Inhibition Effects of Glucose on Clathrate Hydrate (H-Lw...

Inhibition Effects of Glucose on Clathrate

Hydrate (H-Lw-V) Equilibrium

by

Anthony Stephen Carbone

Department of Chemical Engineering

McGill University

Montreal, Quebec, Canada

May, 2011

A thesis submitted to McGill University in partial fulfillment of the

requirements of the degree of Masters of Engineering

© Anthony Stephen Carbone, 2011

Inhibition Effects of Glucose on Clathrate Hydrate (H-Lw-V) Equilibrium McGill University

ii Master’s Thesis © Anthony Stephen Carbone, 2011

Abstract

The three phase (H-Lw-V) equilibrium conditions for methane clathrate hydrates and

carbon dioxide clathrate hydrates in aqueous solution containing glucose were experimentally

determined. The glucose aqueous solutions varied between 10-30 wt% glucose, with the pressure

ranging from 3.4 to 7.8 MPa for the methane systems and 1.6 to 3.8 MPa for the carbon dioxide

systems. The experimental temperature for both hydrate systems in question ranged from 275.15

to 281.25 K.

Experimental results showed that systems containing glucose did in fact exhibit an

inhibiting effect on hydrate equilibrium conditions. This is shown by the equilibrium plots where

equilibrium curves of the glucose containing systems reside above (at higher pressures for equal

temperatures) the equilibrium curve of pure water. It was also found that as the glucose

concentration increased in the aqueous solution present, so did the degree of inhibition. All

systems containing glucose were also found to follow the same trend as the systems containing

pure water in terms of the apparent curvature of the equilibrium line. When compared to other

common thermodynamic hydrate inhibitors, only the 30 wt% glucose solution was able to exhibit

a level of hydrate inhibition at par with the others.


iii Master’s Thesis © Anthony Stephen Carbone, 2011

Résumé

Les conditions d’équilibre pour les trois phases H-Lw-V ont été déterminées

expérimentalement pour des systèmes d’hydrates de méthane et d’hydrates de dioxyde de

carbone en présence d’une solution aqueuse de glucose. La concentration de la solution aqueuse

de glucose était variée entre 10-30 pourcentages massiques sous une gamme de température de

275.15 à 281.25 K. La gamme de pression résultante du système de méthane est 3.4 à 7.8 MPa,

et 1.6 à 3.8 MPa pour le système de dioxyde de carbone.

Les résultats des expériences avec les systèmes contenant le glucose ont démontré que le

glucose a un effet bloquant à propos de l’équilibre de la phase d’hydrate. Cela est évident après

une analyse des graphiques des données d’équilibre des systèmes contenant le glucose. La

courbe d’équilibre pour tous les systèmes contenant le glucose résident au dessus de la courbe

d’équilibre d’eau pure. Cela indique que pour la même température, les systèmes contenant le

glucose requièrent une pression plus élevée pour atteindre l’équilibre avec la phase d’hydrate. En

plus, les systèmes contenant le glucose suivent la même tendance que les systèmes contenant de

l’eau pure. En comparant les systèmes de glucose en question avec d’autres systèmes d’hydrates

contenant des bloquantes thermodynamiques populaires, il est possible de constater que

seulement la solution de glucose à 30 pour cent massique démontre un niveau bloquant

comparable.


iv Master’s Thesis © Anthony Stephen Carbone, 2011

Acknowledgements

I would like to take this opportunity to thank all of the individuals who have helped me

throughout my studies and research, ultimately leading to the successful completion of my

masters degree in engineering. Without their help and support, my time as a graduate student

would not have been as fulfilling and enjoyable as it has been.

Firstly, I would like to sincerely thank my supervisor, Dr. Phillip Servio, for giving me

the opportunity to work under someone as talented as he is. I am extremely grateful for all of his

support, trust, and patience, as well as for his ability to motivate his students and elevate them to

be better researchers, and more importantly, better people. This dissertation, as well as the

journey it represents, would not have been the same without him.

Secondly, I would like to thank Dr. Servio’s hydrate research team, in particular Hallvard

Bruusgaard and Dany Posteraro. Mr. Bruusgaard took the time and had the patience to teach me

all I needed to know about the experimental set-up used throughout my masters degree, while

Mr. Posteraro was always there to lend a helping hand. Both of their friendships will not be

forgotten.

Thirdly, I would like to thank the technical staff as well as the administrative staff in the

chemical engineering department at McGill University who were always fast to act on anything

needing their attention.

Finally, I cannot end without giving a great big thank you to all of my family and friends

who have loved and supported me every step of the way. You are my inspiration to keep working

as hard as I do, and for that, I cannot thank you enough.


1 Master’s Thesis © Anthony Stephen Carbone, 2011

Table of Contents

1.0 Introduction 4

1.1 Gas Hydrates 4

1.2 History of Gas Hydrates 4

1.3 Naturally Occurring Gas Hydrates 5

1.4 Hydrate Structure 6

1.5 Phase Equilibrium 8

1.6 Methane Hydrates 10

1.7 Carbon Dioxide Hydrates 11

1.8 Nucleation and Growth 11

1.9 Hydrates Inhibition 12

1.9.1 Antifreeze Proteins, Sugars, and Fatty Acids 13

1.10 Research Objectives 15

2.0 Materials and Methods 16

2.1 Experimental Set-Up 16

2.2 Materials 18

2.3 Experimental Procedure 18

2.4 Experimental Conditions 19

3.0 Results and Discussion 20

3.1 Methane Experiments 20

3.2 Carbon Dioxide Experiments 25

4.0 Conclusions and Recommendations 31

4.1 Conclusions 31

4.2 Recommendations 32

5.0 References 33

Publications 38



List of Figures

Figure 1: Structure I, structure II, and structure H cavities and unit cells (hydrate structures) 7

Figure 2: Sample three-phase equilibrium curve of a methane + butane system 9

Figure 3: Partial phase diagram for the methane-water system as published by Deaton

and Frost 10

Figure 4: Partial phase diagram for the carbon dioxide-water system as published by Deaton and

Frost 11

Figure 5: Proposed disaccharide core structure of the thermal hysteresis factor isolated from Upis

ceramboides 14

Figure 6: Experimental set-up 17

Figure 7: Measured equilibrium conditions of the water + glucose + methane system

(0 wt% glucose) 21

Figure 8: Measured equilibrium conditions of the water + glucose + methane system (various

glucose concentrations) 23

Figure 9: Measured and reported values of various methane hydrate systems in the presence

of a common thermodynamic hydrate inhibitor 24

Figure 10: Measured equilibrium conditions of the water + glucose + carbon dioxide

system (0wt% glucose) 26

Figure 11: Measured equilibrium conditions of the water + glucose + carbon dioxide

system (various glucose concentrations) 29

Figure 12: Measured and reported values of various carbon dioxide hydrate systems

in the presence of a common thermodynamic hydrate inhibitor 30



List of Tables

Table 1: Hydrate-liquid-vapour equilibrium data for the system CH4-0 wt% glucose/H20 20




Table 5: Hydrate-liquid-vapour equilibrium data for the system CO2-0 wt% glucose/H20 25






1.0 Introduction

1.1 Gas Hydrates

Gas hydrates are non-stoichiometric crystalline solids formed by water molecules in the

presence of gaseous or volatile liquid molecules at suitable temperatures and pressures. Under

appropriate hydrate forming conditions, water molecules become linked together through

hydrogen bonding, producing a cage-like structure stabilized by the presence of hydrate forming

gas molecules. Stabilization occurs through physical bonding of the present gas via weak van der

Waals forces. The formation of gas hydrates is thermodynamically unstable without the presence

of the guest gas molecules.

1.2 History of Gas Hydrates

Sir Humphry Davy was the first to discover gas hydrates. His discovery occurred in 1810,

with his results confirmed in 1823 by Michael Faraday [1]. Research in the field of gas hydrates

initially began as a purely academic pursuit, identifying which molecules possessed the

capability of forming hydrates, and under which temperature and pressure conditions these

hydrates existed. Since the 1930s, hydrates have gained industrial interest in the oil and gas

industry when it was discovered by Hammerschmidt that gas hydrate formation occurs in oil

pipelines, causing partial or total blockage of flow [2]. Due to these implications, a significant

portion of hydrate research is devoted to hydrate inhibition through the use of chemical, thermal,

mechanical and hydraulic inhibitors [3].

The oil and gas industry is continuously searching for hydrate inhibitors that will allow

oil and gas extraction in deeper waters without the threat of hydrate formation impeding



everyday operations [14]. As Kelland reported, low dosage hydrate inhibitors (LDHIs), which

are subdivided into anti-agglomerates (AAs) and kinetic hydrate inhibitors (KHIs), are a more

recent alternative technology to thermodynamic inhibitors which prevent gas hydrates from

plugging oil and gas production wells and pipelines [15]. AAs and KHIs are generally water

soluble polymers [14]. Other potential hydrate inhibitors include anti-freeze proteins (AFPs) and

anti-freeze glycoproteins (AFGPs), which are proteins that bind to the surface of ice nuclei, thus

acting as an inhibitor to ice formation. Recent work has shown that AFPs also inhibit hydrate

formation [14].

1.3 Naturally Occurring Gas Hydrates

Gas hydrates occur naturally around the world and can be found below the permafrost

zone and in subsea sediments in the Arctic, the Antarctic, and tropical and subtropical oceans [4].

These hydrates, composed mostly of methane, exist in very large quantities, making them a

potential alternative source of energy [5]. Estimates predict that global hydrate reserves contain

more than twice as much organic carbon than current global fossil fuel reserves [6]. Hydrates

formed with carbon dioxide are also currently being researched as a means of carbon dioxide

sequestration into the ocean, thus reducing the amount of carbon dioxide released into the

atmosphere, mitigating the effect of global warming [7].

One of the most recent technologies proposed with regards to gas hydrates is their

implementation in storage and transportation of gases. As suggested by Thomas and Dawe, who

have reviewed ways of transporting natural gas, gas hydrate gas storage and transportation is a

promising technology. As compared to Liquefied Natural Gas (LNG) or Compressed Natural

Gas (CNG), gas hydrates do not require cryogenic temperatures (-160 oC) or extremely high



pressures (200 bar) as they can be produced at normal temperatures (3 to 5 oC) and pressures (10

to 40 bar) [8]. Gudmundsson et al. have shown that hydrate based transportation technology is

appropriate for low to medium volumes (0.1 to 1 billion cubic meters) and short to medium

distances (100 to 5,000 km) [9].

The need for more fundamental research in the field of gas hydrates has been highlighted

by the growing interest in the field. The prospect of deep sea oil and gas extraction requires a

durable and inexpensive hydrate inhibitor to facilitate the extraction of these valuable resources.

1.4 Hydrate Structure

Gas hydrates take the form of different structures depending on the size of the guest

molecule present. To date, four hydrate structures have been observed: structure I (sI), structure

II (sII), structure H (sH), and tetragonal structure TS-I (only formed with Br2 guest molecule)

[10]. The figure below shows the three most common structures, along with their relative size

and potential guest molecules.



Figure 1: Structure I, structure II, and structure H cavities and unit cells (hydrate structures) (Image by Sloan

[11]).

Structure I hydrates are composed of 2 small cavities and 6 large cavities, leading to a

unit cell containing a total of 46 water molecules. The small and large cavities have an average

radius of 3.95 Å and 4.33 Å respectively [11]. Structure I hydrates form in the presence of guest

molecules smaller than propane, such as methane, ethane, and carbon dioxide, explaining why

this is the predominant structure found to occur naturally below the ocean floor. A body centered

cubic unit cell is typical of structure I hydrates [11].

Structure II hydrates contain 16 small cavities and 8 large cavities per unit cell, with a

total of 136 water molecules present [11]. The small and large cavities have an average radius of

3.91 Å and 4.73 Å respectively [11]. Structure II hydrates form in the presence of molecules



larger than those found in the formation of structure I hydrates, such as propane and iso-butane,

but smaller than pentane.

The structure H hydrate was discovered more recently by Ripmeester et al. [12]. Unlike

hydrates of structure sI and sII, structure H hydrates contain three different types of cavities, 3

small, 2 medium, and 1 large [11]. These cavities have average radii of 3.91 Å, 4.06 Å, and 5.71

Å respectively and due to their larger size, require larger guest molecules to ensure

thermodynamic stability. Hydrates of structure H contain 34 water molecules per unit cell [11].

1.5 Phase Equilibrium

Gas hydrates are often grown in stirred tank reactors or PVT cells (a cell where pressure,

temperature and volume are variable). These reactors are pressurized with a hydrate forming gas

or gas mixture which is consumed by the liquid water, leading to the formation of suspended

hydrate crystals.



Figure 2: Sample three-phase equilibrium curve of a methane + butane system [13].

The line presented in Figure 2 represents the three-phase equilibrium line between liquid

water, gas, and hydrate. Above the equilibrium line, in the formation region, hydrate and liquid

water is present, while below the equilibrium line, in the decomposition region, hydrate growth

or formation will not occur.

Decomposition region

Formation region



1.6 Methane Hydrates

Deaton and Frost [36] were the first to publish equilibrium data about a methane hydrate

system in 1946. Their experimental method for determining such equilibrium data involved the

use of the isothermal pressure-search method, a method still employed today to determine

equilibrium data of various gas hydrate systems. Below is a plot of the data published by Deaton

and Frost for a methane hydrate system:

Figure 3: Partial phase diagram for the methane-water system as published by Deaton and Frost [1] (1946).

The curve separates the system conditions where methane hydrates can exits. Above the

curve, methane hydrates will nucleate and grow, whereas below the curve, hydrates cannot

nucleate and will decompose if present. On the curve itself, the methane hydrates are in

equilibrium with the system (allowing for the co-existence of methane gas, methane hydrates,

and water). At equilibrium, the hydrate growth rate is equal to the rate of hydrate decomposition,

allowing for no net apparent hydrate growth within the system.

0

0.5

1

1.5

2

2.5

3

273 274 275 276 277 278 279

p / M

Pa

T / K

Methane Hydrateand Liquid Water

Methane Gasand Liquid Water



1.7 Carbon Dioxide Hydrates

In 1946, Deaton and Frost also published equilibrium results for the carbon dioxide (CO2)

hydrate system [36]. The same experimental method used to discover the methane hydrate

equilibrium data was also used for the CO2 hydrate system, with their experimental data plotted

below:

Figure 4: Partial phase diagram for the carbon dioxide-water system as published by Deaton and Frost [1]

(1946).

Once again, the pressure is plotted on the y-axis in MPa and the temperature on the x-axis

in K.

1.8 Nucleation and Growth

Hydrates are generated in a supersaturated aqueous solution of hydrate forming gas

dissolved in a bulk liquid phase (water). Like other crystal growth patterns, gas hydrate nuclei

are constantly forming and decomposing during the nucleation phase until a critical size or stable

hydrate nuclei is formed [21]. Critical radii for stable hydrate nucleation have been reported to be

0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

5

272 274 276 278 280 282 284

p / M

Pa

T / K

CO2 Hydrateand Liquid Water

CO2 Gasand Liquid Water



between 15 nm to 50 µm [22]. The nucleation stage of hydrate growth is widely accepted as

being stochastic in nature, with factors such as driving force masking the unpredictability of the

system’s behaviour.

Vysniauskas and Bishnoi first assumed an Arrhenius-type dependence for the

temperature, pressure, water-gas interfacial area, and degree of supercooling on the reaction rate

constant [26]. The first to acknowledge hydrate growth as being crystalline in nature was

Englezos et. al. [23]. This model is based on crystallization and two-film theory. It describes the

growth of hydrate crystals in constantly stirred systems where agglomeration is assumed to be

negligible, using the reaction rate constant as the only adjustable parameter [23]. Skovborg and

Ramussen later proposed their own model based on the assumption that the transport of gas in

the gas phase to the liquid water phase is the rate determining step in the overall hydrate

formation process [24]. While there is some controversy regarding the rate determining step for

hydrate growth in a semi-batch stirred tank crystallizer, the most recent model of Bergeron and

Servio circumvents this ambiguity [25]. The Bergeron and Servio model relies on experimental

measurements of mole fractions of the hydrate forming gas in the bulk liquid phase to describe

hydrate growth, regardless of the rate determining step [25].

1.9 Hydrate Inhibition

The inhibition of hydrates in the oil and gas industry has traditionally been accomplished

by adding thermodynamic inhibitors such as methanol and glycol [17]. Thermodynamic

inhibitors are additives that change the hydrate’s thermodynamic forming conditions, thus

preventing hydrate growth under normal hydrate forming conditions [14]. One of the major

drawbacks of thermodynamic inhibitors is that they must be used in relatively large



concentrations (10-50 wt%) [14]. To circumvent this problem, low dosage hydrate inhibitors

(LDHIs) such as anti-agglomerates (AAs) and kinetic hydrate inhibitors (KHIs) have been

developed [15]. LDHIs are commonly water soluble polymers that are effective at concentrations

below 1 wt% [14]. AAs allow hydrates to form yet prevent them from accumulating into large,

uncontrollable masses while KHIs prevent hydrate formation by delaying their nucleation and

growth [14]. Nevertheless, the use of AAs and KHIs has been limited due to environmental

restrictions as a result of their lack of biodegradability [18].

1.9.1 Antifreeze Proteins, Sugars, and Fatty Acids

Antifreeze proteins (AFPs) and antifreeze glycoproteins (AFGPs) have recently been

under study as hydrate inhibitors due to their environmental compatibility. AFPs lower the

freezing point of ice, but do not affect its melting point; the difference between these two

temperatures is called thermal hysteresis [19]. The growth of the ice crystal is thus easily

controlled by keeping the system within the thermal hysteresis range. Zeng et. al. has concluded

that AFPs act on hydrate crystals via an adsorption-inhibition mechanism, as has been suggested

for the interaction of AFPs and ice, thus inhibiting hydrate formation and growth [20]. However,

protein denaturing is a major obstacle due to their natural limited shelf-life under ideal

conditions, and due to their quick unfolding outside of their narrow ideal operating condition

range. Therefore, finding a function, durable, cost-effective, and environmentally-friendly

hydrate inhibitor is currently a major priority for the oil and gas industry.

Chun et. al. have shown that simple sugars, like glucose and sucrose, have a substantial

inhibiting effect on a hydrate system [31]. Experiments were conducted on a CHClF2 (R22)

hydrate system in the presence of sucrose, glucose and lactic acid at concentrations of 20 and 40



wt%. The equilibrium hydrate dissociation results obtained by Chun et. al. show that the three

phase (H-Lw-V) systems containing sucrose and glucose are shifted from the equilibrium curve

of the pure R22-water system [31]. This shows that sucrose and glucose exhibit an inhibiting

effect.

Recently, xylomannan, a non-protein thermal hysteresis-producing antifreeze, has been

discovered in the freeze tolerant Alaskan beetle Upis cerambodies [28]. The thermal hysteresis

factor isolated from the Alaskan Upis bettle produced 3.7 +/- 0.3 oC of thermal hysteresis at

5mg/ml, comparable to that of the most active insect antifreeze proteins [28]. Compositional and

structural analyses indicate that this antifreeze contains a β-mannopyranosyl-(14) β-

xylyopyranose backbone and a possible fatty acid component (Figure 5).

Figure 5: Proposed disaccharide core structure of the thermal hysteresis factor isolated from Upis ceramboides

[28].

These antifreeze capabilities, coupled with stability of the sugar and (possible) fatty acid

chain, show great potential as the basis for future hydrate inhibitor development.



1.10 Research Objectives

The objective of the research is to understand the effect of glucose on high pressure gas

hydrate systems, similar to those found in deep sea oil and gas pipelines. This objective stemmed

from the discovery of xylomannan, a non-protein thermal hysteresis-producing disaccharide

produced by the Alaskan Upis cerambodies [28]. Since xylomannan will be difficult to obtain in

useful quantities at the present time, another, more available simple sugar will be tested for its

antifreeze effects with regards to hydrate inhibition.

A methane gas hydrate system and a carbon dioxide gas hydrate system will be tested in

the presence of glucose at various concentrations. Equilibrium plots of these systems will be

produced based on the experimental data acquired.



2.0 Materials and Methods

2.1 Experimental Set-Up

All experiments were carried out in a Jefri-DBR Phase Behaviour System (Oilphase-

DBR-Schlumberger), as used by Beltran et. al. [35] to acquire data for Equilibrium Studies for

the System Methane + Carbon Dioxide + Neohexane +Water (Figure 6). The system contains a

high-pressure PVT cell built with a 20 cm tall glass cylinder (150 cm3 void volume) secured

between two full length viewing windows, inside a stainless steel frame. The full length viewing

windows allow for the constant monitoring of the contents of the reactor. An automated, high-

pressure, positive displacement pump (Oilphase-DBR-Schlumberger) [35] was used to regulate

the experimental pressure inside the reactor [35]. The fluid inside the pump comes into contact

with a floating stainless steel piston, which isolates the hydraulic fluid from the process side of

the cell. This allowes the piston to exert a pressure on the experimental contents without

contaminating the process side with the pump fluid. The high-pressure PVT cell is mounted

inside a temperature-regulated air bath by a bracket and horizontal shaft, attached to an electric

motor. The motor powers the shaft, allowing the cell to oscillate through sixty degrees about its

center of gravity, at a rate of forty oscillation cycles per minute [35].



Figure 6: Experimental set-up.

The pressure and temperature inside the high-pressure PVT cell was monitored using a

platinum RTD probe and a pressure transducer (both supplied with the phase behaviour system).

Using a coverage factor of k = 2 and assuming the corresponding standard uncertainty had a

normal distribution, each expanded uncertainty was estimated to be UT = 0.2 K and UP = 14 kPa,

for temperature and pressure respectively [35].



2.2 Materials

The methane and carbon dioxide gases used during the equilibrium experiments were

both purchased from MEGS Specialty Gases with research grade purity specifications of

>99.995%. The D-(+)-Glucose used during these same experiments was purchased from Sigma-

Aldrich Canada Ltd with a purity rating of ≥99.5 % (GC). For all experiments, the reverse

osmosis water was used as the base for the aqueous phase in the PVT cell.

2.3 Experimental Procedure

The process side of the PVT cell was loaded with 30 cm3 of 0 wt%, 10 wt%, 20 wt% and

30 wt% glucose solution, as required by each set of experiments. These solutions were prepared

using a Denver Instrument S-4002 top loading scale. Once these solutions were injected into the

process side of the PVT cell, the gas being used in the experiment was brought into contact with

the glucose mixture, then pressurized to a value just below the expected hydrate forming region

and left overnight to saturate and equilibrate. Agitation of the PVT cell began once all the

necessary chemicals were loaded into it. In order to induce hydrate crystallization, the positive

displacement pump was used to move the isolation piston forward, thus increasing the pressure

of the system above the equilibrium value for pure water. Once hydrates were formed, the

system was allowed to equilibrate (at least 2.5 hours) and pressure, temperature, system volume,

and hydrate presence in the bulk were monitored. The liquid (aqueous) + hydrate + gas

equilibrium conditions were measured using the isothermal pressure-search method [37, 38, 39].

This pressure search technique satisfies the phase rule for the systems in question . The system’s

pressure was decreased in steps of 0.05 MPa, allowing for the cell’s temperature to equilibrate

for one hour between each pressure decrease. In order to decrease the pressure in the cell, the



positive displacement pump was used to retract the isolation piston located within the PVT cell.

Care was taken to avoid large temperature fluctuations, maintaining the cell temperature within ±

0.4 K of the original value. As carried out by Beltran et. al. [35], this procedure was repeated

until visual confirmation of hydrate disappearance was achieved; the equilibrium temperature

and pressure conditions were then taken as the mean of the conditions at which the final hydrate

crystal was observed and the conditions at which the final hydrate crystal disappeared.

The estimated standard uncertainties were as follows: for temperature uT = 0.2 K, for

pressure uP = 0.03 MPa. Since the coverage factor used is k = 2 and assuming the corresponding

standard uncertainty had a normal distribution, each expanded uncertainty was estimated to be

UT = 0.4 K and UP = 0.06 MPa [35].

2.4 Experimental Conditions

All experiments were performed around the hydrate forming region in order to correctly

identify the equilibrium conditions of the hydrate systems in question. Temperatures commonly

experienced by underwater gas pipelines were chosen as a basis for data point acquisition due to

the isothermal pressure-search method being employed (temperature was set, pressure was

varied). Due to the isothermal nature of the procedure, temperature was not allowed to fluctuate

by more than 0.4 K. For both the methane and carbon dioxide systems, data points were acquired

at 2, 4, 6, and 8 K. Replicates of the endpoints were acquired for the methane/glucose system

while replicates for all four temperatures were acquired for the carbon dioxide/glucose system in

order to validate the accuracy of the results.



3.0 Results and Discussion

3.1 Methane Experiments

The accuracy of the data acquired using the previously described equipment was

confirmed through a comparison of equilibrium data obtained by Deaton and Frost [36] in a 0

wt% glucose system (Table 1). The data published by Deaton and Frost was within the

experimental uncertainty of the presented work within the same temperature range. This data is

presented in Table 1 as well as in Figure 7. It is used as a lower system boundary to show the

inhibition effect of the aqueous solutions used in this work. Note that an inhibition effect is

observed when hydrate equilibrium conditions are shifted to higher pressures and/or lower

temperatures due to the presence of inhibitor(s) in aqueous solution [16].

Table 1: Hydrate-liquid-vapour equilibrium data from this work. Temperature T, Pressure p, for the system CH4-

0 wt% glucose/H20 under H-Lw-V equilibrium.

T / K p / MPa

275.80 3.35

277.44 3.98

279.47 4.85

281.14 5.81



Figure 7: Measured equilibrium conditions of the water + glucose + methane system.

The experimental H-Lw-V hydrate equilibrium pressure and temperature conditions are

presented in Tables 1, 2, 3, and 4 based on their respective glucose concentrations, as well as in

Figure 8. It is observed in Figure 8 that systems containing glucose have equilibrium values

shifted above the pure water (0 wt% glucose) + methane system values. Also, as the glucose

concentration increases between systems, the respective equilibrium curves shift further above

the pure water + methane system equilibrium values. This indicates that glucose has an inhibiting

effect on methane hydrate systems and that as the concentration of glucose increases, so does the

degree of inhibition. There seems to be an exponential increase of inhibition as the glucose

concentration increases between systems, but due to a lack of detailed modelling, this claim

cannot be substantiated with any significant statistical data.

3.00

3.50

4.00

4.50

5.00

5.50

6.00

6.50

275 276 277 278 279 280 281 282

p / M

pa

T / K

▬, 0 wt% glucose, this work;

X, 0 wt% glucose, Deaton and Frost [1];





T / K p / MPa

275.35 3.43

275.50 3.48

277.34 4.23

279.46 5.12

280.91 5.87

281.25 6.02



T / K p / MPa

275.29 3.81

275.36 3.85

277.32 4.71

279.20 5.67

280.87 6.61

280.94 6.60



T / K p / MPa

275.23 4.43

275.25 4.43

277.35 5.46

279.31 6.65

280.83 7.74

280.88 7.77



Figure 8: Measured equilibrium conditions of the water + glucose + methane system.

Experiments were not run for aqueous solutions containing more than 30 wt% glucose

since the maximum solubility of glucose in pure water within the temperature range used is

around 40 wt% glucose [40]. Therefore, in order to prevent glucose crystallization in the PVT

cell, a 30 wt% glucose aqueous solution was used as an upper boundary for data acquisition.

Figure 9 shows the relative inhibitive strength of glucose as compared to other simple

common thermodynamic inhibitors, methanol (CH3OH) [41] and magnesium sulphate (MgSO4)

[42]. A 30 wt% glucose solution shows a slightly stronger inhibitive action to hydrate formation

in comparison to methanol in low concentrations as well as to magnesium sulphate in

approximately equal proportions to glucose. When methanol is present at 35 wt% [41], it shows

a stronger inhibitive effect on hydrate formation than does glucose.

3

4

5

6

7

8

9

275 276 277 278 279 280 281 282

p / M

pa

T / K

▬, 0 wt% glucose, this work X, 0 wt% glucose, Deaton and Frost, 1946 ♦, 10 wt% glucose ■, 20 wt% glucose ▲, 30 wt% glucose



Figure 9: Measured equilibrium conditions of the water + glucose + methane system and reported equilibrium

conditions of the water + methanol + methane system, water + MgSO4 + methane system and water + methanol +

methane system.

0

2

4

6

8

10

12

250.0 255.0 260.0 265.0 270.0 275.0 280.0 285.0 290.0

p / M

Pa

T / K

■, 30 wt% glucose, this work;

♦, 4.2 wt% CH3OH, [41];

▬, 28 wt% MgSO4, [42];

X, 35 wt% CH3OH, [41];



3.2 Carbon Dioxide Experiments

As was carried out in the methane experiments, after the PVT cell was cleaned out and

loaded with the new components required for the carbon dioxide experiments, the accuracy of

the measurements was again observed by confirming the data published by Deaton and Frost for

a 0 wt% glucose/H2O carbon dioxide hydrate system. These results are presented in Table 5 and

have been plotted in Figure 10.

Table 5: Hydrate-liquid-vapour equilibrium data from this work. Temperature T, Pressure p, for the system CO2-


T / K p / MPa

275.37 1.61

277.49 2.05

279.21 2.47

280.98

3.16

275.24 1.55

277.16 1.92

279.13 2.43

281.10 3.23

Figure 10 compares the data acquired from my experimental set-up with the data

published by Deaton and Frost in 1946 [36]. This data lies in agreement with the data published,

thus proving the validity of future experiments with carbon dioxide. As in the methane

experiments, this data is used as the lower boundary for the carbon dioxide hydrate system due to

the lack of inhibiting glucose present. All future experimental results for systems containing a

glucose solution as the aqueous phase are expected to lie above the curve presented in Figure 10,

thus demonstrating glucose’s inhibiting effect on hydrate equilibrium.



Figure 10: Measured equilibrium conditions of the water + glucose + carbon dioxide system.

Experiments were then carried out with aqueous glucose solutions containing 10 wt%, 20

wt%, and 30 wt% glucose. The experimental data pertaining to these experiments are presented

in Tables 6, 7, and 8 respectively. A plot of this data (Figure 11) has been produced and will be

subsequently analyzed.

0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

5

272 274 276 278 280 282 284

p / M

Pa

T / K

▲, 0 wt% glucose, this work;

─, 0 wt% glucose, Deaton and

Frost [1];





T / K p / MPa

275.16 1.66

277.14 2.09

279.14 2.71

281.06

3.50

275.20 1.64

277.16 2.12

279.12 2.67

281.15 3.61



T / K p / MPa

275.12 1.91

277.10 2.43

279.01 3.12

280.61

3.95

275.06 1.91

277.05 2.41

279.09 3.13

280.55 3.92



T / K p / MPa

275.35 2.40

277.25 2.98

278.98 3.74

275.34 2.34

277.21 2.95

278.92 3.71



As predicted, all carbon dioxide hydrate experiments carried out in the presence of a

glucose aqueous solution did in fact reside above the pure water/carbon dioxide hydrate

equilibrium line (Figure 11), indicating that glucose does have a distinct inhibiting effect on

hydrate equilibrium in the presence of carbon dioxide. While the inhibitive effect of glucose is

relatively low at the 10 wt% concentration level, glucose does show a much stronger inhibitive

effect as its solution concentration increases. As was also experienced in the methane hydrate

system discussed previously in section 3.1, there seems to be an exponential increase in degree

of inhibition as the glucose concentration increases, but due to a lack of detailed modelling, this

claim cannot be substantiated with any significant statistical data.

Again, experiments were not run with aqueous solutions above 30 wt% glucose in order

to avoid crystallization in the PVT cell and experimental set-up tubing. A more detailed

explanation for this decision can be found in section 3.1.

It is important to note that there are no experimental values for hydrate equilibrium

around 281 K. This is due to the inability to form hydrates within the system at that temperature.

While it is not impossible to form carbon dioxide hydrates at 281 K, I was unable to induce

stable hydrate nucleation, possibly due to a lack of pressure. As was done for all other

experiments, the system was brought to a pressure above the assumed equilibrium point to

induce hydrate nucleation. Once hydrate growth began, the system was slowly brought down to a

pressure near the equilibrium point, where it was allowed to reach steady state before the

stepwise decreases in pressure commenced. At 281 K, the system was allowed to reach pressures

in excess of 10 MPa, far above the assumed equilibrium range of 4.9-5.2 MPa, yet hydrate

nucleation still did not occur.



Figure 11: Measured equilibrium conditions of the water + glucose + carbon dioxide system.

The relative strength of glucose as an inhibitor in a carbon dioxide hydrate system is

shown in Figure 12. In this figure, glucose is compared to other thermodynamic inhibitors

commonly used in the oil and gas industry: alcohols and salts. The alcohols used in the

comparison are methanol and ethanol, while the salts used are sodium chloride and potassium

chloride.

From the figure, it is clear that the best performing inhibitor of those being compared is

the 10 wt% methanol (CH3OH) solution. The 30 wt% glucose solution tested seemed to have the

same relative inhibitive strength as the 10 wt% potassium chloride (KCl) solution and the 10

0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

5

272 274 276 278 280 282 284

p / M

Pa

T / K

▲, 0 wt% glucose, this work;

─, 0 wt% glucose, Deaton and

Frost;

X, 10 wt% glucose, this work;

*, 20 wt% glucose, this work;

♦, 30 wt% glucose, this work;



wt% ethanol (C2H5OH) solution, while displaying a stronger inhibitive strength than the 5 wt%

sodium chloride (NaCl) solution.

While the 30 wt% glucose solution does show the same inhibitive strength as some of the

other common thermodynamic inhibitors, it is being used in a higher concentration than the

others, making it a less effective inhibitor on a per mass basis.

Figure 12: Measured equilibrium conditions of the water + glucose + carbon dioxide system and reported

equilibrium conditions of the water + methanol + carbon dioxide system, water + ethanol + carbon dioxide

system, water + potassium chloride + carbon dioxide system, and water + sodium chloride + carbon dioxide

system.

0

0.5

1

1.5

2

2.5

3

3.5

4

268 270 272 274 276 278 280 282

p / M

Pa

T / K

▬, 30 wt% glucose, this work; ▲, 10 wt% CH3OH [43]; X, 10 wt% C2H5OH [43]; ♦, 10 wt% KCl [44]; ■, 5 wt% NaCl [44];



4.0 Conclusions and Recommendations

4.1 Conclusions

Experiments were conducted to measure the effect of aqueous D-(+)-glucose on hydrate

formation in the presence of methane and carbon dioxide. The accuracy of the equipment used

was verified by reproducing data for a pure water + methane system as well as a pure water +

carbon dioxide system found in the literature.

Experiments were conducted near the assumed hydrate equilibrium conditions of both the

methane and carbon dioxide systems. The experimental pressures ranged from 3.4 to 7.8 MPa for

the methane systems and 1.6 to 3.8 MPa for the carbon dioxide systems, with an experimental

temperature range of 275.15 to 281.25 K and glucose mass fractions of 0.1 to 0.3.

In both the methane and carbon dioxide systems, an inhibiting effect was observed at all

glucose concentrations. The inhibiting effect of the glucose was then perceived to increase as the

glucose concentration in the aqueous solution increased. All systems containing glucose were

also found to follow the same trend as the systems containing pure water in terms of the apparent

curvature of the equilibrium line.

In comparison to other thermodynamic hydrate inhibitors, glucose did not perform very

well on inhibition strength to mass fraction ratio. The 30 wt% glucose solution used was the only

solution of the three tested to yield a hydrate inhibition level at par with the other thermodynamic

inhibitors considered and employed in industry (10 wt% KCL, 5 wt% NaCl, 4-10 wt% methanol,

10 wt% ethanol). As seen in Figure 8, if methanol is present in the same mass fraction as

glucose, its inhibitive effects on hydrate equilibrium are much greater.



4.2 Recommendations

Experimentation with sugars as thermodynamic inhibitors should be further pursued with

sugars of increasing complexity since they have been proven to be a viable hydrate inhibitor; one

that is readily available, thermally stable, environmentally friendly, and inexpensive when

compared to other types of synthetic hydrate inhibitors being used, such as anti-freeze proteins

and surfactants. Perhaps if sugars of longer length or containing side chains were employed,

greater inhibition could be achieved, and at lower sugar concentration.

Future studies on the equilibrium of hydrate systems in the presence of these sugars

should be carried out with gas mixtures that more closely duplicate the conditions present in

natural gas pipelines in terms of gas phase composition.

Morphology experiments can also be conducted for hydrates systems using sugars as an

inhibitor to get an insight as to how the sugar molecules interact with the hydrate cages. A better

understanding of the physical interaction between sugars and hydrates at the molecular level will

lead to a greater the ability to synthesize a very effective sugar molecule for hydrate inhibition.



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Publications

1. Bruusgaard, H., Carbone, A. & Servio, P., “H-Lw-V equilibrium measurements for the

CH4+C2H6+H2O hydrate forming system”, J. Chem.Eng. Data, 55 (9), 3680–

3683, 2010.

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