Chemical Analysis can be of two types: Chapter 11- Measurement and Data Processing: - : Substances are classified on the basis of their or properties, such as - : The amount of the sample determined in with appropriate units. Uncertainty in Measurement: Numerical data can be divided into types: - Data involving exact numbers and data involving inexact numbers. When we carry out experiments, there will always be some uncertainty associated with the measured data, that data has numbers. Such uncertainty is most likely due to the instruments used in the lab or human error. Difference between accuracy and precision: Accuracy and Precision: Accuracy: An accurate measurement is a measurement. It Is the closeness between the result of a measurement and a value of the measured. Precision: A precise measurement is a measurement. (Measurement that appears over and over again). The more precise a measurement, it has. Precision is the closeness between
test results. In other words: Accuracy means getting a result that is. Precision means getting a every time you try. Think of shooting at a target: Being accurate means you hit the bull's eye. Being precise means hitting the same spot on the target every time. Examples: Accuracy is about arrows being on the target (Correct spot). Precision is about arrows being. (Almost same spot over and over again). NOTE: however; that even precise measurements can be inaccurate, if the equipment is faulty or.
Significant Figures: It refers to the number of digits reflecting. The greater the number of significant numbers, the greater is the certainty of a measurement. Rules: 1. All digits are significant 1.22 9456 8 8.2 2. are significant eg. 2005 3.06 900.0035 3. All zero s to the of the are NOT significant. (called leading zero s) eg. 0.005 0.00000506 4) Zero s to the of the (trailing zero s) ARE significant IF the is SHOWN. Eg. 4.00 300.000 4000. 0.0560 0.0050600 5) Zero s to the RIGHT of the last non-zero digit (trailing zero s) ARE NOT significant if there is an UNDERSTOOD decimal point. Eg. 600 9 800 000 10
50.0060 0.005600200 59 000 6) In Scientific Notation, the of the number is NOT significant. Eg. 3.45 x 2.000 x 4.903 x 0.05600 x 7) Find the number of significant digits in each of the following measurements: 0.002060 3.500 x 98 000 876 000. 200.003 3.2 x 5678.94 1.000000 0.0000012 0.002 x 8) Rules for Calculations with Significant Digits For Multiplication or Division: Answer is rounded to the OF SD S in the question. Perform the following calculation and round the answer to the correct number of significant digits: (5.6 x ) (3.651 x 10-7) (Calculator answer )
Final answer 9) Perform the following calculation and round the answer to the correct number of significant digits: (Calculator answer ) Final answer 10) Rules for Calculations with Significant Digits for Addition or Subtraction: Answer is rounded to the LOWEST # OF DECIMAL PLACES in the data. Do the following calculation and round the answer off correctly: 4456.9833 + 32.41 = Note : The exponent indicates how many tens are multiplied together. - The exponent can also be interpreted as an indication of how many places to move the decimal point in the decimal portion of the number. - A POSITIVE exponent requires the decimal point to be moved to the RIGHT. - A NEGATIVE exponent requires the decimal point to be moved to the LEFT. Examples: 1. Change the following numbers to decimal notation: a. 2.75 x 103 =
b. 5.143 x 10-2 = 2. Change the following numbers to scientific notation: a. 6547 = b. 0.00168 Hint: A number correctly expressed in scientific notation should always be written so that there is only one non-zero digit in front of the decimal point. 11) Significant figures associated with logs: Log of a number: Is the to which the can be raised to get that number. A log has 2 parts: The number of digits in the mantissa indicates the number of sig figs in the question. Ex: Since the number has SFs, so there should be SFs in the decimal part of the answer. Experimental Errors: Every measurement has associated with it. There are 2 types of experimental errors:
- - Systematic Errors: Flaw in the or with the. Systematic errors can be further classified into: - Instrumentation errors - Experimental methodology errors - Personal Errors Examples of systematic errors: - Systematic errors affect the of the results. Random Errors: Random errors occur due to. These errors can be reduced by. Random errors affect the of the results. Examples: Absolute and relative uncertainty: Absolute uncertainty: Is the uncertainty associated with the result from a measurement. Relative uncertainty: Is the ratio of the absolute uncertainty to the measured experimental result. Example: Absolute uncertainty =.02 cm3 The recorded volume at the end point= 12.25 cm3 Find the relative uncertainty.
Percentage Error: % Error= literature value-experimental value * 100 % Literature value Example: If the literature value of calcium carbonate is 178.1 kj and the experimental value was 172.0 kj. Find its percent error. 11.2 Graphical Techniques: Graphs are used for representing data. Independent Variable: Is plotted on the axis, is usually the Dependent Variable: Is plotted on the axis, is usually the 3 main features of graphs:
- The slope (m): Example: Find slope- - The intercept ( c) : Is the point where the line cuts the y-axis at x=0 The intercept can be found by two methods: a) Using extrapolation This method involves extending a line back to the y-axis to find the intercept. B) Using the equation of a line: y= mx+ c - A best-fit line: When you plot data from an experiment you may find not all points lie exactly on the line. It is a straight line that best represents the data on a scatter plot.
AP Molar mass (amu, g/mol) often inaccurately referred to as molecular weight (MW) Most common volume units in chemistry: Molar mass (g/mol or amu) often accurately referred to as relative molecular mass (M r ) Most common volume units in chemistry: IB ml cm 3 (this is the same as a ml) L dm 3 (this is the same as a L) Most common concentration units in chemistry: Most common concentration units in chemistry: Molarity (M): mol/l Molality (m): mol/kg mol dm -3 (notice that this is the same as M) g dm -3 (notice that this is NOT the same as m)