1. Introduction: Matter and Measurement
Chapter 1
Introduction: Matter and Measurement
Dr. RAJANI SRINIVASAN
Tarleton State University

2. Chemistry the Central Science
Chemistry is the central to fundamental understanding of other science and technologies.
It shares ties with almost all the branches of science and engineering
“Biology, Medical, Agriculture, Pharmacy, Forensic science, Rocket science, Engineering, Environmental Science etc.”

3. Chemistry ?????
Chemistry is the branch of science which studies the properties and behavior of anything that comes in contact with you whether you can see it or not.
That anything is “Matter”

4. Matter
We define matter as anything that has mass and takes up space

5. Matter Matter is composed of the smallest particle called Atom
Two or more atoms combine together in a particular shape and angle to form Molecules

6. Classification of Matter
Physical State
Composition
Element , Compound and mixture
Gas, Liquid and solid

7. States of Matter Gas: has no fixed volume, takes the shape of the
container, Can be compressed and expanded
Example : water vapor
Liquid : Has distinct volume independent of the container , takes the shape o f the container
Example : water
Solid : Has definite volume and Shape
Example : Ice
Note: Both Liquid and solid cannot be compressed or expanded.

8. States of Matter
Element : Substances which cannot be decomposed further . They are formed of only one type of atoms.
Example : Oxygen , Silicon etc.
Currently 118 elements are known
Over 90% of Earth’s crust is composed of Oxygen, silicon, aluminium, iron and calcium
Oxygen Carbon and hydrogen forms 90% of human body mass

9. States of matter
Compound: Substances formed by the combination of two or more elements . They can be decomposed to their individual components
Example :
When Oxygen and Hydrogen combine together they form water and when electricity is passed through water they decompose to Oxygen and Hydrogen

10. States of matter

11. States of Matter
Mixture: Made up of different substances and each substance retains its chemical identity and property. Example : Coffee
Mixture
Heterogeneous
Homogenous
Substances which do not have same composition throughout .
Eg: rocks and wood
Substances which has uniform composition throughout .
Eg: Air

12. Classification of matter

13. Properties of matter Physical Properties… Chemical Properties…
Can be observed without changing a substance into another substance.
Boiling point, density, mass, volume, etc.
Chemical Properties…
Can only be observed when a substance is changed into another substance.
Flammability, corrosiveness, reactivity with acid, etc.

14. Properties of matter Intensive Properties… Extensive Properties…
Are independent of the amount of the substance that is present.
Density, boiling point, color, etc.
Extensive Properties…
Depend upon the amount of the substance present.
Mass, volume, energy, etc.

15. Changes in matter Physical Changes
These are changes in matter that do not change the composition of a substance.
Changes of state, temperature, volume, etc.
Chemical Changes (Chemical Reaction)
Chemical changes result in new substances.
Combustion, oxidation, decomposition, etc.

16. Chemical Changes
In the course of a chemical reaction, the reacting substances are converted to new substances

17. Separation of mixtures
Criteria for separation of mixtures
Different methods for different types of mixtures
Steps: Determine
Use the difference in the properties in the mixture to separate them.
Heterogenous
Homogenous

18. Separation of Mixture
Example : Heterogeneous mixture of gold and Iron fillings: Properties (Physical):1) Iron magnetic gold non magnetic This property can be used to separate them 2) Chemical Property: Gold dissolves in some acids in which iron does not

19. Filtration Useful for heterogeneous mixture
Solid and liquid mixture can be separated using filtration
In filtration, solid substances are separated from liquids and solutions

20. Distillation Very useful for separating homogenous mixture
Used the property of substance to form gasses.
Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components
Boiling water and salt at 100C water stars evaporating we can condense and get water salt remains in the flask and can be derived from it.

21. Chromatography
This technique separates substances on the basis of differences in solubility in a solvent.

22. Units of measurement Système International d’Unités
There are seven base units from which all the other units are derived
A different base unit is used for each quantity
In this chapter we will consider Length , mass and temperature

23. Metric System

24. Length and Mass SI unit for Length is meter
Mass = measure of amount of material in an object. (Note mass is different from weight)
[ Weight = Force that is exerted by its on its mass by its gravity]
SI unit for mass is Kilo gram

25. Brain teaser !!!!
Question: What is the name of the unit that uses the prefix
a) gram – Nano- ng
b) 10-6 Second – micro - µg
C) 10-3 meter – milimeter- mm

26. Temperature
Temperature is the measure of degree of hotness and coldness of the object
Heat always flows from higher temperature to the lower temperature spontaneously.
Units for measurement of temperature
Celsius Scale - ◦ C
Kelvin
Fahrenheit

27. Temperature
In scientific measurements, the Celsius and Kelvin scales are most often used.
The Celsius scale is based on the properties of water.
0 C is the freezing point of water.
100 C is the boiling point of water

28. Temperature …… The kelvin is the SI unit of temperature.
There are no negative Kelvin temperatures.
K = C
The Fahrenheit scale is not used in scientific measurements.
F = 9/5(C) + 32
C = 5/9(F − 32

29. Temperature…
Note : Zero Kelvin is the lowest attainable temperature = Absolute Zero = ◦C
Celsius and Kelvin have equal sized units

30. Derived units
Units that are derived from the basic seven SI units are called derived units.
Density is a physical property of a substance.
It has units (g/mL, for example) that are derived from the units for mass and volume.
d = m V

31. Derived units Speed = Ratio of distance travelled to lapsed time .
Speed = meter /second = Length /time

32. Uncertainty in measurement
Exact Numbers – Exact Known value- Dozens
Inexact Numbers – Numbers with uncertainty
Instrument errors – all the equipment have measuring error
Human errors – no two human beings can have same measurement
Accuracy – How closely individual measurements agree with the correct value
Precision –measure of how closely individual measurements agree with one another

33. Significant figures
When we measure a dime in the balance we report the mass as ±0.0001g
Symbol ± represents the uncertainty in measurements.
All the digits in the measured quantity including the uncertain ones are called Significant figures.
Greater the significant figures greater certainty implied for the measurement

34. Significant numbers
When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.
All nonzero digits are significant –
significant figures
Zeroes between two significant figures are themselves significant sf; 7.03 = 3 sf
Zeroes at the beginning of a number are never significant = sf ; 0.03= 1 sf they just indicate the number of decimal
Zeroes at the end of a number are significant if a decimal point is written in the number = 2.0 – 2sf

35. Significant numbers
If a number ends with zero but has no decimal then zeros are not considered significant.
Note: Exponential notation helps us to understand whether zeros are significant or not
Example: 10,300 based on the measurements they can have variable significant figures
A) 1.03*10 4 g = 3 sf
B) * 10 4 g = 4 sf
C) * 105g = 5 sf

36. Significant figures
When addition or subtraction is performed, answers are rounded to the least significant decimal place.
When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

37. Dimensional analysis
Using same units to solve the problems to have accurate answers
We use dimensional analysis to convert one quantity to another.
Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm

38. Dimensional analysis
Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm version factors to be used
1 in.
2.54 cm or

39. Dimensional analysis We know 100 cm = 1m or 1m = 1 /100 cm
Q: Convert 8.00 meter rod to inches
We know 100 cm = 1m or 1m = 1 /100 cm
1inch = 2.54 cm

40. Problems
Use the form of the conversion factor that puts the sought-for unit in the numerator:
Given unit 
 desired unit
desired unit
given unit
Conversion factor

41. Problems For example, to convert 8.00 m to inches, convert m to cm
convert cm to in
8.00 m 
100 cm
1 m
1 in.
2.54 cm 
315 in.

42. Conversion involving volume
We can also have conversion factor from one measure to different one
Example Density D= m/V
What will be the mass in grams of 2 cubic inches of gold with density of 19.3 g/cm3
Conversion factor= 19.3g/1cm3 and 1cm3/19.3 g
We know 2.54 cm3/1inch3 = cm3/1inc3
Mass in grams = 2in3*16.39cm3/1in3* 19.3g/1cm3 = 633 g