Suggestions for Lab report “Topic 11”

1. Purpose

2. Procedure ( Brief description in point form)

3. Data collection and Processing

a) Data collection
Trial Water volume ± uncert cm3 Balance reading ± uncert
Empty cup 0.00 3.045
1 + 10.00 13.128
2 + 10.00 23.205
3 + 10.00 …

Table 1: Successive masses of cup and 10.00cm3 of water


b) Data processing
• Show formulas and calculation ( Sample calculation if repetitive)
• Tabulate results
V water ± uncert cm3 Mass of 10.00 cm3 portion ± uncertg Deviation │m – X │ ± uncertg
10.00 … …
10.00


Table 2: ……………

• Calculate the mean deviation “d”

• Test the results for precision

• Test the results for possible rejection of a measurement. If any of the measurements should be rejected, prepare a new table of results excluding the rejected data.

Correction extra exercises

#1
0.28 b) 6.0x10-2 c) 2.2 d) 11

#2
(M2-M1)/2 = 50.3g
50.3 ± 0.2 g
δ (M2-M1)/2 = ±0.3/2 = 0.15 = 0.2


#3
X = 14.91% and d= 0.01 ; d/x = 6.7x10-4 < 2x10-3 Measurements are precise
d/xtrue = 6.3x10-4 < 2x10-3 Measurements are accurate

Topic 11

Uncertainty and Errors in measurements "Raw Data"
a) Uncertainty
 No Science can offer absolute certainty; an uncertainty range applies to any experimental value. (Measured value ± Uncertainty Unit)
 Uncertainty in a measurement depends on the precision of the measuring device:
 In graduated instruments, the uncertainty is ± of the smallest division
(Unless given a different value)
 In digital instruments the uncertainty is ± the smallest scale division

b) Significant figures
 Certain and uncertain digits are called significant figures. When making a measurement it is important to record the result to the appropriate number of significant figures. Ex: for a burette 25.00 ml and not 25 ml.
 Rules for counting significant figures

 Refer to Book page 216/423

c) Experimental Errors
 Random errors:
- Are caused by the way we read measuring instruments or by changes in surrounding conditions
- Can be reduced by repeating measurements
 Systematic errors are of 3 types:
- Instrument errors that can be corrected by calibration
Ex: Temperature measured by a thermometer that has an imperfection.
- Method errors, these are difficult to detect and control
Ex: Side reactions during a Chemical change.
- Personal errors that can be minimized by self-discipline.
Ex: Reading a burette or detecting a color change.
 Accuracy and Precision
 Precision:
- Degree of agreement among several measurements of the same quantity.
- Precise measurements have small random errors.
 Accuracy:
- Agreement of a particular value with the true value.
- Accurate measurements have small systematic errors.
 Refer to Book page 217and 218/424 and 425

Uncertainties in calculated results "Processed Data"
a) Rules for significant figures in Mathematical operations
 In a multiplication or a division; the number of significant figures in the result is the same as the number of significant figures in the least precise measurement.
 In an addition or a subtraction; the number of decimal places in the result is the same as the number of decimal places in the least precise measurement.

PS: For rounding look at the digit to be dropped; if it is ≥ 5 then round up the digit before it, if it is < 5 drop it without changing the digit before it.


Graphical Techniques
 Dependant variable and independent variable
 Best fit line (correction of random errors)
 Displaced line (As a result of systematic errors)
 Extrapolating a straight line
 Gradient of a straight line:
 Gradient of a curve at a point = Gradient of tangent at this point
 Plotting data to produce a straight line
Example: Ideal gas equation PV = nRT; Plot P (vs) 1/V
 Using Data logger to plot graphs
Example: Titration curve of an acid
 Refer to Book pages 222---225/429---432

Orientation session IB-1

IB-I Orientation Session

• Attendance and SL/HL

• Establish a group e-mail

• Course Description
Objectives
It is the intention of all the Diploma Programme experimental science courses that students achieve the following objectives.
1. Demonstrate an understanding of:
a) scientific facts and concepts
b) scientific methods and techniques
c) scientific terminology
d) methods of presenting scientific information.
2. Apply and use:
a) scientific facts and concepts
b) scientific methods and techniques
c) scientific terminology to communicate effectively
d) appropriate methods to present scientific information.
3. Construct, analyze and evaluate:
a) hypotheses, research questions and predictions
b) scientific methods and techniques
c) scientific explanations.
4. Demonstrate the personal skills of cooperation, perseverance and responsibility appropriate for effective scientific investigation and problem solving.
5. Demonstrate the manipulative skills necessary to carry out scientific investigations with precision and safety.

Outline and options
- Core material: 11 topics and 2 options
- Additional HL material for 9 of the previous topics and the 2 options

External Assessment
• Papers 1, 2 and 3
(Book page 502 and 503/866 and 867)
• Quality of answers (Key words, extended answers….)
• Action verbs or command terms
(Refer to given H-out)
• Extended Essay: in any topic and contributes with TOK to overall diploma score. (Book page 497---501/ 861---865)

Internal Assessment
• 24% of final assessment
• Lab activities according to IA criteria
(Book page 492---496/856---860)
• Group 4 project : interdisciplinary and stressing Scientific investigation


Material Needed
• Textbook
"PEARSON/HEINEMANN BACCALAUREATE"
CHEMISTRY SL/HL
• Notebook
• Folders For lab and H.out

Work Habits
• Concentration in class
• Group work: Division of tasks, involving all members.
• Daily work: being organized and setting schedules to meet deadlines.