPhysics. Georgia State University. Retrieved March 1, 2008.
Jump up ^ Eisberg and Resnick (1985). Quantum Physics. John Wiley and Sons. p. 97.
Jump up ^ "CODATA Recommended Values of the Fundamental Physical Constants: 2006" (PDF). Committee on Data for Science and Technology (CODATA). NIST.
[hide]
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Lesley Fera
From Wikipedia, the free encyclopedia
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Lesley Fera
Born
(1971-11-23) November 23, 1971 (age 46) [1]
Nationality
American
Occupation
Actress
Years active
1998–present
Lesley Fera (born November 23, 1971) is an American actress. She is best known for her role as Veronica Hastings on the ABC Family series Pretty Little Liars.[2] Fera also had a recurring roles in 24, CSI: Miami and Southland, as well as performing in a number of theater productions.[2]
Contents
[hide]
1 Biography
2 Filmography
2.1 Film
2.2 Television
2.3 Stage
3 References
4 External links
Biography[edit]
A native Californian, Lesley is a lover of the stage, she is very active in L.A. theatre and has also enjoyed working at numerous regional theaters. Played the role of Veronica Hastings (PLL).
Filmography[edit]
Film[edit]
Year
Title
Role
Notes
2005 Herstory Marly
2017 The Lovers Susan
2017 Boone: The Bounty Hunter Madam Sylvia
Television[edit]
Year
Title
Role
Notes
1998 3rd Rock from the Sun Woman #2 "The Great Dickdater" (season 3: episode 13)
1999 Absence of the Good Beth Television film
2000 Titus Paula "Sex with Pudding" (season 1: episode 1)
2001 Gideon's Crossing Janet Beckman "The Crash" (season 1: episode 19)
2001 Strong Medicine Anne "Impaired" (season 2: episode 7)
2001 Thieves Carla Patterson "Home Is Where the Heist Is…" (season 1: episode 10)
2001 The Drew Carey Show Judy "Bus-ted" (season 7: episode 6)
2001 Citizen Baines Eileen McGuire "Out in the Rain" (season 1: episode 7)
2002 Judging Amy Mrs. Moran "Picture of Perfect" (season 4: episode 21)
2003 The Practice A.D.A. Kate Barron "We the People" (season 8: episode 1)
2004–10 CSI: Miami Dr. Joyce Carmel 4 episodes
Guest star (Seasons 2, 4, 5, 9)
2004 ER Helen Kolber "NICU" (season 10: episode 12)
2004 Medical Investigation Gladys Halperin "Coming Home" (season 1: episode 3)
2005 NYPD Blue Debra Campo "Sergeant Sipowicz' Lonely Hearts Club Band" (season:12 episode 17)
2006 Numb3rs ATF Agent Annie Wilson "Guns and Roses" (season 2: episode 20)
2006 Cold Case Dana Taylor '04 "The War at Home" (season 4: episode 2)
2006–07 Justice DA Jennifer Reese 2 episodes
2009–10 24 Angela Nelson 6 episodes
Recurring role (Season 7) Guest Star (Season 8)
2009 Without a Trace Cynthia "Heartbeats" (season 7: episode 19)
2009 The Mentalist Leslie Sloop "Throwing Fire" (season 2: episode 10)
2010 Nip/Tuck Daniella Creighton "Christian Troy II" (season 6: episode 17)
2010 Criminal Minds Leslie McBride "The Fight" (season 5: episode 18)
2010–17 Pretty Little Liars Veronica Hastings 53 episodes;
Recurring role (Seasons 1–7)
2010 Detroit 1-8-7 Dr. Maya Ehrlich "Shelter" (Season 1: episode 10)
2011 House M.D. Kay "Bombshells" (Season 7: episode 15)
2013 Southland Sergeant Waters 4 episodes
Recurring role (Season 5)
2013 Dating in LA and Other Urban Myths Penelope 9 episodes
Main role (Season 1)
2014 Stalker Mrs. McCoy "Whatever Happened to Baby James?" (Season 1: episode 2)
2017 Battle of the Network Stars Herself Contestant
Stage[edit]
Year
Title
Role
Location
1995–96 Noises Off Poppy Asolo Repertory Theatre
1996 The Life and Adventures of Nicholas Nickleby Miss Knag (et al)
1997 The Children’s Hour Martha Provincetown Theater Foundation
1997 The Mousetrap Mollie Pennsylvania Centre Stage
1997 Misalliance Hypatia
1998 A Sense of Place Mary Tate American Stage
1999 Communicating Doors Poopay / Phoebe San Jose Repertory Theatre
2000–01 Lady Chatterley’s Lover Lady Chatterley Pacific Resident Theatre
2002 Anna Christie Anna Christie
2002–03 Big Love Olympia
2004 Dealing With Clair The Matrix Theatre Company
2005 Happy End Lillian Holiday Pacific Resident Theatre
2007 The Hasty Heart Margaret
2011 Barrie: Back to Back Mrs. Beatrice Page
2012 Our Town The Eli and Edythe Broad Stage
2015 The Homecoming Ruth Pacific Resident Theatre
2016 Two Dollar Bill Jessica McGovern Pioneer Theatre Company
2016 Bull Isobel Rogue Machine Theatre
2016 Two Henrys Pacific Resident Theatre
2017 Dinner with Friends
References[edit]
Jump up ^ "My birthday cake! :)". Twitter. Retrieved 2015-08-03.
^ Jump up to: a b "Veronica Hastings by Lesley Fera - Pretty Little Liars - ABCFamily.com". Abcfamily.go.com. Retrieved 2014-06-03.
External links[edit]
Lesley Fera on IMDb

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Pretty Little Liars (TV series)

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