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Acid strength
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The strength of an acid refers to its ability or tendency to lose a proton (H+). A strong acid is one that completely ionizes (dissociates) in a solution (provided there is sufficient solvent). In water, one mole of a strong acid HA dissolves yielding one mole of H+ (as hydronium ion H3O+ and higher aggregates) and one mole of the conjugate base, A−. Essentially, none of the non-ionized acid HA remains. Examples of strong acids are hydrochloric acid (HCl), hydroiodic acid (HI), hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3) and sulfuric acid (H2SO4). In aqueous solution, each of these essentially ionizes about 100%
Image of a strong acid mostly dissociating. The small red circles represent H+ ions.
In contrast, a weak acid only partially dissociates. Examples in water include carbonic acid (H2CO3) and acetic acid (CH3COOH). At equilibrium, both the acid and the conjugate base are present in solution.
Image of a weak acid partly dissociating
Stronger acids have a larger acid dissociation constant (Ka) and a smaller logarithmic constant (pKa = −log Ka) than weaker acids. The stronger an acid is, the more easily it loses a proton, H+. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base.
While the pKa value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water and DMSO), the tendency of an acidic solvent to transfer a proton to a reference solute (most commonly a weak aniline base) is measured by its Hammett acidity function, the H0 value. Although these two concepts of acid strength often amount to the same general tendency of a substance to donate a proton, the pKa and H0 values are measures of distinct properties and may occasionally diverge. For instance, hydrogen fluoride, whether dissolved in water (pKa = 3.2) or DMSO (pKa = 15), has pKa values indicating that it undergoes incomplete dissociation in these solvents, making it a weak acid. However, as the rigorously dried, neat acidic medium, hydrogen fluoride has an H0 value of –15, making it a more strongly protonating medium than 100% sulfuric acid and thus, by definition, a superacid.[1] (To prevent ambiguity, in the rest of this article, "strong acid" will, unless otherwise stated, refer to an acid that is strong as measured by its pKa value (pKa < –1.74). This usage is consistent with the common parlance of most practicing chemists.)
When the acidic medium in question is a dilute aqueous solution, the H0 is approximately equal to the pH value, which is a negative logarithm of the concentration of aqueous H+ in solution. The pH of a simple solution of an acid in water is determined by both Ka and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, which may be determined by an equilibrium calculation. For concentrated solutions of acids, especially strong acids for which pH < 0, the H0 value is a better measure of acidity than the pH.
Sulfonic acids, which are organic oxyacids, are a class of strong acids. A common example is p-toluenesulfonic acid (tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable.
Superacids are acid solutions that are more acidic than 100% sulfuric acid.[2] Examples of superacids are fluoroantimonic acid, magic acid and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations.
Contents
[hide]
1 Strong acids in water
1.1 Determining acid strength
1.2 Common strong acids
1.3 Almost strong acids
1.4 Extremely strong acids (as protonators)
2 Weak acids in water
2.1 Dissociation
2.2 Calculating the pH of a weak acid solution
2.2.1 Simplification
2.2.2 Comparison of the full and simplified methods
2.3 Conjugate acid/base pair
3 Factors determining acid strength
3.1 Polarity and the inductive effect
3.2 Atomic radius and bond strength
4 Acids in nonaqueous solvents
5 Corrosivity
6 References
7 External links
Strong acids in water[edit]
In aqueous solution, a strong acid is an acid that ionizes completely by losing one proton. The proton lost is captured by a water molecule to form a hydronium ion:
HA(aq) + H2O → H3O+(aq) + A−(aq)
For sulfuric acid, which is diprotic, the "strong acid" designation refers only to dissociation of the first proton
H2SO4(aq) → H+(aq) + HSO4−(aq)
The acid must be stronger in aqueous solution than hydronium ions, so strong acids are acids with a pKa < −1.74. An example is HCl for which pKa = −6.3.[3] This generally means that, in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution.
Due to the complete dissociation of strong acids in aqueous solution, the concentration of hydronium ions in the water is equal to the total concentration (ionized and un-ionized) of the acid introduced to solution: [H+] = [A−] = [HA]total and pH = −log[H+]. No acid species stronger than H3O+(aq) can exist in aqueous solution, as any stronger acid is converted to H3O+. The acidity of the stronger acid is said to be leveled to the acidity of hydronium ion.[4]
Strong acids are distinguished from weak acids, in which dissociation is incomplete and is represented as an equilibrium, not a completed reaction. The typical definition of a weak acid is any acid that does not dissociate completely. The large difference between the acid dissociation constants of strong acids and those of all other acids makes this demarcation possible.
Determining acid strength[edit]
The strength of an acid, in comparison to other acids, depends on several characteristics:
Electronegativity: The higher the electronegativity of a conjugate base in the same period, the more acidic. In other words, the more electronegative A- is, more acidic (where HA → H+ + A−).
Atomic Radius: With increasing atomic radius, acidity also increases. For example, HCl and HI, both strong acids, ionize 100% in water to become their respective ionic constituents. However, HI is stronger than HCl. This is because the atomic radius of an atom of iodine is much larger than that of a chlorine atom. As a result, the negative charge over the I− anion is dispersed over a larger electron cloud and its attraction for the proton (H+) is not as strong as the same attraction in HCl. Therefore, HI is ionized (deprotonated) more readily.
Charge: The more positively charged a species is, the more acidic (neutral molecules can be stripped of protons more easily than anions, and cations are more acidic than comparable molecules).
Equilibrium: The strength of an acid can be defined by the equilibrium position of its dissociation reaction:
HA(aq)+ H2O(l) → H3O+(aq) + A−(aq)
In a strong acid, equilibrium lies far to the right, meaning that almost all of the original HA is dissociated at equilibrium. A strong acid yields a weak conjugate base (A−), so a strong acid is also described as an acid whose conjugate base is a much weaker base than water.[5]
Common strong acids[edit]
This is a list of strong acids with pKa < −1.74, which is stronger than the hydronium ion, from strongest to weakest.
Perchloric acid HClO4 (pKa ≈ −10)[6]
Hydroiodic acid HI (pKa = −9.3)[3]
Hydrobromic acid HBr (pKa = −8.7)[3]
Hydrochloric acid HCl (pKa = −6.3)[3]
Sulfuric acid H2SO4 (first dissociation only, pKa1 ≈ −3)[7]
p-Toluenesulfonic acid (pKa = −2.8) Organic soluble strong acid
Methanesulfonic acid (pKa = −1.92) Liquid organic strong acid[8]
Almost strong acids[edit]
These do not meet the strict criterion of being more acidic than H3O+, although in very dilute solution they dissociate almost completely, so sometimes they are included as "strong acids"
Hydronium ion H3O+ (pKa = −1.74). Hydronium is often used as an approximation of the state of protons in water.
Nitric acid HNO3 (pKa = −1.64)[7]
Chloric acid HClO3 (pKa = −1.0)[7]
Some chemists include bromic acid (HBrO3),[9] perbromic acid (HBrO4),[9] iodic acid (HIO3),[9] and periodic acid (HIO4)[9] as strong acids, although these are not universally accepted as such.
Extremely strong acids (as protonators)[edit]
(Strongest to weakest)
Fluoroantimonic acid H[SbF6]
Magic acid FSO3HSbF5
Carborane superacid H(CHB11Cl11)
Fluorosulfuric acid FSO3H (pKa = −6.4)[8]
Triflic acid CF3SO3H (pKa = −5.9)[8]
Weak acids in water[edit]
Most acids are weak acids. A weak acid is an acid that dissociates incompletely, releasing only some of its hydrogen atoms into the solution. Thus, it is less capable than a strong acid at donating protons. These acids have higher pKa than strong acids, which release all of their hydrogen atoms when dissolved in water. Examples of weak acids include acetic acid (CH3COOH), phosphoric acid (H3PO4), hydrofluoric acid (HF) and oxalic acid (C2H2O4).
Dissociation[edit]
Weak acids ionize in water solution to only a moderate extent; that is, if the acid was represented by the general formula HA, then in aqueous solution a significant amount of undissociated HA still remains. Weak acids in water dissociate as:
HA ( aq ) ↽ − − ⇀ H ( aq ) + + A ( aq ) − {\displaystyle {\ce {HA_{(aq)}<=>{H+_{(aq)}}+A_{(aq)}^{-}}}}
The strength of a weak acid is represented as either an equilibrium constant or as a percent dissociation. The equilibrium concentrations of reactants and products are related by the acid dissociation constant expression, (Ka):
K a = [ H + ] [ A − ] [ HA ] {\displaystyle K_{a}={\frac {{\ce {[H+][A^-]}}}{{\ce {[HA]}}}}}
The greater the value of Ka, the more the formation of H+ is favored, and the lower the pH of the solution. The Ka of weak acids varies between 1.8×10−16 and 55.5. Acids with a Ka less than 1.8×10−16 are weaker acids than water.
The other way to measure acid strength is to look at its fractional dissociation, which is symbolized as α (alpha) and which can range from 0% < α < 100%. The dissociation ratio is defined as
α = [ A − ] [ A − ] + [ HA ] {\displaystyle {\ce {\alpha = {\frac {[ A^- ]}{[ A^- ] + [ HA ]}}}}}
Unlike Ka, α is not constant and does depend on the [HA]. In general, α will increase as [HA] decreases. Thus acids become stronger as they are diluted. If acids are polyprotic, then each proton will have a Ka. For example: H2CO3 + H2O → HCO3− + H3O+ has two Ka values because it has two acidic protons. The first Ka value is 4.46×10−7 (pKa1 = 6.351) and the second is 4.69×10−11 (pKa2 = 10.329).
Calculating the pH of a weak acid solution[edit]
The pH of a solution of a weak acid depends on the strength of the acid and the other components in the solution. In the simplest case, the weak acid is the only compound in water. In this case, the pH can be found from the concentration of the acid (symbolized as F {\displaystyle F} ), from the Acid dissociation constant (symbolized as K a {\displaystyle K_{a}} ), and by solving for concentration of H+ (symbolized by x and represented more accurately as H3O+). Below is a table that organizes the information. On the first line, the balanced reaction is written. On the second line, the initial conditions are written below each compound. Note that a value of water is not given because its term (activity) in the K a {\displaystyle K_{a}} expression is technically equal to 1, but is often (conveniently) omitted. The third line shows how the value changes as the reaction goes to equilibrium. Then the last line gives the equilibrium concentrations and is simply the sum of each column.
HA(aq) + H2O(l) → A−(aq) + H3O+(aq)
initial F — 0 0
change -x — +x +x
equilibrium F - x — x x
Applying the equilibrium line to the K a {\displaystyle K_{a}} expression yields
K a = [ H + ] [ A − ] [ HA ] = x 2 F − x {\displaystyle K_{a}={\ce {{[H+][A^{-}]} \over {[HA]}}}={{x^{2}} \over {F-x}}}
rearranging yields x 2 + K a x − K a F = 0 {\displaystyle x^{2}+K_{a}x-K_{a}F=0} , which can be solved for x using the quadratic equation. The pH is then calculated as pH = − log  ( x ) {\displaystyle {{\ce {pH}}}=-\log(x)} .
Simplification[edit]
However, if F is more than a thousand times Ka, then (1) the acid will not deprotonate much, (2) the value of x will be small, and therefore (3) F - x ≈ F. This is also known as the 5% rule. This simplifies the Ka expression to...
K a = x 2 F {\displaystyle K_{a}={{x^{2}} \over {F}}}
Solving for x yields
x = K a F = [ H + ] {\displaystyle x={\sqrt {K_{a}F}}=\left[{{\ce {H+}}}\right]}
Then the pH = − log  [ H + ] {\displaystyle {{\ce {pH}}}=-\log[{{\ce {H+}}}]} . The following equation then follows, but is only true if F >>> Ka
pH = − log  K a F {\displaystyle {{\ce {pH}}}=-\log {\sqrt {K_{a}F}}}
Comparison of the full and simplified methods[edit]
A certain weak acid has a Ka = 1×10−5 and the pH of two solutions needs to be found. One solution has a concentration of 0.10M and another has a concentration of 5×10−4M. The pH for both solutions will be calculated using both methods to yield 4 values, which will be compared.
0.1M Solution
The full method gives the following quadratic:
x 2 + 1 × 10 − 5 x − 1 × 10 − 6 = 0 {\displaystyle x^{2}+1\times 10^{-5}x-1\times 10^{-6}=0}
which gives x = 9.95×10−4 M and a pH = 3.00. The simplified method gives
pH = − log  10 − 5 ⋅ 0.1 = 3.00 {\displaystyle {{\ce {pH}}}=-\log {\sqrt {10^{-5}\cdot 0.1}}=3.00}
So both methods yield the same result, but again F is more than a 1000 times Ka. The next case does not have this condition and the results will differ.
5×10−4M Solution
The full method gives the following quadratic:
x 2 + 10 − 5 x − 5 × 10 − 9 = 0 {\displaystyle x^{2}+10^{-5}x-5\times 10^{-9}=0}
which gives x = 6.6×10−5 M and a pH = 4.18. The simplified method gives
pH = − log  10 − 5 ( 5 × 10 − 4 ) = 4.15 {\displaystyle {{\ce {pH}}}=-\log {\sqrt {10^{-5}\left(5\times 10^{-4}\right)}}=4.15}
Here, the results differ by 0.03 pH units. As F becomes closer in value to the Ka, then the difference will increase even more. However, in practice, it is rare to work with such dilute acids and the pH is also dependent on ionic strength and temperature. So in reality, the simplified method works well.
Conjugate acid/base pair[edit]
It is often stated that "the conjugate of a weak acid is a strong base". This statement can be misleading. Most weak acids that textbooks discuss have weak (not strong) conjugate bases. Truly, only the very weakest of acids have strong conjugate bases. For example, if a weak acid has a Ka = 10−5, then its conjugate base would have a Kb = 10−9 (from the relationship Ka × Kb = 10−14), which certainly is not a strong base. A very weak acid with a Ka = 10−20 would indeed have a strong conjugate base.
Factors determining acid strength[edit]
Polarity and the inductive effect[edit]
Polarity refers to the distribution of electrons in a bond, the region of space between two atomic nuclei where a pair of electrons is shared. When two atoms have roughly the same electronegativity (ability to attract electrons) the electrons are shared evenly and spend equal time on either end of the bond. When there is a significant difference in electronegativities of two bonded atoms, the electrons spend more time near the nucleus of the more electronegative element and an electrical dipole, or separation of charges, occurs, such that there is a partial negative charge localized on the electronegative element and a partial positive charge on the electropositive element. Hydrogen is an electropositive element and accumulates a slightly positive charge when it is bonded to an electronegative element such as oxygen or bromine. As the electron density on hydrogen decreases, it is more easily abstracted, in other words, it is more acidic. Moving from left to right across a row on the periodic table elements become more electronegative (excluding the noble gases), and the strength of the binary acid formed by the element increases accordingly:
Formula
Name
pKa[10]
HF hydrofluoric acid 3.17
H2O water 15.7
NH3 ammonia 38
CH4 methane 48
The electronegative element need not be directly bonded to the acidic hydrogen to increase its acidity. An electronegative atom can pull electron density out of an acidic bond through the inductive effect. The electron-withdrawing ability diminishes quickly as the electronegative atom moves away from the acidic bond. The effect is illustrated by the following series of halogenated butanoic acids. Chlorine is more electronegative than bromine and therefore has a stronger effect. The hydrogen atom bonded to the oxygen is the acidic hydrogen. Butanoic acid is a carboxylic acid.
Structure
Name
pKa[11]
butanoic acid or butyric acid ≈4.8
4-chlorobutanoic acid 4.5
3-chlorobutanoic acid ≈4.0
2-bromobutanoic acid 2.93
2-chlorobutanoic acid 2.86
As the chlorine atom moves further away from the acidic O—H bond, its effect diminishes. When the chlorine atom is just one carbon removed from the carboxylic acid group, the acidity of the compound increases significantly compared to butanoic acid (a.k.a. butyric acid). However, when the chlorine atom is separated by several bonds, the effect is much smaller. Bromine is much more electronegative than either carbon or hydrogen, but not as electronegative as chlorine, so the pKa of 2-bromobutanoic acid is slightly greater than the pKa of 2-chlorobutanoic acid.
Perchloric acid (HClO4) is an oxoacid and a strong acid.
The number of electronegative atoms adjacent to an acidic bond also affects acid strength. Oxoacids have the general formula HOX, where X can be any atom and may or may not share bonds to other atoms. Increasing the number of electronegative atoms or groups on atom X decreases the electron density in the acidic bond, making the loss of the proton easier. Perchloric acid is a very strong acid (pKa ≈ -8) and completely dissociates in water. Its chemical formula is HClO4 and it comprises a central chlorine atom with three chlorine-oxygen double bonds (Cl=O) and one chlorine-oxygen single bond (Cl—O). The singly bonded oxygen bears an extremely acidic hydrogen atom, which is easily abstracted. In contrast, chloric acid (HClO3) is a weaker acid, though still quite strong (pKa = -1.0), while chlorous acid (HClO2, pKa = +2.0) and hypochlorous acid (HClO, pKa = +7.53) acids are weak acids.[12]
Carboxylic acids are organic acids that contain an acidic hydroxyl group and a carbonyl (C=O bond). Carboxylic acids can be reduced to the corresponding alcohol; the replacement of an electronegative oxygen atom with two electropositive hydrogens yields a product which is essentially non-acidic. The reduction of acetic acid to ethanol using LiAlH4 (lithium aluminium hydride or LAH) and ether is an example of such a reaction.
The pKa for ethanol is 16, compared to 4.76 for acetic acid.[11][13]
Atomic radius and bond strength[edit]
Another factor that contributes to the ability of an acid to lose a proton is the strength of the bond between the acidic hydrogen and the atom that bears it. This, in turn, is dependent on the size of the atoms sharing the bond. For an acid HA, as the size of atom A increases, the strength of the bond decreases, meaning that it is more easily broken, and the strength of the acid increases. Bond strength is a measure of how much energy it takes to break a bond. In other words, it takes less energy to break the bond as atom A grows larger, and the proton is more easily removed by a base. This partially explains why hydrofluoric acid is considered a weak acid, while the other hydrohalic acids (HCl, HBr, HI) are strong acids. Although fluorine is more electronegative than the other halogens, its atomic radius is also much smaller, so it shares a stronger bond with hydrogen. Moving down a column on the periodic table, atoms become less electronegative but also significantly larger, and the size of the atom tends to dominate its acidity when sharing a bond to hydrogen. Hydrogen sulfide, H2S, is a stronger acid than water, even though oxygen is more electronegative than sulfur. Just as with the halogens, this is because sulfur is larger than oxygen and the H—S bond is more easily broken than the H—O bond.
Acids in nonaqueous solvents[edit]
The strength of an acid also depends on the solvent basicity (tendency to accept a proton), so that acids which are strong in water may be weak in less basic solvents, and acids which are weak in water may be strong in more basic solvents. According to Brønsted–Lowry acid–base theory, the acidity of an acid HA in a solvent S is due to the acid-base reaction HA + S → A− + HS+.
For example, the solvent glacial (anhydrous) acetic acid is less basic than water, and the extent of ionization of the hydrohalic acids varies in the order HI > HBr > HCl. These acids lose a proton to a CH3COOH molecule which forms its conjugate acid, CH3C(OH)2+, but the reaction is only partial. In contrast these three acids are all strong (completely ionized) in water. Acetic acid is said to be a differentiating solvent for the three solute acids, while water is not.[14]
An example of a solvent more basic than water is liquid ammonia, NH3, (at low temperature), in which acids react to form the ammonium ion.
HA + NH 3 ⟶ A − + NH 4 + {\displaystyle {\ce {{HA}+NH3->{A^{-}}+NH4+}}}
In this solvent, acetic acid as solute is a strong acid and ionizes completely, and has the same strength as HCl (at the same concentration. The acidities of both acetic acid and HCl are leveled to that of NH4+ which is the strongest acid that exists in liquid ammonia solution.[15]
Corrosivity[edit]
While strong acids are generally assumed to be the most corrosive, this is not always true. The carborane superacid H(CHB11Cl11), which is thousand times stronger than the strength of sulfuric acid,[16][17] is entirely non-corrosive, whereas the weak acid hydrofluoric acid (HF) is corrosive and can dissolve, among other things, glass[18] and most metals.
References[edit]
Jump up ^ Liang, Joan-Nan Jack (1976). The Hammett Acidity Function for Hydrofluoric Acid and some related Superacid Systems (Ph.D. Thesis) (PDF). Hamilton, Ontario: McMaster University. p. 109.
Jump up ^ Miessler G.L. and Tarr D.A. Inorganic Chemistry (2nd ed., Prentice-Hall 1998, p.170) ISBN 0-13-841891-8
^ Jump up to: a b c d William L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill, 1984), p.177
Jump up ^ Porterfield, William W. Inorganic Chemistry (Addison-Wesley 1984) p.260 ISBN 0-201-05660-7
Jump up ^ Zumdahl, Steven S. (2011). Chemical Principles: Enhanced Edition, (6th ed.). Brooks/Cole Cengage Learning. p. 236.
Jump up ^ Kathleen Sellers; Katherine Weeks; William R. Alsop; Stephen R. Clough; Marilyn Hoyt; Barbara Pugh (2006). Perchlorate: environmental problems and solutions. CRC Press. p. 16. ISBN 0-8493-8081-2.
^ Jump up to: a b c Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 171. ISBN 978-0130399137.
^ Jump up to: a b c Guthrie, J. P. Hydrolysis of esters of oxy acids: pKa values for strong acids. Can. J. Chem. 1978, 56, 2342-2354.
^ Jump up to: a b c d Chieh, Chung. "Strong Acids and Bases". University of Waterloo. Retrieved 27 May 2015.
Jump up ^ pKa's of Inorganic and Oxo-Acids
^ Jump up to: a b Section 8: Electrolytes, Electromotive forces and Chemical Equilibrium
Jump up ^ pKa values for HClOn from Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. ISBN 978-0130399137.
Jump up ^ pKa Data Compiled by R. Williams Archived 2010-06-02 at the Wayback Machine.
Jump up ^ Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 217. ISBN 978-0130399137.
Jump up ^ Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 219. ISBN 978-0130399137.
Jump up ^ George A. Olah, et al. Superacid Chemistry, 2nd ed., Wiley, p. 41.
Jump up ^ That is, the ability of the carborane superacid to protonate a given base (B) is one million times that of a solution of sulfuric acid, so that the ratio [BH+] / [B] is one million times higher. The relative acidities of strong acids can be evaluated using the Hammett acidity function.
Jump up ^ CID 14917 from PubChem
General
Hill, John W., et al. "General Chemistry." 4th ed. New Jersey: Prentice Hall, 2005.
External links[edit]
https://web.archive.org/web/20031027165745/http://www.cm.utexas.edu/academic/courses/Spring2002/CH301/McDevitt/strong.htm
http://jchemed.chem.wisc.edu/Journal/Issues/2000/Jul/abs849.html
Titration of acids - freeware for data analysis and simulation of potentiometric titration curves
Acids and Bases - definitions
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End-systolic volume
From Wikipedia, the free encyclopedia
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End-systolic volume (ESV) is the volume of blood in a ventricle at the end of contraction, or systole, and the beginning of filling, or diastole.
ESV is the lowest volume of blood in the ventricle at any point in the cardiac cycle. The main factors that affect the end-systolic volume are afterload and the contractility of the heart.
Uses[edit]
End systolic volume can be used clinically as a measurement of the adequacy of cardiac emptying, related to systolic function. On an electrocardiogram, or ECG, the end-systolic volume will be seen at the end of the T wave. Clinically, ESV can be measured using two-dimensional echocardiography, MRI (magnetic resonance tomography) or cardiac CT (computed tomography)or SPECT (single photon emission computed tomography).
Sample values[edit]
Along with end-diastolic volume, ESV determines the stroke volume, or output of blood by the heart during a single phase of the cardiac cycle.[1] The stroke volume is the difference between the end-diastolic volume and the end-systolic volume. The end-systolic values in the table below are for the left ventricle:
Ventricular volumes
view
talk
edit
Measure
Right ventricle
Left ventricle
End-diastolic volume 144 mL(± 23mL)[2] 142 mL (± 21 mL)[3]
End-diastolic volume / body surface area (mL/m2) 78 mL/m2 (± 11 mL/m2)[2] 78 mL/m2 (± 8.8 mL/m2)[3]
End-systolic volume 50 mL (± 14 mL)[2] 47 mL (± 10 mL)[3]
End-systolic volume / body surface area (mL/m2) 27 mL/m2 (± 7 mL/m2)[2] 26 mL/m2 (± 5.1 mL/m2)[3]
Stroke volume 94 mL (± 15 mL)[2] 95 mL (± 14 mL)[3]
Stroke volume / body surface area (mL/m2) 51 mL/m2 (± 7 mL/m2)[2] 52 mL/m2 (± 6.2 mL/m2)[3]
Ejection fraction 66% (± 6%)[2] 67% (± 4.6%)[3]
Heart rate 60–100 bpm[4] 60–100 bpm[5]
Cardiac output 4.0–8.0 L/minute[6] 4.0–8.0 L/minute[6]
The right ventricular end-systolic volume (RVESV) normally ranges between 50 and 100 mL.[6]
References[edit]
Jump up ^ Boron and Boulpaep 2005 Medical Physiology Updated Edition p. 521 ISBN 0-7216-3256-4
^ Jump up to: a b c d e f g Maceira, Alicia (2006). "Reference right ventricular systolic and diastolic function normalized to age, gender and body surface area from steady-state free precession cardiovascular magnetic resonance" (PDF). European Heart Journal. 27: 2879–2888. PMID 17088316. doi:10.1093/eurheartj/ehl336. Retrieved 6 November 2015.
^ Jump up to: a b c d e f g Maceira, Alicia (2006). "Normalized Left Ventricular Systolic and Diastolic Function by Steady State Free Precession Cardiovascular Magnetic Resonance". Journal of Cardiovascular Magnetic Resonance. 8: 417–426. doi:10.1080/10976640600572889. Retrieved 6 November 2015.
Jump up ^ Normal ranges for heart rate are among the narrowest limits between bradycardia and tachycardia. See the Bradycardia and Tachycardia articles for more detailed limits.
Jump up ^ Normal ranges for heart rate are among the narrowest limits between bradycardia and tachycardia. See the Bradycardia and Tachycardia articles for more detailed limits.
^ Jump up to: a b c Edwards Lifesciences LLC > Normal Hemodynamic Parameters – Adult 2009
[hide]
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Physiology of the cardiovascular system
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Heart rate
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Stroke volume
End-diastolic volume
End-systolic volume
Afterload
Preload
Frank–Starling law
Cardiac function curve
Venous return curve
Wiggers diagram
Pressure volume diagram
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Fractional shortening = (End-diastolic dimension
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Flag of Trinidad and Tobago
From Wikipedia, the free encyclopedia
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Trinidad and Tobago
Name

Flag of the Republic of Trinidad and Tobago
The Sun-Sea-Sand Banner
Use
National flag
Proportion
3:5
Adopted
31 August 1962
Design
A red field with a white-edged black diagonal band from the upper hoist-side to the lower fly-side.
Variant flag of Trinidad and Tobago
Name
Civil ensign of the Republic of Trinidad and Tobago
Use
Civil and state ensign
Proportion
1:2
Design
A red field with a white-edged black diagonal band from the upper hoist-side to the lower fly-side.
Variant flag of Trinidad and Tobago
Name
Naval ensign of the Republic of Trinidad and Tobago
Use
Naval ensign
Proportion
1:2
The flag of Trinidad and Tobago was adopted upon independence from the United Kingdom on 31 August 1962. Designed by Carlisle Chang (1921–2001),[1][2][3] the flag of Trinidad and Tobago was chosen by the independence committee of 1962. Red, black and white symbolise fire (the sun, representing courage), earth (representing dedication) and water (representing purity and equality).[4]
Contents
[hide]
1 Design
2 Construction
3 Other flags
3.1 British colonial flag
4 See also
5 References
Design[edit]
Trinidadian flags flying at the University of the West Indies in Saint Augustine, Trinidad and Tobago
The flag of Trinidad and Tobago is a red field with a white-edged black diagonal band from the upper hoist-side to the lower fly-side. In blazon, Gules, a bend Sable fimbriated Argent.
Construction[edit]
The flag of Trinidad and Tobago flying at the San Fernando Hill, San Fernando in July 2009.
The width of the white stripes is 1⁄30 of the flag length and the width of the black stripe is 2⁄15. The total width of the three stripes together is therefore 1⁄5 of the length.[5]
Other flags[edit]
The civil ensign is the national flag in a 1:2 ratio. The naval ensign (used by coast guard vessels) is a British white ensign with the national flag in the canton.
Presidential Standard
Prime Minister's Standard
Queen Elizabeth II's personal Trinidadian flag (1966–1976)
British colonial flag[edit]
Prior to independence from the United Kingdom in August 1962, Trinidad and Tobago used a British blue ensign defaced with a circular badge depicting a ship arriving in front of a mountain.
Colonial flag of Trinidad and Tobago (1889-1962)
Red ensign of Trinidad and Tobago (1889-1962)
See also[edit]
Flag of the West Indies Federation (January 1958 – May 1962)
Coat of arms of Trinidad and Tobago
References[edit]
Jump up ^ "Carlisle Chang" Archived 2016-09-11 at the Wayback Machine., Ministry of Community Development, Culture and the Arts, Government of the Republic of Trinidad and Tobago.
Jump up ^ "National flag of Trinidad and Tobago - Carlisle Chang", YouTube.
Jump up ^ Carlisle Chang biography Archived 2015-07-11 at the Wayback Machine., NALIS.
Jump up ^ National Flag". Government of the Republic of Trinidad and Tobago. Retrieved May 12, 2017.
Jump up ^ Flagspot
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The Master of Disguise
From Wikipedia, the free encyclopedia
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"Master of Disguise" redirects here. For the album, see Master of Disguise (album). For the game, see Wario: Master of Disguise. For the book, see Tony Mendez.
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The Master of Disguise
Theatrical release poster
Directed by
Perry Andelin Blake
Produced by
Barry Bernardi
Sid Ganis
Todd Garner
Adam Sandler
Alex Siskin
Written by
Dana Carvey
Harris Goldberg
Starring
Dana Carvey
Jennifer Esposito
Harold Gould
James Brolin
Brent Spiner
Narrated by
Harold Gould
Music by
Marc Ellis
Cinematography
Peter Lyons Collister
Edited by
Peck Prior
Sandy Solowitz
Production
company
Revolution Studios
Happy Madison Productions
Distributed by
Columbia Pictures
Release date
August 2, 2002
Running time
81 minutes
Country
United States
Language
English
Budget
$16 million[1]
Box office
$43.4 million[1]
The Master of Disguise is a 2002 American adventure comedy film directed by Perry Andelin Blake in his directorial debut and written by Dana Carvey and Harris Goldberg. Starring Carvey, Jennifer Esposito, Harold Gould, James Brolin, and Brent Spiner, the film was produced by Adam Sandler through his Happy Madison production company.
Although it was a modest box office success, the film was poorly received by most critics[2] and has appeared on several lists of the worst films ever made.
Contents
[hide]
1 Plot
2 Cast
2.1 Celebrities as themselves
3 Release
3.1 Box office
3.2 Critical reception
4 Accolades
5 Home media
6 Soundtrack
7 References
8 External links
Plot[edit]
In Italy, 1979, Fabbrizio Disguisey, the latest in a long line of secret agents known as "Masters of Disguise", breaks up a smuggling ring run by the evil Devlin Bowman. Not wanting his infant son Pistachio to receive the same dangerous future lifestyle as he and his lineage, Fabbrizio decides to keep his family's identity a secret.
Twenty-three years later, Fabbrizio runs an Italian restaurant in America with his wife and Pistachio. Bowman, fresh out of jail, kidnaps Fabbrizio and "Mama", forcing Fabbrizio to use his powers to steal legendary artifacts. After Fabbrizio's disappearance, Pistachio is visited by his grandfather, who reveals Pistachio's heritage and begins training him.
Pistachio gets the basics down and his grandfather gets him an assistant, Jennifer Baker, who is a little confused about what the job entails. The two find one of Bowman's cigars and learn of Bowman's scheme, as well as that he will be at an antiques fair the next day.
Pistachio and Jennifer go to the fair, with Pistachio disguised as an elderly woman. Bowman invites Jennifer to a party at his house. Pistachio goes to the party in disguise and distracts Bowman while Jennifer looks for clues.
That night, Pistachio and Jennifer look through the clues. Pistachio takes Jennifer home and they kiss. Bowman's men kidnap Jennifer. Pistachio breaks into Bowman's house to rescue Jennifer and his parents. Bowman has attached a mask of his own face to Fabbrizio's head. While the real Bowman escapes, Pistachio fights his father, who is brainwashed to think he is Bowman.
In the end, Pistachio helps his father snap out his trance, they free Mama, return the artifacts, Pistachio marries Jennifer and becomes an official Master of Disguise. However, Bowman still has the United States Constitution. The Disguiseys locate Bowman, defeat him and retrieve the Constitution.
Cast[edit]
Dana Carvey as Pistachio Disguisey
Dylan and Cole Sprouse as young Pistachio Disguisey
Dane Morris as teenage Pistachio Disguisey
Jennifer Esposito as Jennifer Baker
Harold Gould as Grandfather Disguisey
James Brolin as Fabbrizio Disguisey
Edie McClurg as "Mama" Disguisey
Brent Spiner as Devlin Bowman
Michael Bailey Smith as Bald Henchman
Austin Wolff as Barney Baker
Kenan Thompson as Kenan
Jay Johnston as Rex
Maria Canals as Sophia
Mark Devine as Trent
Virginia Hawkins as Interview Woman
Leland Crooke as Appraiser
Erick Avari as Cigar Maker
Vincent Castellanos as Art Dealer
Celebrities as themselves[edit]
Bo Derek
Michael Johnson
Jessica Simpson
Jesse Ventura
Release[edit]
The Master of Disguise was granted a PG certificate by the British Board of Film Classification for "mild language and mild comic violence". However, to obtain the PG certificate, seven seconds of material was cut, the reasons being "dangerous imitable technique, a series of head butts".[3]
Box office[edit]
The film opened in Australia on July 31, 2002, in the United Kingdom on January 17, 2003, and in the United States on August 2, 2002 in 2,565 theaters and earned $12,554,650 in the domestic box office, ranking third for the weekend, behind Signs and the second weekend of Austin Powers in Goldmember.[4] The film closed on November 28, 2002, having grossed $40,388,794 domestically and $3,022,207 internationally for a worldwide total of $43,411,001.[1] When the film was released in the United Kingdom, it originally opened at #14, before falling out of the charts completely by the next weekend.[5][6]
Critical reception[edit]
Rotten Tomatoes ranked the film as the 18th worst of the 2000s decade, with a 1% rating, its consensus reading: "An ill conceived attempt to utilize Dana Carvey's talent for mimicry, The Master of Disguise is an irritating, witless farce weighted down by sophomoric gags."[7] On Metacritic, the film has a 12 out of 100 score based on 24 amateur critics, indicating "overwhelming dislike".[8]
Roger Ebert reserved some praise for director Perry Andelin Blake's production design, as well as for Carvey's imitative abilities, but saw the overall movie as being "like a party guest who thinks he is funny and is wrong". Ebert also noted that the film is only 65 minutes long, but includes 15 minutes of credit cookies in order to qualify as a feature presentation.[9] Jamie Russell at the BBC described the film as being so bad as to make other movies such as Deuce Bigalow: Male Gigolo seem like comic equivalents to Citizen Kane: "Never have so many jokes clunked off the screen to such a silent audience. And never has 80 minutes seemed like such an eternity."[10]
Peter Bradshaw, writing for The Guardian, awarded the film two stars out of five, citing a limited number of amusing moments, but criticizing Carvey's characterizations and suggesting to the potential audience-member that "you might want to put the shotgun in your mouth".[11] Adam Smith in the Radio Times drew attention to the "lame screenplay" and "barely coherent plot", and noted that "when the nearest thing to a genuine joke is the bad guy's propensity for breaking wind whenever he laughs, you can be pretty sure that you're not in the hands of comedy geniuses".[12]
Alan Morrison, writing for Empire, suggested that it was no more than a feeble imitation of the character comedy of Mike Myers' Austin Powers series, concluding that The Master of Disguise was the single worst movie ever made: "a film about idiots, made by idiots, for idiots".[13]
Accolades[edit]
Bo Derek's cameo in the film earned her a Golden Raspberry Award nomination for Worst Supporting Actress, but lost to Madonna in Die Another Day.
Home media[edit]
The film was released on VHS and DVD on January 28, 2003.
Soundtrack[edit]
The Master of Disguise: Music from the Motion Picture
Soundtrack album by Various Artists
Released
July 23, 2002
Length
38:20[14]
Label
Music World, Sony
Professional ratings
Review scores
Source
Rating
Allmusic [15]
No.
Title
Writer(s)
Recording artist
Length
1. "M.A.S.T.E.R., Pt. 2" (featuring Lil' Fizz from B2K) Bryan Bonwell / Marques Houston / Jerome Jones / Kelton Kessee / Tony Oliver Play 2:56
2. "Fun" Rose Falcon / Billy Falcon Rose Falcon 2:55
3. "Happy Face" Beyoncé Knowles / Michael Cooper / Rob Fusari / Calvin Gaines / Falonte Moore Destiny's Child 4:18
4. "Eenie Meenie Minie Mo" Brian Adams / Teron Beal / Eddie Berkeley / Keir Gist Strong 3:22
5. "Walking on Sunshine" Kimberley Rew Val-C 2:49
6. "Master of Disguise" Colleen Fitzpatrick / Jimmy Harry Vitamin C 3:00
7. "Double Dutch Bus" Frankie Smith Devin Vasquez 3:21
8. "Conga" Enrique García Miami Sound Machine 4:14
9. "This Could Be Love" Troy Johnson / Solange Knowles Solange Knowles 4:02
10. "Cherry Pie" Marques Houston / Jerome Jones Jhené Aiko 3:07
11. "M.A.S.T.E.R., Pt. 1" (featuring Play) Bryan Bonwell Hardhedz / Hardhedzz 4:16
Total length: 38:20
References[edit]
^ Jump up to: a b c "The Master of Disguise (2002)". Box Office Mojo. Internet Movie Database. November 29, 2002. Retrieved July 22, 2015.
Jump up ^ "The Master of Disguise". Rotten Tomatoes. Retrieved December 28, 2017.
Jump up ^ "THE MASTER OF DISGUISE (PG)". British Board of Film Classification. October 10, 2002. Retrieved July 22, 2015.
Jump up ^ "Weekend Box Office Results for August 2-4, 2002". Box Office Mojo. Internet Movie Database. August 5, 2002. Retrieved July 22, 2015.
Jump up ^ "Weekend box office 17th January 2003 - 19th January 2003". 25thframe.co.uk. Retrieved 10 January 2017.
Jump up ^ "Weekend box office 24th January 2003 - 26th January 2003". 25thframe.co.uk. Retrieved 10 January 2017.
Jump up ^ "The Master of Disguise". Rotten Tomatoes. Retrieved December 28, 2017.
Jump up ^ "The Master of Disguise". Metacritic.com. Retrieved December 28, 2017.
Jump up ^ "The Master Of Disguise". Rogerebert.com. August 2, 2002.
Jump up ^ "The Master of Disguise (2003)". Bbc.co.uk. BBC. January 13, 2003.
Jump up ^ "The Master of Disguise". Guardian.co.uk. The Guardian. January 17, 2003.
Jump up ^ "The Master Of Disguise". Radiotimes.com. Radio Times. Retrieved May 9, 2017.
Jump up ^ "The Master Of Disguise Review". Empireonline.com. Empire. October 14, 2015.
Jump up ^ "The Master of Disguise (2002)". The Moviemusic Store. Retrieved August 31, 2014.
Jump up ^ "The Master of Disguise - Original Soundtrack - Songs, Reviews, Credits". AllMusic. Retrieved December 28, 2017.
External links[edit]
The Master of Disguise on IMDb
The Master of Disguise at Box Office Mojo
The Master of Disguise at Rotten Tomatoes
The Master of Disguise at Metacritic
[hide]
v
t
e
Happy Madison Productions
Films
Deuce Bigalow: Male Gigolo (1999)
Little Nicky (2000)
The Animal (2001)
Joe Dirt (2001)
Mr. Deeds (2002)
The Master of Disguise (2002)
Eight Crazy Nights (2002)
The Hot Chick (2002)
Anger Management (2003)
Dickie Roberts: Former Child Star (2003)
50 First Dates (2004)
The Longest Yard (2005)
Deuce Bigalow: European Gigolo (2005)
Grandma's Boy (2006)
The Benchwarmers (2006)
Click (2006)
I Now Pronounce You Chuck And Larry (2007)
Reign Over Me (2007)
Strange Wilderness (2008)
The House Bunny (2008)
You Don't Mess with the Zohan (2008)
Bedtime Stories (2008)
Paul Blart: Mall Cop (2009)
The Shortcut (2009)
Funny People (2009)
Grown Ups (2010)
Just Go with It (2011)
Zookeeper (2011)
Bucky Larson: Born to Be a Star (2011)
Jack and Jill (2011)
That's My Boy (2012)
Here Comes the Boom (2012)
Grown Ups 2 (2013)
Blended (2014)
Paul Blart: Mall Cop 2 (2015)
Pixels (2015)
Joe Dirt 2: Beautiful Loser (2015)
The Ridiculous 6 (2015)
The Do-Over (2016)
Sandy Wexler (2017)
The Week Of (2018)
Television series
Rules of Engagement (2007–13)
The Gong Show with Dave Attell (2008)
Nick Swardson's Pretend Time (2010–11)
Breaking In (2011–12)
The Goldbergs (2013–present)
Imaginary Mary (2017)
Key people
Adam Sandler
Allen Covert
Jack Giarraputo
Tim Herlihy
Heather Parry
Sony Pictures Entertainment
Retrieved from "https://en.wikipedia.org/w/index.php?title=The_Master_of_Disguise&oldid=841276686"
Categories:
English-language films
2002 films
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2000s comedy films
American comedy films
American satirical films
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Films set in 1979
Films set in 2002
Slapstick films
Happy Madison films
Revolution Studios films
Columbia Pictures films
Hidden categories:
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