Category: Intermediate Chemistry

Posted on

Mass spectrometry application: structure elucidation

Mass spectrometry is used to identify structures of organic compounds. The sample to be analysed is vaporised, fed into the ioniser and fragmented by bombarding electrons. With knowledge on the relative masses and abundances of isotopes making up the compound or by referencing against the spectra of known substances, the structure of the compound can be determined. For example,

Question

With reference to the spectrum below, deduce the structure of an organic compound that forms a sweet smelling liquid with acetyl chloride.

Answer

The sweet smelling liquid could be an ester and hence the organic compound could possibly be an alcohol. Assuming that it is a simple alcohol, its general formula is CnH2n+1OH.

Look for the peak representing the molecular ion (M+), i.e. the unfragmented sample molecule with an electron knocked out. This peak must be one with one of the highest u/z and with a relatively high abundance.

The highest peak is at 33 u/z but has a negligible abundance. Hence, the peak at 32 u/z is most probably the molecular ion. If so,

M_r\left ( C_nH_{2n+1}OH \right )=32

12n+2n+1+16+1=32

n=1

Hence, the organic compound could be methanol, CH3OH. We can verify our guess by analyzing the other peaks:

u/z

 Molecule Formula

Logic

33

 M++1 13CH3OH

13C has an abundance of only 1.07%, but much higher than that of 17O, 18O and 2H.

32

 M+ 12CH3OH

One of the highest u/z and high abundance

31

 M+-1 12CH3O+

Cleavage of polarised O-H bond; relatively stable and thus high abundance

30

 M+-2 12CH2O+

Loss of 1 proton attached to C and one to O

29

 M+-3 12CHO+

Loss of 2 protons attached to C and one to O

28

 M+-4 12CO+

Full deprotonation

15

 M+-7 12CH3+

Cleavage of polarised -OH bond

With the above deductions, we can conclude that the spectrum is one for methanol.

 

 

next article: Other forms of mass spectrometry
Previous article: Applcation: relative atomic mass
Content page of mass spectrometry
Content page of intermediate chemistry
Main content page
Posted on

Other forms of mass spectrometry

Mass spectrometers can be combined with other instruments to form powerful analytical tools. One such example is the Gas Chromatography-Mass Spectrometer (GC-MS), which combines a gas chromatograph with a mass spectrometer.

The concept of gas chromatography is similar to that of paper chromatography with the exception of the sample and the mobile phase being gases instead of liquids. The sample travels through gas chromatograph part of GC-MS first, where the different types of molecules in the sample are retained by the stationary phase for different durations before they flow into the mass spectrometer portion where they are fragmented and analysed. The data is plotted on a three dimensional graph with retention time, mass-to-charge ratio and relative abundance as the axes. As such, the GC-MS spectrum contains more information regarding the sample as compared to that of TIMS. GC-MS is commonly used by border security forces to detect narcotics and explosives in luggage or on human beings and by environmental agencies to monitor pollutants in the air.

Another important instrument is the Inductively Coupled Plasma Mass Spectrometer (ICP-MS), which is capable of detecting trace amounts of metals and non-metals. The main difference between an ICP-MS and a TIMS is the use of an inductively coupled plasma (a concentration of argon ions and electrons that is produced by inductively heating argon with an electromagnetic coil) to ionise the sample. ICP-MS has greater speed, precision, and sensitivity over TIMS and is used in the fields of medicine, toxicology, forensics, radiometric dating, pharmacy, etc.

 

Previous article: Application: structure elucidation
Content page of mass spectrometry
Content page of intermediate chemistry
Main content page
Posted on

Mass spectrometry application: relative atomic mass

Mass spectrometry plays a crucial role in determining relative atomic mass by accurately measuring the masses of isotopes in a sample, providing essential data for understanding elemental abundance and chemical properties.

As mention in an earlier article, the relative masses of isotopes can be determined by comparing the charge-to-mass ratio of an isotope of interest with that of carbon-12. For example, mass spectrometric data for the ratio of the mass-to-charge ratio (u/z) of 2H to that of 12C is 0.167842. Thus, the relative mass of 2H on the carbon-12 unified atomic mass unit scale is:

u\; of \; ^{2}H=\frac{u/z\; of\;\; ^2H}{u/z\; of\; \; ^{12}C}\; \times \; u\; of\; ^{12}C

u\; of \; ^{2}H=0.167842\times 12\, u

u\; of \; ^{2}H=2.014104\, u

Consequently, the relative atomic mass of an element can be calculated from its isotopic abundance spectrum.

With reference to the spectrum above, the relative atomic mass of chlorine is:

A_{r,Cl}=\frac{100\times 34.969}{100+31.96}+\frac{31.96\times 36.966}{100+31.96}=35.45

Question

With reference to the table below, deduce the spectrum for the chlorine gas ion, Cl2+.

Isotope

 Relative isotopic mass

Relative abundance, %

35Cl

 35

75.78

37Cl

 37

24.22
Answer

Assuming a random distribution of isotopes in the sample of chlorine gas used and that no fragmentation occurred, the possible permutations of chlorine isotopes in a chlorine gas ion are:

Permutation, XClYCl+

 Relative isotopic mass

No. of permutations

35Cl35Cl

 70

1

35Cl37Cl or 37Cl35Cl

 72

2

37Cl37Cl

 74

1

The probability of observing each permutation is:

Permutation, XClYCl+

Probability = abundance X x abundance Y x no. of permutations

35Cl35Cl

0.7578 x 0.7578 x 1 = 0.5742

35Cl37Cl or 37Cl35Cl

0.7578 x 0.2422 x 2 = 0.3671

37Cl37Cl

0.2422 x 0.2422 x 1 = 0.0587

Hence, the spectrum has the following characteristics:

Permutation

 Relative isotopic mass Relative abundance (%)

% of predominant peak

35Cl35Cl

 70 57.42

100.00

35Cl37Cl or 37Cl35Cl

 72 36.71

63.93

37Cl37Cl

 74 5.87

10.22

The values for ‘% of predominant peak’ are calculated relative to the abundance of 35Cl35Cl.

 

Here’s another way to explain the permutation of isotopes in chlorine gas:

Chlorine gas is formed by the collision of two chlorine atoms. Consider an infinitesimal two-dimensional space in a vessel that is filled with chlorine atoms, with the space large enough to accommodate two chlorine atoms (see diagram below).

If we further consider a single axis of collision through that space, and that chlorine atoms are approaching the space along that axis from the left and from the right, we end up with four possible products: 35Cl35Cl, 35Cl37Cl, 37Cl35Cl and 37Cl37Cl. If the abundances of 35Cl and 37Cl are equal, we have twice as many chlorine molecules with mixed isotopes than 35Cl35Cl or 37Cl37Cl, i.e. in the ratio of 2:1:1. If the abundances of 35Cl and 37Cl are unequal, we multiply the numbers in the ratio with the respective abundances of isotopes to get the probabilities of occurrence of the three molecular species. The same result is obtained if we rotate the axis of collision in any direction in three dimensions, or if the collision happens at an angle.

 

next article: application: structure elucidation
Previous article: Modern mass spectrometer design
Content page of mass spectrometry
Content page of intermediate chemistry
Main content page
Posted on

Modern mass spectrometer design

The modern mass spectrometer, an advanced analytical instrument, precisely measures the mass-to-charge ratios of ions, enabling detailed analysis of molecular composition and structure across a wide range of scientific fields.

As mentioned in an earlier article, J. J. Thomson constructed one of the earliest mass spectrometers and conducted the well-known experiment to determine the mass-to-charge ratio of an electron. Since then, different mass spectrometer designs have been developed. For instance, a typical modern Thermal Ionisation Mass Spectrometer (TIMS) consists of the following:

    • Ioniser, where the sample to be analysed is bombarded by electrons to form ions and ion fragments, which are then accelerated into the mass analyser.
    • Mass analyser that utilises a magnetic field to deflect and separate the ions according to their mass-to-charge ratios (u/z). Note that the mass-to-charge ratios before the 1980s were quoted in u/e where e is the charge of an electron instead the number of charges, which is z.
    • Detector that measures the abundance of ions with reference to their mass-to-charge ratios and converts the data into electrical signals.

The result is a plot of ion abundance versus mass-to-charge ratio, with the height of the peaks normalised to the most abundant ion in the spectrum.

 

next article: application: relative atomic mass
Previous article: application: relative isotopic mass
Content page of mass spectrometry
Content page of intermediate chemistry
Main content page
Posted on

How to select an appropriate pH indicator for titration?

How do we select an appropriate pH indicator for titration?

Let’s consider the same indicator as in the previous article: bromocresol green. If the pH of the analyte containing bromocresol green is between 0 and 3.8, it appears yellow; if the pH is between 5.4 and 14, it appears blue. Between 3.8 and 5.4, the colour changes from yellow to blue as the concentration ratio of [In–]/[HIn] changes from 10 to 0.1. In fact, when [In–] = [HIn], the colour of the titrand is green, indicating equal amounts of the yellow acid and the blue conjugate base.

Therefore, bromocresol green is a suitable indicator for a titration with a stoichiometric point (SP) that involves a sharp change of pH from less than 3.8 to more than 5.4 when about 0.1 cm3 of solution (about two drops) is added to the analyte. For example, the titration of 10 cm3 of 0.200 M of HCl with 0.100 M of NaOH sees a colour change of yellow to blue at the SP (see table and graph below).

One drop before SP

SP

One drop after SP

Volume, cm3

19.95

20.00

20.05

pH

3.78

7.00

10.22

In other words, a suitable pH indicator for titration is one whose pH range is narrower than and within the range of the change in pH of the analyte at the stoichiometric point (ideally with one or two drops of solution added).

Another example is the titration of 10 cm3 of 0.200 M of aqueous NH3 (Ka = 1.8 x 10-5) with 0.100 M of HCl, which sees a colour change from blue to yellow at the SP (see table and graph below).

One drop before SP

SP

One drop after SP

Volume, cm3

19.95

20.00

20.05

pH

6.65

5.22

3.78

The following table lists some common indicators and their pH ranges:

Indicator

 Low pH colour pH range High pH colour

pKa

Thymol blue (first transition)

 red 1.2 – 2.8 yellow

1.65

Thymol blue (second transition)

 yellow 8.0 – 9.6 blue

8.96

Methyl yellow

 red 2.9 – 4.0 yellow

3.30

Bromophenol blue

 yellow 3.0 – 4.6 blue 3.85
Congo red blue-violet 3.0 – 5.0 red

4.10

Methyl orange

 red 3.1 – 4.4 yellow 3.46

Screened methyl orange (second transition)

 purple-grey 3.2 – 4.2 green 3.70

Bromocresol green

 yellow 3.8 – 5.4 blue 4.66

Methyl red

 red 4.4 – 6.2 yellow

5.00

Bromothymol blue (second transition)

 yellow 6.0 – 7.6 blue

7.10
Phenol red yellow 6.4 – 8.0 red

7.81

Phenolphthalein (second transition)

 colorless 8.3 – 10.0 purple-pink 9.30
Thymolphthalein (second transition) colorless 9.3 – 10.5 blue

9.90

Alizarine Yellow R

 yellow 10.2 – 12.0 red 11.2
Indigo carmine blue 11.4 – 13.0 yellow

12.2

 

Next article: structures of common indicators
PREVIOUS ARTICLE: THE PH RANGE OF AN INDICATOR
Content page of ph indicators
Content page of Intermediate chemistry
Main content page
Posted on

Structures of common indicators

The diverse chemical structures of common pH indicators, ranging from simple organic dyes to complex conjugated systems, play a crucial role in their ability to visually signal changes in acidity and alkalinity.

Methyl orange is a commonly used indicator. At low pH, protonation of the nitrogen at the para position of the benzenesulphonate moiety is preferred over protonation of the negatively charged oxygen in the sulphonate group, as this confers resonance stabilization (benzenoid-quinonoid tautomerism) to the acid form of methyl orange.

Litmus, extracted from lichen, is processed and applied onto filter paper. It has a pH range of 4.5 to 8.3 and contains the chromophore, 7-hydroxyphenoxazone, which occurs in the following states:

Phenolphthalein has three pKa values with four different forms. However, stoichiometric points of titrations are most commonly monitored with the middle pKa value of 9.30, which involves the lactone and phenolate (dianionic) forms:

Thymol blue has two transition range with pKa1 = 1.65 and pKa2 = 8.96.

 

Previous article: How to select an appropriate ph indicator for titration?
Content page of ph indicators
Content page of Intermediate chemistry
Main content page
Posted on

The pH range of an indicator

The pH range of an indicator is the pH interval in which the indicator changes colour.

To elaborate further, we will use the Henderson-Hasselbalch equation to study the equilibrium of a pH indicator, which is a weak acid.

pH=pK_{In}+log\frac{In^-}{HIn}

Consider a titration using two drops of bromocresol green with pKin = 4.7. We have

pH=4.7+log\frac{In^-}{HIn}\; \; \; \; \; \; \; \; (1)

Since the concentration of the indicator in the analyte is very low, the contribution of H+ from the indicator is negligible and does not affect the total concentration of H+ in the solution. Hence, the pH value in eq1 is solely due to the H+ of the analyte; that is, the pH of the analyte determines the position of the equilibrium of the indicator.

When bromocresol green is predominantly in the form HIn, it appears yellow. If it is predominantly in the form In–, we see it as blue.

\begin{matrix} HIn(aq)\\yellow \end{matrix}\rightleftharpoons H^+(aq)+\begin{matrix} In^-(aq)\\blue \end{matrix}

As a rule of thumb, our eyes can perceive a complete change in the colour of the indicator from its acid form to the conjugate base and vice versa, when the concentration of one form is ten times greater than the other. So, bromocresol green appears blue when:

[In^-]\geqslant 10[HIn]\; \; \; \; \; \; \; \; (2)

which corresponds to the indicator in a pH environment of pH ≥ 5.7 (by substituting eq2 in eq1).

At the other limit, bromocresol green appears yellow when:

[HIn]\leqslant 10[In^-]\; \; \; \; \; \; \; \; (3)

which corresponds to the indicator in a pH environment of pH ≤ 3.7 (by substituting eq3 in eq1).

In other words, if a few drops of bromocresol green are added to an analyte with a pH ≥ 5.7 and another with a pH ≤ 3.7, the solutions will appear blue and yellow, respectively.

We call this difference in pH (3.7 to 5.7 for bromocresol green) the pH range of the indicator.

NEXT ARTICLE: HOW TO SELECT an APPROPRIATE PH INDICATOR FOR TITRATION
Previous article: Overview
Content page of ph indicators
Content page of Intermediate chemistry
Main content page
Posted on

Strong base to weak acid titration curve: overview

We have explained the shape of a strong-base-to-strong-acid titration curve in a basic level article. We shall now describe the shape of a strong-base-to-weak-acid titration curve using simple equations like the Henderson-Hasselbalch equation. To do so, we need to analyse the curve at different stages of the titration. These stages are:

    • Start point
    • After start point but before stoichiometric point
    • Maximum buffer capacity point
    • Stoichiometric point
    • Beyond stoichiometric point

Proceed to the next few articles to understand the mathematics and assumptions behind the formulae for the various stages of the pH curve.

 

next article: start point of titration
Content page of strong base to weak acid titration curve
Content page of intermediate chemistry
Main content page
Posted on

Stoichiometric point of titration

The stoichiometric point of titration is the moment when the moles of titrant added precisely neutralise the moles of analyte in the solution, indicating a complete reaction between the two species.

Generally, the salt of a weak acid and a strong base (e.g. CH3COOH and NaOH) is a weak conjugate base, which hydrolyses in water, i.e. reacts with water to reform the acid according to the following reaction:

CH_3COO^-(aq)+H_2O(l)\; \begin{matrix}K_h\\ \rightleftharpoons \end{matrix}\; CH_3COOH(aq)+OH^-(aq)

where Kh is the hydrolysis constant. This results in a basic solution since OH– is formed. During the stage of a strong base to weak acid titration between the start point and the stoichiometric point, the presence of unneutralised acid in the reaction flask causes the equilibrium of the hydrolysis reaction to shift to the left. However, when the stoichiometric point is reached, the weak acid is completely neutralised by the base and we can no longer ignore the effects of water on the pH of the solution. Since,

K_a=\frac{[CH_3COO^-][H^+]}{[CH_3COOH]}\; \; \; and\; \; \; K_w=[H^+][OH^-]

K_h=\frac{[CH_3COOH][OH^-]}{[CH_3COO^-]}=\frac{K_w}{K_a}

From the hydrolysis equation, [CH3COOH] = [OH–], so

\frac{K_w}{K_a}=\frac{[CH_3COOH][OH^-]}{[CH_3COO^-]}=\frac{[OH^-]^2}{[CH_3COO^-]}

[OH^-]=\sqrt{\frac{K_w}{K_a}[CH_3COO^-]}

\frac{K_w}{[H^+]}=\sqrt{\frac{K_w}{K_a}[CH_3COO^-]}

Taking the logarithm on both sides of the above equation and applying the definitions of pH, pKa and pKw,

pH=\frac{pK_w+pK_a+log[A^-]}{2}\; \; \; \; \; \; \; \; (4)

where [A–] = [CH3COO–], the concentration of the salt at the stoichiometric point before hydrolysis. Eq4 is the general equation to calculate the pH of a strong base to weak acid titration at the stoichiometric point. 

Question

Calculate the pH of the titration of 10 cm3 of 0.200 M of CH3COOH (Ka = 1.75 x 10-5) with 0.100 M of NaOH at the stoichiometric point.

Answer

Using eq4,

pH=\frac{-log10^{-14}-log\left ( 1.75\times 10^{-5}\right)+log\left ( \frac{0.100\times 0.02}{0.01+0.02} \right )}{2}=8.79

Note that [A–] can be computed using either the acid or the base.

 

The same logic applies when determining the equation for the pH of a strong acid to weak base titration at the stoichiometric point, with the expected pH lower than 7.

 

next article: Beyond stoichiometric point of titration
Previous article: After start point and before stoichiometric point of titration
Content page of strong base to weak acid titration curve
Content page of intermediate chemistry
Main content page
Posted on

Beyond stoichiometric point of titration

Beyond the stoichiometric point of titration, the pH of the solution experiences significant shifts, often leading to rapid changes that highlight the excess of titrant and the resultant dominance of the strong base or acid in the solution.

To characterised this region, we make the following assumption:

    • [OH–] from water is negligible, i.e. pH of the solution is solely determined by [OH–]ex from the excess base added after the stoichiometric point as the dissociation of water is again suppressed at this stage.

Even though the auto-dissociation of water is suppressed, water is still equilibrating between its molecular form and its ionic components. Applying the above assumption, the equilibrium constant of water is:

K_w=[H^+][OH^-]\approx [H^+][OH^-]_{ex}

Taking the logarithm on both sides of the above equation and applying the definitions of pH and pKw,

pH=pK_w+log[OH^-]_{ex}\; \; \; \; \; \; \; \; (5)

Eq5 is the general equation to calculate the pH of a strong base to weak acid titration beyond its stoichiometric point. The same logic applies when determining the equation for the pH of a strong acid to weak base titration beyond the stoichiometric point.

The diagram below shows the superimposition of eq5 (purple curve), over the complete pH titration curve (blue curve) that disregard the above assumption, for the titration of 10 cm3 of 0.200 M of CH3COOH (Ka = 1.75 x 10-5) with 0.100 M of NaOH.

Even though the two curves appear to fit perfectly, they do not actually coalesce and are still two separate curves (discernible when the axes of the plot are scaled to a very high resolution). However, for practical purposes, eq5 is a very good approximation of a pH curve for the region beyond the stoichiometric point of a strong acid to weak base titration.

 

next article: Combining all equations
Previous article: stoichiometric point of titration
Content page of strong base to weak acid titration curve
Content page of intermediate chemistry
Main content page