Fluorine-enhanced thermal oxidation of silicon
Transcript of Fluorine-enhanced thermal oxidation of silicon
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Lehigh UniversityLehigh Preserve
Theses and Dissertations
1985
Fluorine-enhanced thermal oxidation of silicon /Christine H. WolowodiukLehigh University
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Recommended CitationWolowodiuk, Christine H., "Fluorine-enhanced thermal oxidation of silicon /" (1985). Theses and Dissertations. 4537.https://preserve.lehigh.edu/etd/4537
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FLUORINE-ENHANCED THERMAL OXIDATION OF SILICON
by
Christine H. Wolowodiuk
A Thesis
Presented to the Graduate Committee
of Lehigh University
in Candidacy for the Degree of
Master of Science
in
Metallurgy and Materials Engineering
'Lehigh University
1985
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CERTIFICATE OF APPROVAL
This thesis is accepted in partial fulfillment of
the requirements for the Degree of Master of Science.
~ -/- 85 Date
2)4-<~ A~Yl~ D~partment Chairman
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ACKNOWLEDGEMENTS
My sincere appreciation is extended to my advisor,
Dr. Ralph Jaccodine, for his guidance, assistance, and
encouragement. Many thanks must also go to Taeho Kook,
who acted as a second advisor throughout the various
stages of my research.
I would like to extend my deepest gratitude to Mr.
Fred Stevie and Mr. Peter Kahora of AT&T Bell Labs, who
went above and beyond the call of duty to perform the
SIMS analysis of my oxides. Further, I would like to
thank Dr. Rick Herman for his vapor pressure measure
ments.
A big thank you must go to everyone in the Sherman
Fairchild building, who made working there a pleasure.
I would especially like to thank Phill Goldman for
teaching me everything there is to know about computers,
Patrick McCluskey for babysitting my flowmeter, Philip
Wong and Tom Krutsick for metallizing my samples, Floyd
Miller for supplying answers to all my clean room
questions, Bob Vogel for his general chemistry
knowledge, and Jeanne Loosbrock for being the best
English major and typist in the world.
Mary Ellen deserves a special thanks not only for
doing my C-V measurements, but also for being a great
roommate and dessert maker. June Turkanis also gets a
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big thanks for countless •Dynasty and ice cream• parties
at her apartment, and for being a great lunch buddy.
FinallY, a very special thanks must go to AndY,
who always pushed me to try my hardest and do my best,
and to my parents, whose love, patience, understanding
and encouragement made' all of this possible,
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TABLE OF CONTENTS
TITLE PAGE
CERTIFICATE OF APPROVAL
ACKNOWLEDGEMENTS
TABLE OF CONTENTS
LIST OF TABLES
LIST OF FIGURES
ABSTRACT
1,0 INTRODUCTION
Page
i
ii
iii
V
vii
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1
3
6 2.0 BACKGROUND
2.1 Theoretical Oxidation Model 6
2.2 Experimental Background Investigations with HCl/02 and c1
2;o
2 10
2,3 Thermodynamic Analysis of the Cl-H-0 Ambient 20
2,4 Chlorine Profiles in Silicon Dioxide 23
2.5 Investigations with Flourine 26
3,0 EXPERIMENTAL PROCEDURE
3,1 Description of Apparatus
3,2 General Oxidation Procedure
3,3 Oxidations with c2H3c1 2F
3,4 Oxidations with NF 3
3,5 SOLGAS
3,6 SIMS Analysis
3,7 Electrical Characterization
V
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4.0 RESULTS 45
5.0 DISCUSSION 80
5.1 Standard Dry Oxidations 80
5.2 Oxidations with c2H3c12F 80
5.3 Oxidations with NF3 85
5.4 SOLGAS 88
5.5 SIMS Analysis 91
5.6 Electrical Characterization 93
6.0 SUMMARY 97
1.0 RECOMMENDATIONS FOR FUTURE RESEARCH 100
REFERENCES 102
VITA 106
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Number
I
II
III
IV
V
LIST OF TABLES
Title
Experimental Conditions Used to Grow Oxides and Oxide Thicknesses of SIMS Samples.
Experimental Conditions Used to Grow Oxides and Oxide Thicknesses of C-V Samples.
Oxide Thicknesses of the c2H3c1 2F Oxidations.
SOLGAS Results of Oxidations at 900°c with 0.11 vol% c2H3c1 2F Added to the Ambient.
SOLGAS Results of Oxidations at 900°c with 0.11 vol% NF 3 Added to the Ambient, in the Presence of 0.005 vol% H2.
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VI SOLGAS Results of Oxidations at 900°c with 0.11 vol% HF Added to the Ambient. 74
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Number
2.1
2.2
2.3
2.4
2.5
2.6
2.7
2.8
2.9
2. 10
2. 11
2. 12
2. 13
LIST OF FIGURES
Ti t 1 e fgg_g_
Deal-Grove model for the oxidation of silicon. 7
Comparison of general relationship (solid line), its two limiting forms (dotted lines), and experimental data compiled to date. 11
Oxidation rate of (100) Si at 1150°c in the presence of HCl and c12• 12
Schematic of the Si02 structure. 14
Oxide thickness vs. oxidation time for the oxidation of n-type Si in various HCl/02 mixtures at 900°c. 16
Parabolic rate constant as a function of HCl concentration for (111) and (100) Si at goo, 1000, and 1100°c. 18
Linear rate constant as a function of HCl concentration for (111) and (100) Si at goo, 1000, and 1100°c. 19
Arrhenius plot of the parabolic rate constant for (111) Si oxidized in various HCl/02 mixtures. 21
Arrhenius plot of the linear rate constant for (lll) Si oxidized in various HCl/0 2 mixtures. 22
Equilibrium partial pressures in HCl/02 ambients vs. temperature and H/0 or Cl/0 atom ratios. 24
Parabolic oxidation rate as a function of Pei for HCl and c1 2 oxidations. 2 25
NBS profile of a 9 vol% HCl oxide, 780 ~ thick, grown at 1150°c. 27
NBS profile of a 2 vol% Cl~ oxide, 640 ~ thick, grown at 1000 C, 28
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2 .14
2. 15
3. 1
3.2
3.3
4. 1
4.2
4.3
4.4
4.5
4.6
4.7
4.8
AES profile of a 9 vol% HCl oxide, 1000 i thick, grown at 1150°c.
Oxide thickness as a function of NF 3 concentration for 30 minute oxida-tions.
Side view of oxidation furnace.
Temperature controller used for oxidation furnace.
Gas supply system for oxidation (a) in pure oxygen, (b) with a liquid fluorine source, and (c) with a gaseous fluorine source.
Oxide thickness as a function of oxidation time for standard dry oxides grown at 900°c.
Oxide thickness as a function of oxidation time for standard dry oxides grown at 1000°c.
Pinholes in Si0 2 grown for 12 hours at 1000°c with 0.055 vol% c2H3Cl2F (lOOX magnification).
Oxide thickness as a function of oxidation time for 0% c2H3Cl2F·
Oxide thickness as a function of oxidation time for 0.011 vol% c2H3c1 2F.
Oxide thickness as a function of oxidation time for 0.055 vol% c2H3c1 2F.
Oxide thickness as a function of oxidation time for 0.11 vol% c2Hfl2F.
Oxide thickness as a function of oxidation temperature for 1 hour oxidations.
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4.9
4 .10
4.11
4.12
4.13
4.14
4.15
4.16
4.17
4.18
4.19
4.20
Oxide thickness as a function of
oxidation temperature for 2 hour
oxidations.
Oxide thickness as a function of
oxidizing temperature for 4 hour
oxidations.
Oxide thickness as a function of
oxidizing temperature for 8 hour
oxidations.
Oxide thickness as a function of
oxidizing temperature for 12 hour
oxidations.
Oxide thickness as a function of
c2~iClzF concentration for 1 hour
oxiaations.
Oxide thickness as a function of
c2~iClzF concentration for 2 hour
oxiaations.
Oxide thickness as a function of
c2~1c12F concentration for 4 hour
oxiaations.
Oxide thickness as a function of
c2~1c12F concentration for 8 hour
oxiaations.
Oxide thickness as a function of
c2~1c12F concentration for 12 hour
oxiaations.
Oxide thickness vs. oxidation time
for the oxidation of lightly doped
silicon in various gas ambients at
1000°c.
Linear rate constant vs. vol%
c2tt 3c12F and HCl in o2 for the oxida
tion of lightly doped (100) silicon
at 900 and 1000°c.
Parabolic rate constant vs. vol%
CzH3Cl 2F and HCl in o2 for the oxida
tion of lightly doped (100) silicon
at 900 and 1000°c.
X
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4.21
4.22
4-.23
4.24
4.25
4.26
4.27
4.28
Oxide thickness as a function of fluorine concentration for 2 hour oxidations at goo 0c; comparison of liquid and gaseous fluorine sources. Oxide thickness as a function of fluorine concentration for 2 hour oxidations at 1000°c; comparisott of liquid and gaseous fluorine sources. Oxide thickness as a function of NF 3 concentration for 30 minute oxidations at 700 and 900°c.
Fluorine and oxygen profiles of an oxide grown for 1 hour at goo 0 c with 0.011 vol% c 2H3 c12F in o2 (Sample C from Table I).
Fluorine and oxygen profiles of an oxide grown for 2 hours at 1000°c with 0.011 vol% C2_H:~c12F in o2 (SamP 1 e F from T ab 1 e -i: r. Fluorine and oxygen profiles of an oxide grown for 2 hours at 900°c with O.Oll vol% NF3 in o2 (Sample I from Table I).
High frequency C-V curve of a "nonideal" fluorinated oxide (Sample #7 from Table II).
High frequency C-V curve of an "ideal" fluorinated oxide (Sample #5 from Table II).
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ABSTRACT
Experiments were carried out in order to study the
effect of fluorine additions to a dry oxidation ambient.
It was found that very small concentrations of fluorine,
up to 0.132 vol%, yielded significant enhancements in
the oxidation rate. The enhancements achieved were
greater than those observed with 10 vol% HCl. The
increase was evident in both the linear and parabolic
rate cons tan ts.
Liquid dichlorofluoroethane (C 2H3c1 2F) and gaseous
nitrogen trifluoride (NF 3) were the two fluorine sources
investigated. It was experimentally determined that the
C2H3c1 2F additions caused relatively larger increases
than the NF 3 additions. This is because various
chlorine-bearing compounds which are known to enhance
the Tate, such as HCl and c1 2, were present in the
c2tt 3c1 2F-02 system, whereas they were absent in the NF 3-
o2 system. Also the amount of water present in the
c2H3c12F-02 system was much larger than in the NF3-o2
system.
The relative enhancement was found to increase with
oxidation time, oxidation temperature, and fluorine
concentration, up to the point where the competitive
etching process took over.
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Thermodynamic calculations determined that the
active species causing the significant increases was
hydrogen fluoride (HF). This supported the observation
that c2H3c12F resulted in larger enhancements than the
NF 3 , i.e., hydrogen and fluorine were readily available
in the c2H3c1 2F-02 system, whereas in the NF 3-o2 system
there was only a small amount of hydrogen present with
which to form HF,
Secondary Ion Mass Spectrometry (SIMS) was performed
to determine the fluorine profiles in the oxide layer,
c2H3c12F oxides displayed peaks at the silicon-oxide
interface, while NF 3 oxides contained flat fluorine pro
files. The difference was explained based on hydroxyl
group replacement of the bonded fluorine.
High frequency and quasistatic capacitance-voltage
(C-V) and bias-temperature-stress (BTS) curves were used
to ~lectrically characterize the samples. Not enough
samples were tested to report general trends, but it was
determined that fluorinated oxides could be at least as
reliable as standard dry oxides.
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1.0 INTRODUCTION
Building a single chip on a silioon substrate can
involve as many as 150 to 200 complex processing steps.
These steps can be grouped into processing categories,
such as oxidation, diffusion, photolithography, metalli
zation, etc., some of which are repeated many times
during the manufacturing cycle of silicon devices. The
process that this report will focus on is oxidation,
specifically thermal oxidation, which is the principal
method used in integrated circuit (IC) manufacture.
Silicon dioxide layers on silicon have many uses
(l)~ Besides acting as a critical component in MOS
structures, namely as gate insulators, they can be used
to protect certain areas against diffusion or implanta
tion of dopant into the substrate. Si02 layers are also
commonly used to passivate the surface, Finally, they
are used for isolation purposes -- to isolate one device
from another or to provide electrical isolation of multi
level metallization systems.
In today's state-of-the-art IC processing, the
trend has been towards lower temperature processing.
The reason for this is that high temperature can degrade
any characteristics achieved during previous processing
steps, or it can make subsequent processing steps more
difficult. For example, carefully determined diffusion
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profiles can be altered dramatically by exposure to high
temperature. Another example is substrate warpage that
occurs during high temperature processing. The warpage
can make the definition of fine features during upcoming
photolithography steps virtually impossible. Finally,
the lower temperatures can retard the nucleation and/or
growth of stacking faults.
Silicon dioxide layers are required in IC manufac
ture in a range of thicknesses -- from electrically
reliable thin gate oxides to thick isolation oxides.
However, at low temperatures unreasonably long times are
often required to grow the oxides to the thicknesses re
quired. It has been found that the addition of chlorine
to the oxidizing ambient not only enhances the oxidation
rate, but also leads to beneficial electrical, chemical,
and microstruc tural properties (2, 3). Chlorine addi
tives improve the electrical characteristics of the Si02
by trapping and neutralizing sodium ions in the oxide
(4,5). They also enhance the minority carrier lifetime
in the silicon substrate by gettering heavy metal im
purities (6,7). Finally, they lower the density of
interface states ( 8) and suppress the formation and
growth of stacking faults (9).
Since chlorine has been proven to enhance the
oxidation rate, and since it is generally believed that
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the enhancement is due to the chemical affinity of
chlorine, it is reasonable to assume that fluorine may
behave similarly. The purpose of this thesis is to show
that fluorine does in fact enhance the oxidation rate,
without degrading the electrical characteristics. The
oxidations were performed with various fluorine com
pounds, and the results were analyzed in terms of the
oxidation rate constants and thermodynamic calculations.
SIMS has been performed to determine the fluorine profiles
throughout the oxides, and C-V measurements also serve to
check the insulating properties of the oxides.
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2. 0 BACKGROUND
2.1 Theoretical Oxidation Model
Silicon dioxide layers on silicon can be obtained
by a number of different techniques, including chemical
vapor deposition (CVD), anodization, plasma reaction,
and thermal growth (10). Thermal oxidation is a process
whereby silicon reacts with oxygen or water at elevated
temperatures to produce a layer of Si02• The reactions
which take place are represented by the following
simplified reaction equations:
Si + o2 + Si02
Si+ 2H20 + Si02 + 2H 2
(2.1)
(2.2)
Marker experiments (11) have demonstrated that the
oxidizing species diffuses through any oxide present and
then reacts with the silicon at the Si-Si02 interface.
In this manner, the Si-Si02 interface is always moving
into the silicon. Based on this idea, Deal and Grove
(12) derived a kinetic model for the thermal oxidation
of silicon. Assuming steady-state conditions, they
equated the three fluxes, i.e., the transport of the
oxidant to the oxide surface, the diffusion through the
oxide, and· the interface reaction, to determine their
first order model. This model is represented by Fig.
2 .1.
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GAS OXIDE
~---X -- 0
c·k_ - - ~
i C
: X
-c. l.
SILICON
F _L~
Fig. 2.1. Deal-Grove model for the oxidation of silicon.
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The flux of the oxidizing species to the outer
oxide may be expressed as
F1 : h(C* - Co) (2,3)
where his a gas-phase transfer coefficient and C * is the
equilibrium bulk concentration of the oxidant in the
oxide. C0
is defined as the concentration of the oxidant
at the outer oxide surface, as determined by manipula
tions of Henry's law and the ideal gas law.
The flux of the oxidant going through the existing
oxide layer is calculated from Fick's First Law of
diffusion:
F2 = -Deff(dC/dx) (2.4)
where Deff is the effective diffusion coefficient and
dC/dx is the concentration gradient of the oxidant in the
oxide. This equation can be simplified to
F2 Deff (Co-Ci)
(2.5) = XO
where Ci is the concentration gradient of the oxidant at
the interface and x0
is the oxide thickness.
The third flux, i.e., the interface reaction flux,
is given by
(2.6)
where k is the interface reaction rate constant.
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Under steady-state conditions, these three fluxes
a~e equated to obtain
kC* k kx
1 + + __£ h Deff
( 2, 7)
Further manipulations of this equation lead to the
form
where
x~ + Ax 0 = B(T + 1)
A - 2 Derr (1/k + 1/h)
B - 2 DefrC*/N1
1 - (xI + Axi)/B
(2.8)
where N1 iS the number of oxidant molecules incorporated
into a unit volume of oxide layer, xi is the initial
oxide thickness, and 1 is a shift in the time coordinate
which accounts for the presence of an initial oxide
layer.
This equation has two limiting forms. For long
oxidation times, it reduces to
x~ = Bt (2,9)
where diffusion to the interface is the rate-limiting
step. Bis defined as the parabolic rate constant. For
short times, equation (2.8) simplifies to
x0 : ! ( t + T ) (2.10)
where the reaction at the interface is the rate-
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controlling step. B/A is the linear rate constant.
This general linear-parabolic relationship is valid
for both wet and dry oxidations over a wide temperature
range (700-1300°C) and for oxide thicknesses between 300
and 20,000 i. It fits the experimental data very well,
as can be seen in Fig. 2.2.
2.2 Experimental Background Investigations Kith HCl/02
.ruW. li2ffi2
Since chlorine-enhanced oxidation of silicon has
been extensively studied during the past decade, and
since fluorine is expected to behave similarly to
chlorine in the oxidizing ambient, experimental
investigations with chlorine are presented. The first of
these chlorine-enhanced oxidation studies was performed
by Kriegler, Cheng, and Colton (13). They oxidized
(100) Si wafers at 1150°c in oxygen containing small
amounts of HCl and c1 2. Their results are summarized in
Fig. 2.3. The thickness range indicated in this figure
corresponds to a parabolic time dependence in pure oxygen
oxidations. Therefore, it is concluded that the addition
of HCl or c1 2 results in an increase in the parabolic
rate constant, B.
When HCl is added to the incoming gas stream, it
reacts with oxygen to produce water and chlorine,
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K)
~ A/2
10-•
0.1
Fig. 2. 2.
H20 02
• 1300 °C 0 • 1200 °C
A • 1100 °C a • 1000°C 0 • 920°C
X 800°C ~ 700°C
10 102 103
t+T
A2/4B
Comparison of general relationship (solid line), it_s two limiting forms
(dotted lines), and experimental data
compiled to date.
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1700
1500
0~ 1400 -ffl ~,::co ~ ~
~ 1200
w 0
~ 1100
JOO()
900
STANDARD DRY 02
15 20 25 30 35 40
OXJDATJON TIME (min)
Fig. 2.3. Oxidation rate of (100) Si at 11S0°c
in the presence of HCl and c12.
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according to the following equilibrium reaction:
4HC1 + o2 ( 2 .11)
Since water results in a higher oxidation rate than
oxygen, one's first impulse would be to attribute the
increase in rate to the water generated in the system.
However, the water itself could not result in such a
large effect as was observed by Kriegler, et al. (13),
Also, for equal amounts of chlorine in the gas phase,
such as with 5% HCl or 2.5% c1 2, larger enhancements were
observed with the c1 2• This supports the fact that
chlorine is the major factor affecting the oxidation
kinetics.
Hirabayashi and Iwamura (14) observed a similar
increase in oxidation rate with an HCl/02 ratio of up to
10 mole% HCl. They attributed the enhanced rate as being
the result of three effects. The first of these was a
structural difference between standard and HCl oxides.
This was in agreement with Kriegler's (15) theory that
during high-temperature oxidation in the presence of HCl,
an Si-0-Cl complex is formed, with the chlorine
substituting for oxygen. The Cl acts as a network
modifier (16), i.e., it forms an ionic bond with an
oxygen ion, resulting in a nonbridging bond, as is shown
in Fig. 2.4 (17). The difference in structure
allows for more rapid diffusion of o2 through the
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() Bridging oxygen
~ Nonbridging oxygen
• Silicon
Fig. 2.4. Schematic of the Si02 structure.
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chlorinated oxide. This leads to an increase in the
diffusion-limi~ed parabolic rate constant, B.
The second of these effects was that the HCl had a
catalytic effect leading to an enhanced reaction at the
Si-Si02 interface. This could possibly be caused by
chlorine breaking the Si-0 bonds at the interface,·
thereby providing more reactive sites. This would cause
an increase in the linear rate constant, BIA. Finally,
the generation of water in the system enhanced the rate
(18).
It quickly became obvious that there was more than
one factor influencing the thermal oxidation kinetics of
HCl oxidations. Therefore, studies were performed by
Hess and Deal (19) to determine the effect of HCl
concentration, oxidation temperature, and silicon
orientation. Figure 2.5 is a plot of oxide thickness vs.
oxidation time at goo 0 c for varying HCl concentrations.
Similar sets of curves were obtained at 1000 and 1100°c.
It can be seen that (111) silicon always oxidizes faster
than (100) over this concentration range, as is the case
for dry oxidation. Another observation is that although
HCl additions cause a systematic increase in the
oxidation rate, there is a relatively large increase
between O and 1%, with the effect decreasing as more HCl
is added.
15
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..--..
E ::L -
(/) (/)
w z ~ u I r-w 0
X 0
-E :i .__..
(/) (/)
w z :::c u -I r-
w 0 -X 0
1.0 (a) (111) Si ORIENTATION
0.1
ll = 10°/o HCI o = 5°/o HCI o = 1 o/o HCI • = 0°/o HCI
0.01 0.1 1.0 10.0 100.0
OXIDATION TIME (hr)
1.0 rrr,1111
(b) (100) Si ORIENTATION
0. 1
6 = 10°/o HCI D = 5°/o HCI o = 1 O/o HCI • = 0°/o HCI
0.01 0.1 1.0 10.0 100.0
OXIDATION TIME (hr)
Fig. 2.5. Oxide thickness vs. oxidation time for
the oxidation of n-typ5 Si in various
HCl/02 mixtures at 900 C. (a) (111) Si
(b) (100) Si 16
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Figures 2.6 and 2.7 09) show the effect of HCl on
the linear and parabolic rate constants, where the
constants include the effects of chlorine, oxygen and
water. Figure 2.6 shows that there is essentially no
orientation effect on the parabolic rate constant. This
is expected, as B should vary only with the partial
pressure of the oxidant and solubility and diffusion of
the oxidant in the Si02. However, at 0-1% HCl additions,
there is a discrepancy between the three temperatures,
which may be due to the varying effects of generated H2o
and c1 2 on the oxidations.
In Fig. 2.7, the linear rate constant is shown to
depend strongly on the orientation, with the effect
decreasing with increasing temnperature. This is also
expected, as BIA involves the reaction at the Si-Si02
interface. Again, there is a discrepancy between 0-1%
HCl and 1-10% HCl additions, which may be caused by
varying oxidant solubility and chlorine etching.
Hess and Deal also determined the activation
energies for the interface and diffusion-controlled
processes. In the standard dry and wet oxidations, an
Arrhenius plot for B results in a straight line, and
activation energy is nearly equal to that for diffusion
of the oxidant in the silica. Similarly, BIA yields a
straight line with activation energy approximately equal
to the bond-breaking energy of a Si-Si bond.· However, as
17
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0.1.---....----..-----,..----....-------·
...... ... .c. ;:;-E ::l ai 1-z j:! U)
z 8 0.01
w ~ a:: u _J
0 CJ <! a:: a:
2 4 6
•,a,i = (111) 0 10 16 = (100)
8 HCI CONCENTRATION (vol. %)
10
Fig. 2.6. Parabolic rate constant as a function of HCl concentration for (111) 8nd (100) Si at 900, 1000, and 1100 C.
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1.0,....----....---...----....----....-----r---,
-... .s:;. ....... E :i.. -~ CD 1000°c
..... z ~ (/) z 0 u 0
w ~ a:: a:: <( w ~ ••• ,, = (111) .J 0
10 16 = (100)
0.001 L..----L---.1..--......L---..i----J.---'
0 2 4 6 B 10
HCI CONCENTRATION (vol. %)
Fig. 2.7. Linear rate constant as a function of HCl concentration for (111) and (100) Si at 900, 1000, and 1100°c.
19
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depicted in Figs. 2.8 and 2.9, Arrhenius plots for HCl/02
oxidations do not yield straight lines. These curved
lines further support the fact that there are a variety
of species affecting the oxidation kinetics, and as the
oxidation proceeds in time, the relative contribution of
these species is changing.
2.3 Thermodynamic Analysis of th§. Cl-H-0 Ambient
Although various researchers had experimentally .,
shown that HCl additions increase the oxidation rate,
none of them were able to fully explain the reason for
it. Tressler, et al. (20) decided to look at the gases
that existed within the furnace under oxidation
conditions, rather than at the gases being fed into the
furnace. At the elevated oxidation temperatures, the
gases react with each other, leading to a final
composition which differs from the input composition.
This final composition can be determined from
thermodynamic calculations.
A modified SOLGAS (21) computer program was used to
calculate the equilibrium partial pressures of the
gaseous species existing in the Cl-H-0 system. The
thermodynamic data (entropy, enthalpy) for all the
species in the system and the elemental molar
concentrations are an input into the program, along with
20
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oc
O.t {l-00 \ ,i:f) ,c(F # #
A
• -.. .c:.
• ' N
E :1.. - 0.01 CD
..... ... z ~ (/)
8 w ~ a:
~ 0.001 A = 10% HCI ...J . : 5% HCI 0 • = 1% HCI CD <! Q: 0% HCI a: <! a..
QOOOl i___--1 __ --1... __ -1-__ ...1_ _ ___,
0.6 0.7 0.8 0.9 1.0 1. 1 1000 T(°K)
Fig. 2.8. Arrhenius plot of the parabolic rate constant for (111) Si oxidized in various HCl/02 mixtures.
21
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-~ .r:. e to :t -~ m
t- 0.1 z
~ 5 u w 0.01 ti 0:
0: <t
~ 0001 :J
(100)(111) ll,A = 5% or 10% HCI o,• = 1% HCI o,• = 0°/o HCI
0.0001 '------'---'-------''---------0.6 07 0.8
1000 T(°K)
09 1.0 1.1
Fig. 2.9. Arrhenius plot of the linear rate constant for (111) and (100) Si oxidized in various HCl/02 mixtures.
22
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the temperature. SOLGAS goes through a series of
iterations to determine the equilibrium conditions based
on the minimization of the free energy of the system
considered. The partial pressures of the various
components obtained in the output are directly
proportional to their concentrations in the furnace.
The results from the Cl-H-0 system are depicted in
Fig. 2.10 (20). R = 0.01 and R = 0.10 correspond to 2
and 20 volume% HCl, respectively. The partial pressures
of the species vary as a function of the incoming HCl/0 2
ratio and temperature.
Looking back to Fig. 2.3, the parabolic rate
constant does not depend on the amount of chlorine input
into the furnace, as 2.5% c1 2 yields a different rate
than 5% HCl, nor does it depend on the total chlorine
content of the oxidizing ambient. But as shown in Fig.
2.11, the parabolic rate constant increases smoothly with
Pei , regar.dless of whether HCl or c1 2 was the incoming
2 gas. This implies that HCl does not directly react
with silicon to form Si02 in HCl/02 oxidations, but
rather through the gaseous reaction products c1 2 and H2o.
2.4 Chlorine Profiles in Silicon Dioxide
Various analytical techniques such as secondary ion
mass spectroscopy (SIMS) (22), Auger electron
spectroscopy (AES) (23, 24), nuclear backseat tering
23
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-E -C, -w a::: => en en w a::: a.
1()-2
I0-3
O.IO=R
10 ....
Fig. 2.10.
HCI
HOCI
O.IO=R
TEMPERATURE (°C)
Equilibrium partial pressures in HCl/0? ambients vs. temperature and H/0 or Cl/0 atom ratios. R = H/0 = Cl/0. Tota124ressure = 1 atm.
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l\) U1
z,,..... 0 Z ,_.. ..... ~ ~ 10 ~N ,_.. ::E: >< ;:l. 0
.::::t U I ,_.. 0 _J .-I 0 .._, i:t:l c:::e w 5 ~ ~ 0- a:::
1150 °C O HCI - 02 OXIDATION
• Cl2 -~ OXIDATION
2 5 10 20
EQUILIBRIUM CLz PRESSURE (10-3 ATM)
Fig. 2.11. Parabolic oxidation rate as a function of Pei
for HCl and c12 oxidations. 2
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spectroscopy (NBS) (25,26), and electron microprobe analysis (EMA) (8) have been used to determine the chlorine concentration profile within the oxide. Typical profiles are illustrated by the NBS results of Figs. 2.12 and 2.13 (26). For HCl oxides, the chlorine is most concentrated within 200 i of the Si-Si02 interface (see Fig. 2.14) (23) and decreases by about one or two orders of magnitude into the bulk of the oxide. In theory, the incorporation of chlorine appears at the oxidation front, implying that Si-Cl bonds form directly at the interface (24).
For c1 2 oxides, the chlorine profile was found to be more evenly distributed throughout the oxide. Also, a higher percentage of chlorine (more than ten times as much) was incorporated into the c12 oxides than in the HCl oxides, for the same concentration of chlorine in the gas phase. For both HCl and c1 2 oxides, the amount of chlorine incorporated in the oxide increased with oxidation time, temperature, and percentage Cl added. No chlorine could be detected in the silicon substrate.
2.5 Investigations .kLl..1h Fluorine
Surface chemistry and the reactions that take place between silicon and fluorine have recently been investigated by Chaung (27). According to this study, an
26
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40
30
CJ') I-2 20 :::i 0 u
10
0
1500
Fig. 2.12.
s·~ 1
S102(Hcn
- SIGNAL (averaged)
--- Si·toiltbockground
\ --- c135
······ c137 . . \ I
. \ . .
I I . SiOz (Cl:35) I . I .
Si02 (Cl:37) r- ~
. \ . \ • \ ' \ ' . ·, . . ,, : ... ,,....._ \:" ... I -----1-----.. ---
, ;·, ' .,J ,• .. ~· ,,. . /'" ... \,,, ..__. ·-·....J'
1550 1600 1650
ENERGY (KeV)
NBS Rrofile of a 9 vol% HCl oxide, 780 ~ thick, grown at 11so
0
c.
27
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50 ~ Si Si02 (Cl2)
- SIGNAL (averaged) --- Si-tail+ background
40 ---ci37 _ ... c135
1/) 30 1-z ::> 0 u
20
10
0
1500
. \ . \ . \ . \ \ ...... \ ./ X
/ \ .. . ... ' . L ~ \ -,. -------,.L----\----.::_ .. __
,I ' /'°' ',, . •' \ . . I ..... • '-·' -· ... . I \ . . / 1550 1600
1650
ENERGY (KeV)
Fig. 2.13. NBS profile of a 2 vol% c12 oxide, 640 R thick, grown at 1000°c.
28
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(I) I-
20
en z 15 I- :> J: .
. ~~ w~ J:...: ~ ~ 10 WW Cl Cl
a:: 0 Wlc, !lC ::, q'. 5 c:t w
Cl.
• APPROXIMATE 'DISTANCE FROM Si-Si02 tJTERFACE, A
200 0 400
CHLORINE (x 20)
4000 5000 6000
SPUTTERING TIME, sec
Fig. 2.14. AES profile of a 9 vol% HCl oxide,
1000 i thick, grown at 1150°c.
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SiF2-like structure is formed upon the exposure of bare
silicon to either XeF2 or SiF4. This structure is
present only on the surface or up to a few atomic layers
deep into the silicon. This type of chemisorbed layer
preceeds the vapor phase desorption of silicon, and is
therefore thought to loosen the structure at the
interface, enhancing the diffusion of oxygen to the
surface and the reactivity of the oxygen at the surface.
Near the end of this study, a Japanese paper was
published, which addressed the specifics of NF 3
additions to a dry oxidation ambient (28). Figure 2.15
shows that the oxidation rate is initially enhanced by
the fluorine additive, and then decreases due to etching
of the oxide by fluorine radicals. The enhancement was
attributed to both enhanced diffusion of oxygen through
the oxide and increased reactivity at the surface,
similarly to the chlorine oxidations, However, much
smaller concentrations of fluorine w~re needed to bring
about the same degree of enhancement.
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40
C/) C/) w :z: G 20 ...... :::c I-
~ 10 ...... >< 0
o eoo•c Q ~' ~ • 1oo•c :~• t:,. 600• C
? I ~ ; ........ ~ I 1 0----o._ ..... _-L\ • -t:r-6-------
)... ---0- I 0---------/ ,' 2- --o----- --,,. ~ I "'- ...... --.. .---L---~
O'-------'-----L...----~---.-...J 400 600 800
0 200
Fig. 2.15. Oxide thickness as a function of NF
3 concentration for 30 minute
oxidations.
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3.0 EXPERIMENTAL PROCEDURE
3.1 Description of Apparatus
The experimental set-up used in this investigation is
depicted in Fig. 3.1. It consists of a hot wall oxida
tion furnace, gas supply system, and exhaust system. The
Thermco Pacesetter II furnace is divided into three heat
ing zones. Each zone is electrically heated and its
temperature controlled by a Dialette W793B master-slave
temperature controller, as shown in Fig. 3.2. Type R
thermocouples (platinum-platinum,13% rhodium) sense the
temperature in the three zones. The data taken by the
thermocouples is input as d-c millivoltages to the con
troller channels. Each channel is referenced to the
center temperature and then independently maintains the
desired temperature profile. The controller is capable
of maintaining temperatures in each zone to within ±1°c.
The reaction tube, end cap, and pushrod were semicon
ductor-grade clear fused quartz. An alumina liner sur
rounded this quartz tube in order to minimize ionic con
tamination or moisture diffusion through the tube. A
vertical-rack quartz boat was used to hold the wafers
perpendicular to the gas flow in the furnace. The verti
cal rack all9wed for fast heating and cooling of the
wafers. Laminar flow was present in the furnace, with the
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GAS EXHAUST
QUARTZ TUBE
---- SILICON WAFERS
QUARTZ BOAT
" " . . . . . . . . . . ' ( ',1 '/ ' ,1 • / • / • / , I • I • / • / • / • / • / • / • /, /
~ , RESISTANCE HEATED FURNACE " ~ , ',1 • , • I •, • I , • ' ' • , ' , • , ' ' ' • ' .
GAS SUPPLY
QUARTZ END CAP
Fig. 3.1. Side view of oxidation furnace.
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TC.I
c:,
0 0 0 0 0 0 0 0 0 0
[g UCTJl U~ ~qp
TC.2 TC.3
TRANSFORMER
CONTROL PANEL
208 /220/240V. ,I ,50/60"'
TC= thermocouple
SCR = silicon controlled rectifier
Fig. 3.2. Temperature controller used for oxidation furnace.
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gas impinging on the backside of the wafers. The loading
chamber met Class 1000 clean room requirements.
The oxygen used had a minimum purity of 99.6%, with
the major impurity being argon. It contained a maximum of
40 ppm hydrocarbons and 9 ppm water. Ultra-high purity
nitrogen, with a minimum purity of 99,99% was also used.
Both gas streams contained Millipore disc filters
(GSWP02500) with a 0.22 micron pore size. The furnace
was well baffled, in an attempt to prevent the back
diffusion of air into the reaction tube.
The gas supply system shown in Fig. 3(a) was used for
oxidation with pure oxygen. Since this investigation
involved oxidation in fluorine, the gas supply system was
modified as shown in Figs. 3(b) and 3(c). Figure 3(b)
depicts the set-up for oxygen bubbling through dichloro
fluoroethane (C2H3c1 2F) whereas Fig. 3(c) shows the set-up
for feeding nitrogen trifluoride (NF 3).
Flowmeters for the N2 and o2 flows were Brooks type
1355-07ClZAA. For the much smaller flow rate of the
fluorine compound, a Gilmont microflowmeter type F-9760
was used. Very small flows in the ml/min range could be
reproduced with this flowmeter's micrometer adjustment
knob. The construction materials of this flowmeter -
teflon, borosilicate glass, and synthetic ruby -- pre
vented its corrosion.
The microflowmeter was pre-calibrated for air and
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w °'
TO FURNACE TO FURN1\CE TO FURNACE
~ CzH3Cl2F
D 02 N2 02 N2 02 N2
I I I
(c)
Fig. 3.3. Gas supply system for oxidation (a) in pure oxygen, (b) with a liquid fluorine source, and (c) with a gaseous fluorine source.
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water by the manufacturer and shipped along with a cali
bration chart. Any other gas flow could be determined
with this flowmeter by using a simple conversion equation
involving the density and viscosity of the gas.
The silicon wafers used were chem-mechanically
polished p-type, Czochralski-grown crystals of (100)
orientation, with a resistivity of 2-10 ohm-cm. Oxide
thicknesses were determined by a Rudolph Research AutoEl
II ellipsometer, which contained a helium-neon laser.
3.2 General Oxidation Procedure
Before any oxidations were performed, a 15.2 cm (6 11 )
flat zone at the center of the furnace had to be
achieved. This was done by placing a Type R thermocouple
into a quartz sheath, and inserting it into the furnace
such that it measured the temperature at the center of
the furnace and 7.6 cm (3 11 ) to either side of the
center. The furnace settings were determined so that the
temperature of the flat zone was constant with time,
reproducible, and controlled to within ±1°c of the tem
perature desired. The settings were obtained individual
ly at temperatures of 700, 800, 900, and 1000°c.
The oxidation procedure consisted of first dialing in
the furnace settings and allowing approximately 2 hours
for the furnace to come to temperature. After this was
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achieved, the wafers were prepared for oxidation. Any
oxide present was removed by etching in a 10:1 solution
of deionized water to hydrofluoric acid several times,
allowing the water in the beaker to overflow each time,
in order to carry any chemical residue. After etching,
the wafers were transported to the furnace loading cham
ber in a beaker with water as a cover, thereby preventing
any oxide from forming in the laboratory environment.
First, standard dry oxidations were performed at 900
and 1000°c, in order to become familiar with the oxidation
process and to see if the oxide thicknesses obtained
paralleled those reported in literature. This process
consisted of turning on the N2 flow at 2 1/min. The
wafers (two wafers per run) were placed 2.54 cm (l")
apart onto a quartz boat and quickly pushed into the
furnace. They were allowed to warm up for 5 minutes in
the hot zone, while the N2 was flowing. The gas flow
was then changed to oxygen at 1 1/min. Oxidations were
performed for 10, 20, and 30 minutes, and for 1, 2, 4,
8, 12, and 16 hours. After oxidizing for a given time
period, the gas flow was changed back to 2 1/min of N2
for purging purposes.
Three points on each wafer were measured with an
ellipsometer in order to determine oxide thickness.
Ellipsometry is a non-destructive, high precision
technique whereby the determination of film thickness is
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based on optical interference of the incoming laser beam.
The thicknesses reported are actually an average of the
six measurements. The reproducibility of the oxides
measured was ±3%. For ellipsometric determination of film
thicknesses less than 200 i, the refractive index was
assumed to be 1.46. In reality, the refractive index was
probably lower, because the fluorinated oxides are
believed to be less dense than the standard dry oxides,
especially at the silicon-oxide interface. Physical
integrity of the oxides was determined by a Nikon optical
microscope at a magnification of lOOX.
The oxidizing gas flow for these fluorinated
oxidations was comprised of 11/min of o2 along with a
much smaller flow of 1,2-dichlorofluoroethane (C2H3c1 2F).
The bubbler filled with the c2H3c1 2F was kept at room
temperature, while oxygen was bubbled through at rates of
1, 5, and 10 ml/min. This corresponded to fluorine
additions of 0.011, 0.055, and 0.110 volume percent,
respectively (assuming that the o2 is saturated with the
compound). It is important to note that the oxygen and
dichlorofluoroethane vapor were assumed to be ideal gases,
and therefore molar percent and volume percent can be used
interchangeably. The oxidations were performed for 1, 2,
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4, 8, and 12 hours, at temperatures of 900 and 1000°c.
3.4 Oxidations Kith NF 3
Nitrogen trifluoride is a gaseous fluorine source, in
which the gas flow was monitored directly by the
microflowmeter. Again, the oxidizing gas flow as 1 1/min
o2 along with the fluorine compound. The NF 3 additions
were 0.011, 0.022, 0.033, and 0.044 vol%, which
corresponded to 0.033, 0.066, 0.099, and 0.132 vol%
fluorine. These oxidations were performed at 900°c for 2
and 4 hours, and at 1000°c for 2 hours.
In order to check the recent experimental results of
Morita, et al. (28), a range of oxidations were performed
for 30 minutes at 900°c. The NF3 additions in this case
were 0.0027, 0.0055, 0.0082, 0.011, 0.022, 0.033, 0.044,
0.055, and 0.066 vol%. Further oxides were grown for 30
minutes at 700°c, with NF 3 additions of 0.011, 0.022,
0.033, and 0.044 vol%.
3.5 SOLGAS
The concentrations of the various chemical species
existing in the oxidation furnace were determined by
SOLGAS. For the dichlorofluoroethane, the input data
consisted of 0.11 vol% c2H3c1 2F in an o2 ambient, for
temperatures of 700, Boo, 900, and 1000°c. For the N~F-0
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system, the program was initially run with 0.11 vol% NF 3
at 900 and 1000°c, and later rerun with the same
conditions along with 50 ppm of hydrogen. The latter was
thought to be more indicative of the actual conditions
within the furnace. The SOLGAS program was also used to
determine which chemical species would be present upon
adding 0.11 vol% HF to the dry ambient at 900°c.
3.6 SIMS Analysis
Secondary Ion Mass Spectrometry (SIMS) was performed
to determine the concentration depth profile of the
fluorine in the fluorinated oxides. The samples that were
analyzed by SIMS are listed in Table I. The primary beam
was rastered across a 500 µm x 500 µm square, forming a
·crater, and the secondary ions were collected from a 150
µm diameter circle at the center of this crater. The
fluorine profile was determined by monitoring 49 SiF+,
which is equivalent to monitoring 19F+. The fluorine
concentrations were not determined on an absolute basis,
but rather on a relative basis.
3.7 Electrical Characterization
High frequency and quasistatic capacitance-voltage
(C-V) and bias-temperature-stress (BTS) were performed
to evaluate the electrical properties of the oxide.
Table II lists the various experimental conditions that
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TABLE I
Experimental Conditions Used to Grow Oxides
and Oxide Thicknesses of SIMS Samples
Identifying Oxidation
Letter Temperature (OC)
A 800
B 800
C 800
D 900
E 1000
F 1000
G 900
H 900
I 900
Oxidation Time (hr)
2
4
8
1
1
2
2
2
4
42
Flourine Oxide
Concentration Thickness
(vol%) (i)
0.11% c2H3c1 2F 395
0.055% c2tt 3c1 2F 519
0.011% c2H3c1 2F 506
0.011% c2tt 3c12F 359
0.011% c2H3c1 2F 1109
0.011% c2H3c12F 1800
0.033% NF 3 1227
0.011% NF 3 887
0.022% NF 3 1750
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TABLE II
Experimental Conditions Used to Grow Oxides and Oxide Thicknesses of C-V Samples
Identifying Oxidation Oxidation Fluorine Anneal Oxide Number Temperature Time Concentration in o2 Thickness
( 0c) (min) (vol% NF3) (min) (~)
1 700 30 25
2 900 30 116
3 700 30 a.on -- 81
4 700 30 a.on 100 92
5 900 30 a.on 190
6 900 30 a.on 60 286
7 700 30 0.044 179
8 700 30 0.044 100 190
9 900 30 0.044 20 629
10 goo 30 0.044 100 887
4.3
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were used to grow the tested oxides. The samples were pre
pared so that they could be compared on the basis of
oxidizing temperature, post-oxidation treatment, and
fluorine concentration added to the oxidizing ambient.
All samples were metallized by filament
evaporation of aluminum. In all instances, front-side
metallization was performed within 4 hours of oxidation,
in order to prevent the oxides from absorbing any
contaminants and moisture from the air. None of the
samples had a post-metallization anneal.
A 1 MHz high frequency C-V measurement was
performed, followed by a quasistatic C-V at a linear
ramp rate of 30 mV/sec. The BTS procedure consisted of
taking a C-V curve, heating the sample to 200°c,
applying a positive bias of 1 MV/cm for 3 minutes,
cooling the sample to room temperature, and taking
another C-V curve. The shift between the two curves was
indicative of the mobile ion concentration.
44
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4.0. RESULTS
The results of the original control (fluorine
free) oxidations at 900 and 1000°c are plotted in Figs.
4.1 and 4.2, respectively. These graphs were used to
determine the linear and parabolic rate constants, which
were found to agree with published data.
The analysis of the fluorinated oxides was carried
out in a similar manner to the published work on chlori
nated oxides, as presented in Chapter 1. The fluorine
oxidations involved varying four parameters--oxidation
temperature, time, fluorine concentration, and fluorine
additive. The raw data for the c2H3c1 2F study is pre
sented in Table III. The data which are boxed in repre
sent poor quality oxides, i.e. oxides which had a high
density of pinholes. An example of these pinholes,
which were most pronounced in high temperature oxida
tions with large amounts of c2H3c1 2F, can be seen in
Fig. 4.3. The data not enclosed in boxes correspond to
good quality oxides, i.e. smooth oxides, free of pin
holes, as examined under lOOX magnification.
Variation of oxide thickness with oxidation time is
shown in Figs. 4.4 to 4.7, for varying concentrations of
c2H3c12F. Figures 4.8 to 4.12 depict the variation of
oxide thickness with oxidizing temperature. Figures 4.13
to 4.17 show the dependence of oxide thickness on
45
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.-~ ;:I . ...... V) (/)
Lu ct'~ ::'!:.:: <._, ,._ ... -·-~~a. ..
1--
tL-1 c.:.: ,-.. >::::'. c::·
0 I 16 r-----r---....,---r----,.,-----,--·---
0. J.2 //
/ I
~/ ! I /F
I t:l' --1 n 08,. .. .. . I I
' I l 1-
' l I r
o,ot
OXIDATION TIME (H)
Fig. 4.1. Oxide thickness as a function of
oxidation time for standard dry
oxides grown at 900°c.
46
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0.4
0.3
,,.-.. ~ ;:1 ~
(/) (/)
LU z 0.2 ~ L.) ...... :::c !--
L.!..I '--::: _ .. X C
0.1
0 0
4 8 12 OXIDATION TIME (H)
16
Fig. 4.2. Oxide thickness as a function of oxida-~;~~t;:e15~5o~~andard dry oxides
47
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l.2. hours
700°c 8oo 0 c 900°c
1000°c
.8. hours
100°c 8oo 0 c 900°c
1000°c
H. hours
700°c 8oo 0 c 900°c
1000°c
.2. hours
700°c 8oo 0 c 900°c
1000°c
l hour
100°c 8oo 0 c 900°c
1000°c
TABLE III
Oxide Thicknesses(~) of the c2H3c1 2F Oxidations
0 .011 vol% O .055 vol% Control k .2. fl 3. .Ql.2_E. .c. .2. fl 3.Q .2. E.
69 106 20·6 287 675 1247 868 2830 4587
2500 6000 j924o I
55 100 169 217 506 938 792 2118 ·dftb 2033 4329 66 1
41 63 99 138 282 5l9 512 1078 2166
1469 2813 4696
32 46 62 98 156 290
302 652 1264 885 1800 2965
28 35 49 61 97 164
195 359 695 566 1109 1845
48
0.11 vol% .C.,2.fl3.Cl,2.E.
276 1705 6251
[ 12410 I
217 1315
mlliJ 1
126 708
2732 I 5110 I
73 395
1689 I 4834 I
50 209
mlli
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V 0(:) u V
Q) 0 0 0 0
e 0
0 0 0 oO o
0 C:fJo ®o 0 0 Ott 0 0
0
0 G
0
0 oo 0
0 0
~
0
0 (\
Fig. 4.3. Pinholes in Si02 grown for 12 hours
at l000°c with 0.055 vol% c2H3Cl 2F
(lOOX magnification).
49
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,-... z ;1 ,..,,
V,) (/) w z ,.,, u ....... :c !--
LIJ ~ -X 0
0.3
o 100°c D 800°C b. goo0c 'v 1000°c
0.2 -
0.1
0 -w OXIDATION TIME CH)
Fig. 4.4. Oxide thickness as a function of oxidation time for 0% c2H3Cl2F.
50
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....... ~ ;::,. ,._,
U) (/') L!..J z ::::.:::: u ..... ::J': 1--
L!..J c.:::i ,_ .. ><'. 0
0.6
0 100°c D soo0c D. goo0c V J.000°C
0.4 .
.,../ p
0.2
OXIDATION TIME (H)
Fig. 4. 5. Oxide thickne:.s as a function of oxidation time for 0.011 vol% c2H3Cl2F.
51
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0.8
,-.. ~ ;::.l. ..... .,, 0.6
C/) (/) LL.I z ~ :...) ....... ::r.: i-- 0.4 LL.I r..:::i ....... >< 0
o 100°c D 800°C .L\ 9oo0c V l000°c /
/ /V
./ ·/
/
/ /
/ /
/
--·--· _____ -o--
8
OXIDATION TIME (H)
12
Fig. 1+.6. Oxide tbickness a8 a function of oxidation time for 0.055 vol% c2H3cl2F.
52
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,...... ~ ;:l. -(/) (/)
w z ~ u -:::c: f-
w Q ->< 0
1.2
0 7oo0c /
800°C /
D /
fl goo0c "/
/
V l000°C 1.0 /
/ /
/ 0,8 /
/ I
Iv 0.6 I
vi
0.4 - I I
I
0.2 1" I I
00 4 8 12
OXIDATION TIME (H)
Fig. 4.7. Oxide thickness as a function of oxida
tion time for 0.11 vol% c2H3Cl2F.
53
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(./) U) LLJ z ~ u .--. :.r: 1-
1..!.J i:::i ......
0 7 ,)
0.2 -
a 0.1
I I
TEMPERATURE (00
I I
I
I I
I I
I
0
Fig. 4.8. Oxide thickness as a function of oxidizing temperature for 1 hour oxidations.
54
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0.5
0 0% C2H3CL2F D 0.011% C2H3CL2F
I 0.4 .6.0.055% c2H3CL~F 'v O. 11% C2H3CL2 I
,...._ I ::E: ;:l. I - 0.3 C/)
C/) I w z: :::s::::: I u -::r: I I-
w 0.2 I :=l ....... >< 0
0.1
800 900 1000
TEMPERATURE c°C)
Fig. 4.9. Oxide thickness as a function of oxidizing temperature for 2 hour oxidations.
55
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C/) C/) w z ~ u ....... ::c 1-
w ~
0.6
0.5
0.4
0.3
oO% CzH3CL2F
a o. 011% c2 H3CL2F .A0,055% c2H3CL2F 'vO .11% c2H3CL2F
I I
I J
I I
L I
I I
x 0.2 0
0.1
800 900
TEMPERATURE (0C)
Fig. 4.10. Oxide thickness as a function of oxidizing temperature for 4 hour oxidations.
56
1000
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1.2·
O 0% C2H3CL2F o 0.011% C2H3CL2F
1.0 6. 0.055% C2H3CL2F I \J 0.11% c2H3CL2F I
I 0.8 ,...... I
::E: ;:1 I ... _,
U) U)
I I 1..1.J z / Y- 0.6 I /....) I ,_. :--.c I I- I w 0 - I >( 0.4 C:..."'.)
Fig. LL 11. Oxide thickness as a function of oxidizing temperature for 8 hour
oxidations.
57
1000
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....... :E: ;:I. ..._,
U) (/)
w z: ~ u -:::r.: I-L.1J ;=) .... _. >< 0
1.2 O 0% CzH3CL2F
a 0.011% CzH3CLzF I I::,,. 0.055% C2H3CLzF I
'v' 0.11% CzH3CLzF I 1.0 I
I / /1
0.8 I I I I -·
I I 0.6 · I
I
0.4
0.2
0 ~==:.:::::::-:::__...(:,__J ___ ___JL----J
700 800 900 TEMPERATURE c°C)
1000
Fig. 4.12 Oxide thickness as a function of
oxidizing temperature for 12 hour
oxidations.
58
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,.-., ::E: ;::l. ~
U) U) w z ~ u ....... ::::c I-
w ;:::::1 ....... >< 0
0.3 o 100°c
D 800°C .,,.'v /
~ goo0c / ,,,,,,. /
V l000°C /
0.2 / /
0.1
0 ~::0::=:t:::==i=~t=c:==:::::::1c===~=Q._J 0.04 0.08 0.12 0
CzH3CL2F CONCENTRATION C voL%)
Fig. 4.13. Oxide thickness as a function of c
2H
3c1
2F co~centration for 1 hour
oxidations.
59
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"""" :a: ;:I.
'-"
C/) C/) w z ~ u ...... :c f-
w Q ...... >< 0
0.5
0.4
0.3
0.2
0.1
0 0
V 0 700°c /
D 800°C /
~ goo0c /
/ V l000°C /
/ /
/
i--------D -------0-0.04 0.08 0.12
Fig. 4.14. Oxide thickness as a function of c2H3Cl2F concentration for 2 hour oxidations.
60
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,...._ ~ ;:l.
'-J
U) U) w :z: ~ L.) ...... ::c f--
w ~ ...... X 0
0.6 0 100°c ---.,..,....
800°C .,,.,
D / /
~ goo0c /
"v 1000°c 0.4
0.2
0 0 0.04 0.08
Fig. 4.15. Oxide thickness as a function of c
2H
3cl2F concentration for 4 hour
oxidations.
61
_.. ..J:1
0.12
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,..... ::E: ;:l. ..._,
C/) C/) w z ~ u -:I: I-
w A -X 0
1,2
0 100°c 0 soo0c
Cl. 900°c /
/
'1 l000°C /
/ 0.8 /
/ /'1
/ /
/ /
0.4 'V/
0 0 0.04 0.08
Fig. 4.16. Oxide thickness as a function of
c2H3c12F concentration for 8 hour
oxidations.
62
,,,Yl
0.12
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,......_ :E: ;:1. .._,,
Cl) Cl) w z ~ u ....... ::x: I-
w r.:::i ....... X D
0 700°c
800°C 'v
1,2 D / .,..,,,..
I.),. goo0c .,/
/
1000°c /
V ./ /
/~
/ 0.8 ,/
/ /
I
0.4
-----0 -------0
---0 0.04 0.08
Fig. 4.17. Oxide thickness as a function of
c2H3c12F concentration for 12 hour
oxidations.
63
0.12
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C2H3c1 2F concentration. Another way of observing this
same increase is shown in Fig. 4~18, where the c2H3c1 2F
oxides are compared with chlorinated as well as standard
dry oxides.
Linear and parabolic rate constants were determined
for the c2H3c1 2F oxidations, and are plotted in Figs.
4.19 and 4.20, respectively. Included in these graphs are
the comparable constants for HCl oxidations.
Figures 4.21 and 4.22 depict the variation of
oxide thickness with fluorine concentration for two
different fluorine sources. These two graphs serve as a
comparison of a liquid (C2H3c1 2F) and a gaseous (NF3)
fluorine source. Further NF 3 results can be seen in
Fig. 4.23, which can be used as a comparison with the
results of Morita, et al. (28). In this graph, the
oxide thickness is expressed in units of nanometers, and
the fluorine concentration as parts per million of NF 3,
so that a more direct comparison could be made with the
data of Morita, et al. (28).
Table IV contains partial results of SOLGAS for
oxidations at 900°c when 0.11 vol% c2H3c12F was added to
the ambient. The partial pressures of the various
species are proportional to their concentrations in the
system. Out of the 82 possible chemical species in the
F-Cl-H-C-0 system, only the data for the eighteen most
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,-.. ~ ;:1 .._,
U) U) w z ::::.::::: u ...... :::i:: I-
w A ...... >< 0
10,0
1.0
0.1
o 0.055 VOL% C2H3CL2F A 10 VOL% HCL
a 3 VOL% CL2 e CONTROL
10.0
OXIDATION TIME (H)
100.0
Fig. 4.18. Oxide thickness vs. oxidation time for the oxidation of lightly doped 8ilicon in various gas ambients at 1000 C.
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.,...... :c
......... ~ ;:I.
'--'
1--z <::;C 1--(/) z: 0 u L.t.J 1--.::J'.: 0::
0:::: .:::( UJ z ...... _1
HCL CONCENTRATION (VOL%)
1.0 2 4 6 8 10 12
--r- ~ ~,,..-- ..-
---~ ~ _.~ ..
~ c2H3CLzF A900°C
A HCL. , V l000°C
0.001 0 0.04 0.08 o.i2
CzH3CLzF CONCENTRATION (VOL%)
Fig. 4.19. Linear rate constant vs. volio c
2Hfl
2F and HCl in o2 for the oxi-
dation of light~y doped (100) silicon at 900 and 1000 C.
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,....._ :J:
.·-....__ N ~ ;::1. ...._,,
f-z <( I-U) z 0 c __ j
l.lJ I-<C e::::
l..> ,._., _J 0 P:.l .::'.C 0::: <C o_
HCL CONCENTRATION (voL%)
0 2 4 1.0
6 8 10 12
.6. C2H3C1_2F
& HCL ..-...-
---0 . .1 -
" 0.001
0 O.OLI 0.08 0,L2
C2H3CL2F CONCENTRATION (voL%)
Fig. 4.20. Parabolic rate constant vs. vol%
c2H3Cl2F and HCl in Oz for the oxi-
dation of light!y doped (100) silicon
at 900 and 1000 C.
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0.20
...--. ~ 0.15 ;:i .._,,, (/) (/) LL.J :;;::: ::::..:.:: u 0.10 -::c I-Ll.J i::::i ........
°' ><
co c:,
0
.6. C2H3CL2F
A NF3
0.02 0.04 0.06 0.08 0.10 -0.12
FLUORINE CONCENTRATION (voL%)
Fig. 4.21. Oxide thickness as a function of fluorine
con5entration for 2 hour oxidations at
900 C; comparison 9f liquid and gaseous
fluorine sources.
0.14
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0.50 'v
0.40 V
0.3J ,,...._
::E: ;:l. .._,,
C/) (/) lJJ z ~ u ,_. :c i- 0.20
0\ w \,!) p -><
0
0.1
0 0
Fig. 4.22.
V /
CzH3CL2F / /
NF3 /
/ /
/ V
0.02 0.04 0.06 0. 08 0.10 0.12 0.14
FLUORINE CONCENTRATION (voL%)
Oxide thickness as a function of fluorine concentration for·2 hour oxidations at 1000°c; comparison of liquid and gaseous fluorine sources.
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r.--. ~ z '-'
(/) (/) w z ~ u ....... ::c: I-
--.J w 0 A .......
>< 0
70
60 ~ goo0c D 100°c
50
40
30
20
10
0 0 100 200 300 400 500 600 700
NF3
CONCENTRATION· (PPM)·
Fig. 4.23. Oxide thickness as a function of NF 3 concentration for 30 minute
oxidations at 700 and 900°C.
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TABLE IV
SOLGAS Results of Oxidations at goo0
c with 0.11 vol% c
2H
3c1 2F Added to the Ambient
Chemical Species
02
HCl
HF
CO 2
H20
Cl 2
Cl
ClHO
HO
ClO
H0 2
0
H202
H2
03
Cl02
ClF
co
71
Partial Pressure ~ml
O. 99x10°
o .43xlo-2
o.22x10-2
o.22x10-2
o.28x10-4
o.22x10-4
o.12x10-4
0.74xl0-5
0.25xl0-5
o.66x10-6
0.14xl0-7
o,89x10-8
o.22x10-8
o.45xlo-9
o.11x10-9
0.82xl0-lO
0.5lxlO-lO
o.18x10-10
O. 98x10-11
o.11x10-20
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abundant species and the fluorine species are included in
the table. Partial SOLGAS results for the N-F-H-0
system are reported in Table V. These data correspond
to 0.11 vol% NF 3 additions, along with 0.005 vol% H, to
a dry oxidation ambient at 900°c. Again, only the
concentrations for the most abundant chemical species
are contained in the table. Finally, the results of the
F-H-0 system, due to 0.11 vol% additions of HF to an o2
ambient at 900°, are presented in Table VI.
The SIMS results are presented in Figs. 4.24 to
4.26. These three figures represent the three different
types of profiles that were observed for the fluorinated
oxides.
High frequency capacitance-voltage (C-V) curves
are depicted in Figs. 4.27 and 4.28. The behavior
observed here, along with quasistatic C-V and bias
temperature-stress (BTS) results, are discussed.
72.
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TABLE V
SOLGAS Results of Oxidations at 900°c with 0.11 vol% NF 3 Added to the Ambient, in the
Presence of 0.005 vol% H2
Chemical Species Partial Pressure C atm)
02 0.99xl0°
F 0.5lx10-2
ONF 0.76x10-3
N2 o.11x10-3
F2 o.29x10-3
HF o.99x10-4
F02 0.16xl0-4
NO o.11x10-4
N02F o.12x10-5
FO o.11x10-5
N0 2 0.59x10-6
F2o o.22x10-7
0 0. 89x%10-8
NF 3 0.50xl0-8
F2N 0.29xlo-9
03 o.11x10-9
N2o 0.2lxlO-lO
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TABLE VI
SOLGAS Results of Oxidations, at goo 0 c with 0.11 vol% HF Added to the Ambient
Chemical Species Partial Pressure (atm)
02 0.99xl0°
HF O .22x10-2
0 o.9ox10-8
F 0.50xlo-8
HO O. 48x10-8
03 o.11x10-9
H20 O .11x10-9
H02 0. 28x10-lO
F02 0.15xlo-10
FHO 0. 22xl0-11
FO o.11x10-11
H202 0.84xlo-14
H o.31x10-14
H2 o.17x10-14
F2 O. 28x10-l5
F2o O. 21x10-l9
7~
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z C) -~ 0::: 1-z w u z 0 u
lx.1019
lxlo17
L------0 0.1 0.2
DEPTH ( µM)
Fig. lf.24. Fluorine and oxygen profiles gf an oxide grown fo·r 1 hour at 900 C with 0.011 vol% c2H3cl2F in o2
(Sample C from Table I) .
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lxlo23 .,.--~...--~....--~....--,~--.--~-r-~--.-~---::i
S102\ SI 1x1022 ,
11x1021' /
~
! lxl020 r;J ..._, !..,....,-------I-
I 1x1019 ! ~ ~ 8 1x1018 1
lxlo17 i 0 0.1 0.2
DEPTH (µM)
160+
49srF+
0.3
Fig. 4.25. Fluorine and oxygen profiles of0 an oxide grown for 2 hours at 1000 C with 0.011 vol% c2H3Cl2F in Oz (Sample F from Table I).
76
I
' J 3 ..J
~ d
~ ~
~ =-1
~
j 3
1
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.-~ ~ u
.......... (/)
~ 0 I-cl: '~ z 0 ...... I-c:r.: 0::: 1--z L.i..J u z 0 u
lxlo23 I
S102 l S1
lxlo22 I
lxlo21
lxl028
lxlo19 160+
1x1018 49s1F+
lx1017L---------~-'-----~~....__~------
O 0.1 0.2 0.3
DEPTH (11M)
Fig. lt.26. Fluorine and oxygen profiles
of an oxide grown for 2 hours
at 900°c with 0.011 vol% NF 3
in o2 (Sample I from Table I).
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-..1 co
500 +posjtjva to nGgotivG bios
to pos i ti vg bias
400 ~
Lt-a. \ ~
\
w 300 \
\ u \ z < t I-
\ ~ 200 u < i a.. < \
u 100
$.¢-<) ~ e,..~ ~o--0
0 -12 -10 -8 -6 -4 -2 0 2 4 6
Fig. 4. 27.
GATE BIAS CV)
High frequency C-V curve of a "non-ideal" fluorinated
oxide (Sample #7 from Table II).
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---1 \.0
800 •posjtjvg to negctjvg bjcs
011-ioot i V'1 to positJvg bias ""'
"600 lJ_ a.
'"" lJ.J u z 400 < r-..... u \ < a. ' < 200 u l
0 ..._ ......... __ ....._ .............. _.___,_ __ ....__--+-+---+--....---t
-8 -6 -4 -2 0 2 4 6
GATE BIAS CV)
Fig. 4.28. High frequency C-V curve of an "ideal" fluorinated
oxide (Sample #5 from Table II).
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5. 0. DISCUSSION
5.1 Results of Standard Dry Oxidations
The initial standard dry oxidations were performed
as a type of system calibration, so that the subsequent
fluorinated oxides could be compared with published data
on chlorinated oxides. Figures 4.1 and 4.2 were com
pared with similar graphs of other researchers (12,29),
and the oxide thicknesses were found to lie within the
range expected. Using the Deal~Grove (12) method, the
linear and parabolic rate constants at 900°c were calcu
lated to be 0.0191 µm/h and 0.0015 µm 2/h, respectively,
which are very similar to van der Meulen and Cahill's
(30) values of 0.0132 µm/h and 0.0024 µm2/h. Similar
ly, at 1000°c, the values in this study of 0.0559 µm/h
and 0.0095 µm 2/h compare reasonably well with their
published values of 0.0483 µm/h and 0.0101 µm2/h.
During oxidations with c2H3c1 2F, two simultaneous
processes are believed to be taking place, namely growth
and etching. At the lower temperatures and with smaller
amounts of additive, the fluorine-containing compounds
act to increase the oxidation rate, However, at the
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other extreme -- larger concentrations of c2H3c1 2F at
higher temperatures -- these same fluorine-containing
compounds lead to etching of the oxide, resulting in
pinholes (as can be seen in Fig. 4.3).
The remarkable aspect of the c2H3c12F oxidations is
that such a significant enhancement of oxidation rate
was achieved with such small additive concentrations of
the fluorine compound. The data in Table III illustrate
the magnitude of the enhancement.
The variation of oxide thickness enhancement with
time can be observed in Figs. 4.4 to 4.7. In all cases,
the relative enhancement observed upon adding 0.11 vol%
c2H3c1 2F to any dry ambient increased with oxidation
time. For one hour oxidations, the average enhancement
observed over the entire temperature range was 260%.
This increased to 340% for four hour oxidations, and
finally seemed to level off at 450% for twelve hour
oxidations.
In going from Oto 0.11 vol% c2H3c12F, the enhance
ment at 700°c varied between 75% and 300% (depending on
oxidation time). At 8oo 0 c, the relative increase ranged
from 250% to 500%. Temperatures of goo 0 c brought about
enhancements of 400% to 600%. Finally, at 1000°c, the
increase ranged from 300% to 450%. Generally, it can be
said that the relative enhancement increased between 700
and goo 0 c, and then decreased in going from 900 to
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1000°c. In all instances, the decrease at 1000°c was
accompanied by a deterioration of the physical integrity
of the oxide. Those data points which represent poor
quality oxides, in which the thickness measurements may
not have been very accurate, are connected by dotted
lines in the graphs.
The dependence of oxide thickness enhancement on
c2H3c1 2F concentration is illustrated by Figs. 4.13 to
4.17. Generally, additions of 0.055 vol% c2H3c1 2F to
the dry ambients caused increases of approximately 200%
for the shorter oxidation times, and 315% for the longer
oxidation times. Similarly, additions of 0.11 vol%
c2H3c1 2F yielded increases of 250% to 450%. In all
instances, the amount of enhancement observed increased
with c2H3c1 2F concentration.
The magnitude of the enhancement can also be judged
from Fig. 4.18. The c2H3c1 2F data is plotted for
1000°c, so that comparable data for 10 vol% HCl (19) and
3 vol% c1 2 (31) could be included in the graph. It can
be seen that under comparable conditions, the oxidation
rate is increased with approximately two orders of mag
nitude lower vol% c2H3c1 2F over those grown in chlorine
ambients.
The c2H3c1 2F oxidation data was also analyzed in
terms of the linear and parabolic rate constants.
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Figures 4.19 and 4.20 are plots of these constants vs.
vol% c2H3c12F. The data are plotted at 900 and 1000°c
so that comparable data using HCl as an additive (19)
could be included in the graph. It is clearly seen that
the constants for c2H3c1 2F oxidations are larger than
those for the HCl oxidations which have two orders of
magnitude higher concentrations. The behavior of the
two rate constants was similar at 700 and 8oo 0 c; how
ever, there was no comparable chlorine data available at
these lower temperatures.
Both of the rate constants increase as more c2H3c1 2F
is added to the system (at least over the range investi
gated). This suggests that, even at the very low con
centrations of c2H3c1 2F, there is a direct influence on
the several factors that determine the linear and para
bolic rate constants. The definitions of the linear
(B/A) and parabolic (B) rate constants written below for
convenience:
B/A = C*/[N1(1/k + 1/h)]
B: 2 Deff C*/N1
where C* is the equilibrium concentration of oxidant in
the oxide, N1 is the number of oxidant molecules incor
porated into a unit volume of the oxide layer, k is the
surface reaction rate coefficient, his the gas phase
transport coefficient, and Deff is the effective diffu
sion coefficient of the oxidant through the oxide.
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Figures 4.19 and 4.20 show that even at the lowest
additive concentration of 0.011% c2H3c12F there is an
immediate substantial increase in both the linear and
parabolic rate constants. One's first impulse would be
to attribute the increase to the factor that is common
to both of the rate constants, i.e. C*/N1• However,
such small concentrations of the c2H3c1 2F are added that
it is unlikely that either the solubility (equilibrium
concentration) or the number of oxidant molecules incor
porated in the oxide is changing significantly. Also,
in comparing Figs. 4.19 and 4.20, it should be noted
that 0.011 vol% c2H3c12F additions result in different
degrees of enhancement for the two rate constants.
Therefore, it is concluded that Deff and k are also
changing in this concentration range.
One interesting and important aspect to note is the
difference in behavior of the linear rate constant as
the HCl or c2H3c1 2F concentration is increased. The
linear rate constant increased with c2H3c12F concentra
tion over the entire range investigated. However, it
increased only up to 1% HCl additions, and then levelled
off. The levelling off may be caused by saturation,
i.e. 1% HCl in the system may be enough to saturate all
of the active surface sites, so that further additions
may no longer enhance the rate. This idea should be
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tested for the case of chlorine oxidations.
Obviously, fluorine-bearing compounds play a major
role in the growth kinetics through their influence on
both Band BIA. A possible explanation is that fluorine
acts as a hydroxyl anion in the silicon dioxide struc
ture (32). In other words, as the oxidation proceeds,
Si-F bonds are formed directly at the interface, re
sulting in some type of Si-0-F complex. And since Si-F
bonds are formed more readily than Si-0 bonds (based on
electronegati vi ty differences), the reactivity at the
silicon-oxide interface is enhanced, leading to an in
crease in the linear rate constant. The formation of
Si-F bonds also results in a larger number of non
bridging bonds, which leads to a looser structure.
This looser structure leads to enhanced diffusivity of
the oxidizing species, resulting in an increase in the
parabolic rate constant.
5.3 Results of Oxidations Hith NF 3
Results of oxidations performed with a gaseous
fluorine source--nitrogen trifluoride (NF 3)--indicate
that the oxidation rate is also increased with this
additive, but to a lesser extent than with the c2H3c1 2F.
This claim is supported graphically by Figs. 4.21 and
4.22. In comparing these two graphs, it is obvious that
there is a difference in behavior between the oxidations
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performed at 900 and 1000°c. It appears that the degree
of enhancement is greater at the lower temperature of
900°, i.e., an increase of 320% is observed in going
from 0% addition to 0.044 vol% NF 3. At 1000°c, this
same concentration range yields enhancements of only
95%. Similarly, when comparing Fig. 4.23 to the pre
vious figures, it can be seen that at 700°c, increases
of 600% are observed over this same concentration range.
Generally, it appears that the degree of enhancement is
decreasing with temperature between 700 and 1000°c. It
is possible that, at the higher temperature, the "etch
ing" reaction has taken over, thereby decreasing the
resultant growth rate without causing any observable
pinholes. However, much more data need to be collected
on NF 3 oxidations before such a claim can be substan
tiated.
Another interesting point to note in the NF 3 oxida
tions is that in all cases the oxides appeared optically
to be good quality oxides. In observing Fig. 4.22, it
can be seen that the same fluorine concentration yielded
"good quality" NF 3 oxides in one case and "poor quality"
c 2H3c1 2F oxides in the other case. However, this phe
nomenon may also be a function of oxide thickness, i.e.,
the NF 3 oxides were not grown to the same thickness as
the c2H3 Cl 2F oxides.
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The data in Fig. 4.23 were compiled so that a
direct comparison with the results of other researchers
(28) could be made. This graph can be directly compared
and contrasted with Fig. 2.15. It is very apparent that
the general shape of the curves in this study does not
parallel the previously reported data. In both cases,
fluorine was introduced into the dry ambient in the form
of NF3, over the same concentration range. However, the
experimental conditions were slightly different. In the
system of Mori-ta, et al. (28), a cold-walled quartz tube
reactor was employed, and the wafers were placed onto a
silicon susceptor which was heated by irradiation. In
this study, a standard hot-walled oxidation system was
used. As a result, more hot surfaces were competing for
the reactive gas, and the NF 3 concentration that the
wafers were exposed to was thereby probably lowered.
The effectively lower NF 3 concentration may also help to
explain why the oxides did not appear to have pinholes.
In examining Fig. 4.23 more closely, it can be seen
that at 700°c, the oxide thickness increased smoothly
with NF 3 concentration, at least over the tern perature
range investigated. This behavior was similar to that
observed for the liquid fluorine source. However, at
900°c, there appears to be some discontinuity in the
form of a slight peak at approximately 30 ppm. This is
where the peak is expected to lie, based on the findings
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of Morita, et al. (28). Had the data points of this
study been taken at smaller concentration intervals, the
peak may have been more pronounced. However, the rest
of the curve does not follow the general shape of
Morita, et al., i.e., the data show the oxide thickness
increasing smoothly with NF 3 concentration whereas
Morita's data depict a decrease. The decrease was
attributed to etching of the oxide by fluorine radicals.
Again, this discrepancy can be explained by the effec
tively lower NF 3 concentration--in this study, the data
points were still within the growth regime that occurs
at lower fluorine concentrations, whereas their data had
already passed into the higher concentration etching
regime.
5.4 SOLGAS
When c2H3c1 2F is allowed to react with oxygen, and
then this mixture is heated to the elevated oxidation
temperatures, the following reaction is believed to be
taking place:
C2H3Cl 2F + X0 2 t HF+ 2HC1 + 2C0 2 + Y0 2 (4.1)
The oxygen concentration in this reaction equation is
left as X and Y moles of o2 , because it is so much
higher (by two to three orders of magnitude) than that
of the dichlorofluoroethane that it seems inappropriate
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to attach exact numbers to these molar concentrations.
The reaction product HCl then further dissociates to
form c1 2 and H20, as it does in chlorine oxidations.
From this equation and from the results in Table IV,
it can be seen that many of the dissociation products of
c2H3c12F are the same as those formed when trichloro
ethylene (C 2Hc1 3) or trichloroethane (C 2H3c13) is
heated. However, there are additional fluorine-bearing
species in the c2H3c1 2F system that are not present in
the c2Hc1 3 or c2H3c1 3 systems. Table IV indicates that
the only fluorine-bearing species present in non-negli
gible amounts in the furnace is hydrogen fluoride (HF).
Although the table only contains the results of a speci
fic oxidation, namely O.ll vol% c2H3c1 2F additions at
900°c, the relative amounts of the chemical species were
consistent throughout the entire temperature and concen
tration ranges investigated. Therefore, it is postu
lated that HF is the active species that results in such
marked increases in the oxidation rate.
When NF 3 was added to the oxidation ambient, it was
initially expected that there would be no enhancement of
rate, since there was no hydrogen added to the system
which would form HF. In reality, however, there was
some hydrogen present in the furnace due to impurities
in the o2 and to in-diffusion of air through the quartz
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tube. Therefore the SOLGAS program was run for an input
of NF 3, H2, and o2, with the H2 varying between 0.000005
and 0.005 moles. This range was believed to be repre
sentative of the actual conditions existing within the
furnace. The partial results are contained in Table V.
It can be seen that very small hydrogen additions to the
NF 3-o2 ambient resulted in appreciable amounts of HF.
The HF concentration was found to increase as the hydro
gen additions to the ambient increased.
SOLGAS was also run for HF additions to an o2
ambient, and the results are contained in Table VI. No
oxidations were performed with HF; this was merely done
to see if HF could be used as a possible fluorine
source. In comparing the HF results to the c2H3c1 2F and
NF 3 results, it can be noted that with HF additions, the
HF concentration is the same as for c2H3c1 2F additions
but higher than for NF 3 additions. However, the H20
concentration is approximately one order of magnitude
lower for HF additions than for c2H3c1 2F additions. The
HF-02 system also does not contain any of the chlorine
bearing species that act to enhance the oxidation rate.
Therefore, it would be expected that the enhancement
achieved with HF additions would lie somewhere between
that achieved with NF 3 and c2H3c12F additions. This
will be the focus of future research.
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5.5. SIMS Analysis
Based on a SIMS analysis, it appears that the wafers
can be grouped into three categories. The first of
these categories contains samples A, B, C and D from
Table I, and has a typical fluorine profile as is shown
in Fig. 4.24. These oxides all had oxide thicknesses
less than 600 R which were grown using c2H3c1 2F as the
fluorine source. These oxides are characterized by a
fluorine peak just above the silicon-oxide interface,
and a relatively high level of fluorine throughout the
oxide layer.
Figure 4.25 is representative of the second group,
which contains samples E and F of Table I. These oxides
were greater than 1000 R thick, and were grown in a
c2H3c1 2F-02 ambient. The SIMS profiles show a low fluo
rine concentration at the surface, which increases
sharply approximately halfway into the oxide layer. The
fluorine concentration continues increasing up to about
3/4 of the oxide thickness, and it remains at this level
(which is the same as the peak level of the oxides in
the first category) through to the silicon-oxide inter
face.
The last group, which contains samples G, H, and I
of Table I, has fluorine profiles that are typified by
Fig. 4.26. These samples are unique in that they were
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all grown using NF 3 as the fluorine source. They all
exhibit a higher level of fluorine throughout the oxide
layer, with a very slight peak approximately 1/4 of the
way in from the outer surface.
The difference in fluorine profiles of the NF 3 and
c2H3c1 2F oxides can be explained on the basis of a model
derived for chlorinated oxides (33). According to this
model, the initial chlorine profile throughout the
oxides is relatively flat. As the oxidation in HCl or
c2tt3c1 3 proceeds, the bound chlorine is replaced by
hydroxyl groups, which come from the water that is
generated in the system. This results in a decrease in
the chlorine profile. Oxygen can also replace the
bonded Cl, but at a much slower rate than water. It is
expected that this same replacement mechanism occurs in
fluorinated oxides. As a result, oxidations with
c2H3c12F, which contain large amounts of generated water
in the system, should yield peaked fluorine profiles,
whereas oxidations in NF 3, in which very little water is
generated, should result in flat profiles.
In all of the samples, there was a definite drop in
the fluorine concentration at the silicon-oxide inter
face, as is observed for chlorine profiles in chlori
nated oxides. Also, the fairly constant fluorine level
observed throughout all of the samples indicates that
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the fluorine is not mobile in the oxide, but is tied
into the structure.
5.6 Electrical Characterization
High frequency capacitance-voltage (C-V)
measurements indicated that the fluorinated oxides
exhibited electrical properties comparable to those of
standard dry thermal oxides. Only one sample was found
to deviate substantially from an ideal C-V curve (34),
i.e., sample #7 from Table II. An example of this
deviation is shown in Fig. 4.27. This oxide was grown
at 700°c, and did not have a post-oxidation anneal (POA)
or a post-metallization anneal (PMA). It is not
surprising that this sample exhibited unusual C-V
characteristics, as low-temperature oxides have
generally been found to have inferior electrical
properties (35). Also, the looser structure of the
fluorinated oxides (28) may have caused a deterioration
of electrical properties.
Generally, all of the other samples tested
produced "ideal" C-V curves, as is illustrated by Fig.
4.28. This curve characterizes sample #5 from Table II,
which had an oxide grown at 900° and did not have a
POA. It also depicts one characteristic that was
typical of all of the samples, i.e., the dip in curve at
approximately -1 V. This dip was only present in the
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forward sweep (positive to negative voltage), not in the
reverse sweep, and appeared to anneal out with biasing.
The dip has been previously observed by other
researchers (36), and is explained as being caused by
the formation of a depletion edge region around the
aluminum dot.
A comparison of high frequency and quasistatic C-V
measurements determined that the higher temperature
oxides exhibited lower interface state densities. The
low temperature, non-annealed oxides (samples #3 and 117
from Table II) were too leaky to generate quasistatic C
V curves. Annealing for 100 minutes in o2 at 700°c
(samples #4 and 8) lead to lower interface state densi
ties, but they still ranged from 5xloll to 1x1012 ev-1
cm-2. It was determined that this low temperature an
neal did not have much of an effect on the electrical
properties, as only 11 i of thermal oxide were grown
during the anneal, and this 11 i was apparently not
enough to change the properties of the interface.
It was also found that fluorinated oxides could be
grown with electrical properties at least as good as
those of standard dry thermal oxides, as there was
virtually no difference between samples #2 (control) and
#10. Samples #10, which had a high fluorine content,
had a POA in oxygen at goo 0 c, so that the actual MOS
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structure was metal-fluorinated oxide-thermal oxide
silicon. Therefore, the calculated interface state
densities actually reflected a thermal oxide-silicon
interface. One thing that this experiment did prove is
that at the same oxidation time and temperature, thicker
oxides could be grown with the addition of fluorine to
the ambient, and that these thicker oxides were electri
cally comparable to the thinner standard dry oxides.
The bias-temperature-stress (BTS) experiments
~bowed all of the oxides to have sodium ion concentra
tions between 1011 and 1012 cm-2. These numbers appear
to be on the high side, but are actually reasonable
considering the fact that no special care was taken to
ensure that the wafers were cleaned of ionic contami
nants before oxidation. Wafers 117, 8, 9, and 10, which
were all oxidized with the higher fluorine concentra
tion, exhibited the highest Na+ concentration, approxi
mately one order of magnitude larger than the control
samples. Coincidentally, these four wafers were metal
lized together; therefore, no claim can be made as to
whether the elevated Na+ concentration was brought about
by the fluorine concentration or by the metallization
procedure.
Chlorine has been found to passivate sodium (4,5);
however, the oxidation must take place at a high enough
temperature with a high enough chlorine concentration
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for a long enough period of time in order for the Na+ to
be trapped. Well-passivating oxides are 1100-1200 i
thick, grown at 1050°c or higher in an ambient contain
ing at least 8 vol% HCl. Therefore, it is not sur
prising that these oxides, which were less than 900 i
thick, grown at 700 to goo 0 c in an ambient containing up
to 0.044 vol% NF 3, were not effective in trapping and
neutralizing the Na+.
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6.0 SUMMARY
It was experimentally found that small additions of
a fluorine-bearing species to a dry oxidation ambient
increased the oxidation rate considerably. The enhance
ment was observed with two different fluorine compounds,
namely liquid dichlorofluoroethane (C2H3c1 2F) and gaseous
nitrogen trifluoride (NF 3). The accelerated growth rate
was evident in both the linear and parabolic rate con
stants, thereby indicating that both the diffusion of
oxidant through the existing oxide and the reaction at
the silicon-oxide interface were increased by the
fluorine additive. The relative enhancement increased
with oxidation time, oxidation temperature, and fluorine
concentration, but only up to a certain point. At that
point, the etching process which is competing with the
growth process took over, resulting in "poor quality"
oxides in which pinholes may have been observable.
The amount of c2H3c12F added to the ambient varied
up to 0.11 volume percent, which is approximately two
orders of magnitude lower concentration than its
comparable chlorinated oxides. In fact, these
fluorinated oxides were found to increase the rate
substantially more than even 10 vol% HCl or 3 voll c1 2•
A computer program (SOLGAS) was used for a
thermodynamic analysis of the oxidation ambient •. It was
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determined that at elevated oxidation temperatures,
C2H 3c1 2F dissociates to form many of the reaction
·products formed upon the dissociation of c2H3c1 3 or
c2Hc1 3• Some of these reaction products, such as H2o or
Cl 2, are known to accelerate the growth rate. However,
they were present in the furnace in such small quantities
that they could not have created such a large increase.
But there was one compound that was unique to the
C2H 3c1 2F-D2 system, which was present in relatively large
amounts -- hydrogen fluoride (HF). Therefore, it was
postulated that HF was the active species that resulted
in such marked increases.
Additions of NF 3 to the dry ambient also caused the
oxidation rate to increase, although to a lesser degree
than the c2H3c1 2F. This was expected because there were
no chlorine-bearing species present and also because the
HF concentration was lower. The only hydrogen present
with which to form HF came from impurities in the oxygen
or from in-diffusion of moisture through the quartz tube.
Fluorine concentration profiles were determined by
Secondary Ion Mass Spectrometry (SIMS). It was found
that the c2H3c1 2F oxides yielded fluorine profiles which
were peaked at the silicon-oxide interface. The peaks
were sharp for the thinner oxides, and broader for the
thicker oxides. For the NF 3 oxides, the fluorine profile
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was constant throughout the oxide layer. This behavior
was explained based on a model previously derived for
chlorinated oxides. Apparently, the bonded fluorine
ions can be replaced by hydroxyl groups or oxygen, with
the replacement rate of OH being much greater. In a
system where water is readily available, the fluorine
concentration is reduced as it is replaced by OH.
The oxides were also tested for their insulating
properties, and it was found that even with the maximum
fluorine concentration, they could be electrically as
reliable as standard dry oxides.
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7.0 RECOMMENDATIONS FOR FUTURE RESEARCH
First of all, I think it is very important to
substantiate the claim that HF is the chemical species
enhancing the growth kinetics. This can be done by
oxidizing with various fluorine-bearing compounds, and
analyzing the growth rates in conjunction with SOLGAS.
The most obvious fluorine additive to investigate is HF,
in liquid or gaseous form, although other, less toxic
additives could also be used.
Another critical set of experiments should include
"extremely dry" oxidations, performed with ultra-high
purity oxygen in a double-walled quartz tube. These
conditions would prevent any hydrogen from unintention
ally getting into the system, and would further serve to
check if HF was the active species.
Many more oxidations using NF 3 as the fluorine
source need to be carried out, in an attempt to further
explain the discrepancy between my results and those of
other researchers.
Etch rate experiments should also be performed, to
evaluate how the etch rate of the oxides depends on
fluorine concentration and post-oxidation treatment.
A much more extensive study of the electrical
characteristics of the fluorinated oxides needs to be
undertaken. Many more oxides need to be grown, and MOS
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capacitors fabricated, so that general effects of
temperature, post-oxidation treatment, and fluorine
concentration can be determined.
Finally, a model for the growth kinetics, along
with accompanying equations, needs to be derived, or the
existing linear-parabolic model needs to be modified to
include fluorinated oxidations.
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presented at the 167th Meeting of the Electrochemical Society, Toronto, Ontario, May 1985.
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VITA
Christine Helen Wolowodiuk was born to Catherine
and Walter Wolowodiuk on December 12, 1961, in New York
City. Raised and educated in New Providence, New
Jersey, she entered Rutgers University-College of
Engineering in the fall of 1979. She graduated with
high honors in 1983, with a B.S. in Ceramic Engineering.
That same year, she entered into the graduate
program at Lehigh University, in the Metallurgy and
Materials Engineering Department. While studying there,
she has been supported by a teaching assistantship,
Sherman Fairchild Fellowship, and Lehigh University's
Distinguished Scholar Fellowship.
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