1. Fundamentals of Chemistryby
Dr Violeta Jevtovic
Assistant Professor of Inorganic Chemistry

2. Introduction Chemistry: Branch of Science which deals with study of
composition of matter, properties of matter, changes in
matter and law and principles under which these changes
occur. Or
Chemistry is the science of atoms, their structures, their
combination and their interactions.
Science: A process for understanding nature and its
changes to explain phenomena of the physical world.
Matter: Anything which as some mass and occupy
some space. Examples: water, Gold, NaCl, sugar, Air etc.

3. Chapter I: Chemical Foundations
Section 1: Units of Measurement
Making measurements is essential for all sciences.
In chemistry, we generally use measurements of mass, temperature, time, amount of substance (number of moles),…
If measurements were to be useful, a standard system of units has to be adopted.
There are 2 majors standard systems of units which are adopted around the world:

4. Units of Measurement (Cont’d)
The metric system: used by most of the world
The English system: used in the USA
In 1960, an international agreement set up a system of units called the International System (Le System International in French) or the SI system.
The SI system is based on the metric system and units are from the metric system.

5. The Fundamental SI Units
Note that the volume V, which is a very important physical quantity in chemistry, is not a fundamental unit but is derived from length. A cube that measures 1 m on each edge has a volume of 1 m3.
1 m*1m* 1m = 1 m3

6. The Fundamental SI Units (volume conversion)
There are 10 dm in 1m, then: (1m)3 = (10 dm)3 = 1000 dm3 1 dm3= 1 L 1000L = 1 m3 1L = 1000 cm3=1000 mL, therefore 1 mL=1cm3

7. Larger and Smaller Units
p = pico ( )
n = nano (one billionth)
µ = micro ( )
m = milli (0.001)
c= centi (0.01)
10o = 1
Larger units
h= hecto
k = kilo (1000)
M = mega ( ) 106
G = giga ( ) 109 7

8. Section 2: Uncertainty in Measurement (error)
A measurement always has some degree of uncertainty.
The uncertainty of a measurement depends on the precision of the measuring device.
For example, using a bathroom scale, you might estimate the mass of grapefruit to be approximately 1.5 pounds.
Weighing the same grapefruit on a highly precise balance might produce a result of pounds.
In the first case, the uncertainty occurs in the tenths of a pound place;
In the second case, the uncertainty occurs in the thousands of a pound place.

9. Example 1.1: Uncertainty in Measurement

10. Example 1.1: Uncertainty in Measurement

11. Precision and Accuracy
Precision: Determines how closely measurements agree with each other, Reflects the reproducibility of a given type of measurement.
1st series of measurements: 34, 35, 37, 37, 38
2nd series of measurements: 30, 35, 40, 42, 47
The precision of the 1st series is better than the 2nd series.
Accuracy: Determines the agreement of a particular value
with the true value (standard value). Example: True Value =
36.0
Average = Sum of all the values / number of values
example: 34, 35, 37, 37, 38
sum of all the values = 181
number of values = 5
average = 36.2
true value = 36.0  good accuracy. 11

12. Precision and Accuracy
Not Precise, Not accurate
Precise, but not accurate
Accurate, and precise
The Difference between Precision and Accuracy 12

13. Random Vs. Systematic Error
We have 2 main types of errors: Random and Systematic.
A random error, it is also called indeterminate error, means that a measurement has an equal probability of being high or low.
Random error affects the precision of the measurement.
The systematic error, is also called determinate error. This type of error occurs in the same direction each time: It is always high or low.
Systematic error affects the accuracy of the results.

14. Example 1.2
Suppose we weigh a piece of brass 5 times on a very precise balance and obtain the following results:
Normally, we assume that the true mass of the piece of brass is very close to the average of the 5 results:

15. Example 1.2 (cont’d)
However, if the balance has a defect causing it to give a result that is consistently 1.00 gram too high (a systematic error of g), then the measured value of 2.486g would be seriously in error. High precision among several measurements is an indication of accuracy only if systematic errors are absent.

16. Example 1.1: Precision and Accuracy
To check the accuracy of a graduated cylinder a student filled the cylinder to the 25-mL mark using water delivered from a buret, and then read the volume delivered. (This is called calibration). Following are the results of five trials:
Is the graduated cylinder accurate?

17. Example 1.2: Solution
The results obtained for the graduated cylinder are very precise. The student has good technique. However, note that the average value measured (26.54) using the buret is significantly different from 25 mL. Error = – 25= 1.54 mL Thus, this graduated cylinder is not very accurate. It produces a systematic error (the indicated result 25mL is low for each measurement).

18. Rules for Counting Significant Figures
All non-zero digits are significant figures: 1234  4 significant figures
Zeros between non-zero digits (captive zeros) are significant figures: 205  3 significant figures
Zeros beyond decimal point at the end of the number (trailing) are significant figures:  5 significant figures
Zeros preceding the first non-zero digit in a number are not significant figures:  3 significant figures
Zeros at the end of whole numbers are not significant figures unless you are given information to the contrary: 3400  2 significant figures, X 102  4 significant figures,  4 significant figures. 18

19. How many significant figures are in each of the following?
12  2 significant figures (S.F.)
1098  4 S.F.
2001  4 S.F.
2.001 x 103  4 S.F.
 3 S.F.
1.01 x 10-5  3 S.F.
 4 S.F. (because of the decimal point).
 7 S.F. 19

20. Rules for Significant Figures in Mathematical Operations
Multiplication and division: The number of significant figures in the result is the same as that in the quantity with the smallest number of significant figures.
Ex.: x =  6.4
(3 S.F.) (2 S.F.) (3 S.F.) (2 S.F.)
The product should have only two significant figures since 1.4 has two significant figures.
Addition and subtraction: The result has the same number of decimal places as the least precise measurement used in the calculation.
Ex.:   31.1
The correct result is 31.1, since 18.0 has only one decimal place. 20

21. Rules for Rounding of Data
In a series of calculations, carry the extra digits through
to the final result, then round.
If the digit to be removed;
a. Is less than 5, the preceding digit stays the same.
Example rounds to 1.3.
b. Is equal to or greater than 5, the preceding digit is increased by 1.
Example rounds to 1.4. 21

22. Example 1.4: Significant Figures in Mathematical Operations
Carry out the following and give each result the correct number of significant figures.
a) 1.05 x 10-3 ÷ 6.135
The correct answer is: 1.71 x 10-4 , which has three significant figures because the term with the least precision (1.05 x10-3 ) has three significant figures. b)
The result is 7 with no decimal point, because the number with the least number of decimal places (21) has none.

23. Exponential Notation (Scientific Notation)
The number 1.00 x 102 is written in exponential notation. This type of notation has at least 2 advantages: 1- The number of significant figures can be easily determined. 2- Fewer zeros are needed to write a very large or very small number. For example, the number is much more conveniently represented as 6.0 x 10-5 (it has two significant figures).

24. Units Conversion How do you convert 1.53 minutes to seconds?
a. Find a conversion factor (or factors) :
60 sec = 1 min
b. Set up start-up and ending information with units 1.53 min. = sec
c. We need an answer in ‘sec’ and we need to get rid of ‘min’.
Therefore, 1.53 min X 60 sec/1 min = 91.8 sec. 24

25. Units Conversion
It is often necessary to convert from one system of units to another.
This can be done by a method called the unit factor method.
Example: Consider a pin measuring 2.85 cm in length. What is its length in inches?
To accomplish this conversion we must use the equivalence statement cm = 1 in.
If we divide both sides of this equation by 2.54 cm, we get:
This expression is called a unit factor

26. Unit Conversion
Multiplying any expression by this unit factor does not change its value.
The pin has a length of 2.85 cm. Multiplying this length by the appropriate unit factor gives:
Note that the centimeter units cancel to give inches in the result. This is exactly what we want to accomplish.

27. Example 1.7: Unit Conversion III
A student has entered a 10.0-km run. How long is the run in miles? 1km = 1000 m 1 m = yd 1760 yd = 1 mi First step convert 10 km to meter: Then convert meters to yards: Finally, covert the yards to miles: Since we the distance originally has 3 Sig. Fig. then the result should be rounded to 6.22 mi.

28. Temperature Measurement
Three systems for measuring temperature are given below;
The Celsius scale (oC)
The Kelvin scale (K)
The Fahrenheit scale (oF)
Following Four equations (formula) are used for introversion of
temperature scales.
TK = TC (1)
TC = TK – (2)
TF = TC X 9 oF / 5 oC + 32 oF (3)
TC = (TF – 32 oF ) 5 oC / 9 oF (4) 28

29. The Three Major Temperature Scales
Cont.
The Three Major Temperature Scales 29

30. TC = (TF - 32 oF) 5 oC / 9 oF Example:
A person has a temperature of oF. What is this
temperature on the Celsius scale? On the Kelvin scale?
By using equation;
TC = (TF - 32 oF) 5 oC / 9 oF
= (102.5 oF - 32 oF) 5 oC / 9 oF = 39.2 oC
TK = TC
= ( ) K = K 30

31. Density and Matter
Density: The mass of substance per unit Volume of the
substance
Density = mass(g)/volume(cm3)  g/cm3
Density is often used as an “identification tag” for a substance.
Matter: Anything occupying space and having mass. Matter exits in three states:
1. Solid is rigid, it has a fixed volume and shape.
2. Liquid has a definite volume but no specific shape, it assumes the shape of its container.
3. Gas has no fixed volume or shape, it takes on the shape and volume of its container. 31

32. Types of Mixtures
Homogeneous mixture: Having visibly indistinguishable parts. Physical properties are the same throughout the material. A homogeneous mixture a solution (example: vinegar).
Heterogeneous mixtures: Having visibly distinguishable parts. Physical properties are different at different points in a material (example: bottle of ranch dressing). 32

33. Definitions
Pure substance: is one with constant composition. Pure substances can be isolated by separation techniques – distillation, filtration, chromatography.
Compound: is a substance with constant composition that can be broken down into elements by chemical processes. Example: electrolysis of water produces hydrogen and oxygen.
Physical change: is a change in the form of a substance but not in its chemical composition. E.g. conversion of ice into water.
Chemical change: when a given substance becomes a new substance or substances with different properties and different composition. Burning of wood or gas. 33

34. Chart of Organization of Matter

35. Recommended Exercises from the End of Chapter 1:
On page 32, Exercise # : 29, 30, 31, 32, 33, 34, On page 33, Exercise # : 35, 36, 42 On page 34, Exercise # : 51, 52, 53, 55, 63, 65, 67 On page 35, Exercise # : 77