1. FUNDAMENTALS OF CHEMISTRY
Chemical Foundations BY
Dr. Ghulam Abbas
Assistant Professor
UNIVERSITY OF NIZWA

2. Atoms and Electrons
Hydrogen
Nitrogen-7
Carbon-6
Oxygen-8

3. 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
Chemsitry 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 has some mass and occupies
some space. Examples: water, gold, NaCl, sugar, Air etc.

4. Scientific Method 4.Theory: A set of tested hypotheses that explains
To understand any process that lies at the center of
scientific inquiry.
Steps in the Scientific Method are:
1. Making observations: (collecting data).
2. Hypothesis: A possible explanation for an observation.
3. Performing experiments: to test the prediction or
hypothesis (testing the hypothesis).
4.Theory: A set of tested hypotheses that explains
some natural phenomena. ( summary of why it
happened).
5. Law: observed behavior formulated into statement is called Law (Summary of what happened).

5. Observations: are of two types;
1- Qualitative Observations: It does not involve a number.
Example; sky is blue, Water is liquid.
2- Quantitative- involves both a number and a units. Example: Water boils at 100 oC, road length is 100 kilometer.
The Various Parts of the Scientific Method

6. Fundamental Quantities and Units
Physical Quantity Units Abbreviations
Mass kilogram kg
Length meter m
Time second s
Temperature Kelvin k
Electric Current Ampere A
Amount of substance mole mol
Luminous Intensity candela cd

7. SI System
In 1960, an international agreement set up a system of units called the International System or the SI system. This system is based on the metric system and units derived from the metric system.
example: Mass unit is kilogram(kg), length unit is meter(m). Volume is not a fundamental SI unit but is derived from length. Cubic meter (m3) is more commonly used is liter/(L).
1 liter= 1 dm3 = (10cm)3 = 1000 cm3
1 cm3 = 1 mL
1000 mL = 1 Liter

8. 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 8

9. Precision and Accuracy
It determines how closely several Measurements (results) of
the same quantity agree with each other. It shows 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:
It determines how closely several Measurements (results) of the same quantity agree with the true value (standard value). Example: True Value = 37.0
Average = Sum of all the values / number of values
Example: 34, 36, 39, 40
sum of all the values = 149 number of values = 4
average = 149/4 = 37.2
true value = 37.0  good accuracy. 9

10. Precision and Accuracy
No Precision No accuracy
Precision but not accuracy
Accuracy
The Difference between Precision and Accuracy 10

11. Errors
Random Error (indeterminate error): A measurement has an equal probability of being high or low. This type of error occurs in estimating the value of the last digit of measurement.
Systematic Error (Determinate error): This type of error occurs in the same direction each time. It is either always high or always low, often resulting from poor technique. 11

12. Rules for Counting Significant Figures (number of digits that contribute to precision)
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 in different way; 3400  2 significant figures, X 102  4 significant figures,  4 significant figures. 12

13. 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. 13

14. 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. 14

15. 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. 15

16. 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) 16

17. 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. 17

18. The Three Major Temperature Scales
Cont.
The Three Major Temperature Scales 18

19. 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 19

20. Density and Matter
Density: The mass of substance per unit Volume of the
Substance D= m/v
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. 20

21. MIXTURES PURE COMPOUNDS
MIXTURE and COMOUNDS
A MIXTURE is a combination of two or more substances that are not chemically united and do not exist in fixed proportions to each other. Most natural substances are mixtures. Example air is the mixture of several gases, petroleum etc.
 MIXTURES
 PURE COMPOUNDS
A mixture can be physically
separated into pure
compounds or elements.
 A pure compound has a
constant composition with
fixed ratios of elements.
 Mixtures may exhibit a
changing set of physical
properties.
For example, mixture of
alcohol and water boils
over a range of
temperatures.
Physical properties such as
boiling point or melting
point of pure substances
are invariant.
For example, pure water
boils at 100 degrees oC.

22. 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). 22

23. 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: 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. 23

24. Uncertainty in measurement
What is a measurement?
A measurement tells us about a property of something. It might tell us how heavy an object is, or how hot, or how long it is. A measurement gives a number to that property.
What is uncertainty of measurement?
The uncertainty of a measurement tells us something about its quality.
Uncertainty of measurement is the doubt that exists about the result of any measurement. You might think that well-made rulers, clocks and thermometers should be trustworthy, and give the right answers.
Every measurement has some degree of uncertainty. 24

25. Uncertainty in measurement
person
Results of measurements 1
43.30 2
43.40 3
43.35
The above results show that the first two numbers (43) remain the same regardless of who makes the measurement; these are called as CERTAIN digits. However, the digits to the right side of the decimal point must be estimated and therefore varies, called as an UNCERTAIN digits.
Certain digits are known as Significant figures of a measurement. The uncertainty of any measurement is indicated by ±value.
For example, in above case the uncertainty is 43 ± 0.1
You may be interested in uncertainty of measurement simply because you wish to make good quality measurements and to understand the results.
Two terms are usually used to describe the reliability of any measurement; precision and accuracy. 25

26. When you know chemistry, there’s a new level of looking at the world around you.