1. Chemistry 101 : Chap. 1 Matter and Measurement
What is Chemistry and Why we study it
(2) Classification of Matter
(3) Properties of Matter
(4) Units of Measurement
(5) Uncertainty in Measurement
(6) Dimensional Analysis

2. The Study of Chemistry
Chemistry: The study of the properties of matter and the
changes that matter undergoes.
 Matter : Physical material of the universe
Anything that has mass and occupies space
 Changes in Matter : Physical or Chemical changes
Why Chemistry?
 Chemistry is the central science
 Chemistry is a practical science and has profound
impact on our daily living

3. Macroscopic vs. Microscopic
Macroscopic World : Realm of ordinary-sized object.
Things we can see with the naked eye.
 (Sub)Microscopic World : Realm of atoms and molecules
Carbon nanotube (10-9 m)
Chemistry is the science that seeks to understand the properties
and behavior of matter (macroscopic) by studying the properties
and behaviors of atoms and molecules (microscopic)

4. Major Divisions in Chemistry
 Physical Chemistry (CHM321, CHM420)
 Organic Chemistry (CHM211, CHM212)
 Inorganic Chemistry (CHM 455, CHM546)
 Analytical Chemistry (CHM235, CHM435)
 Biochemistry (CHM365, CHM568)
All divisions are interrelated and cannot be
standing alone.

5. Classification of Matter: pure substance vs. mixture
Pure Substance: A sample of matter that has distinct
properties and a composition that doesn’t vary from sample
to sample (either element or compound)
Elements: A pure substance that cannot be
decomposed into simpler substances. The basic unit
of an element is an atom.
Nitrogen atom
Nitrogen molecules
Argon gas (atoms) Nitrogen gas (molecules)

6. Classification of Matter: pure substance vs. mixture
Compound : Substances that are composed of
two or more elements. The basic unit of compound is
a molecule
Mixture : Combinations of two or more substances
in which each substance retains its own chemical identity.
nitrogen atom
Two or more elements
(compound)
Two or more substances
(mixture)
hydrogen atom
ammonium (molecule)

7. Elements  At the present time, there are 116 elements
Periodic Table of the Elements
= H2, N2, O2, F2, Cl2, Br2, I2

8. Elements
 Not all elements are equal…

9. Compounds  Most elements can interact (or react) with other elements
to form compounds
Example: Combine hydrogen & oxygen to generate water
Oxygen Hydrogen water
However, elemental hydrogen and oxygen exist as diatomic
molecules (H2 and O2) in nature. +
O H2
2H2O

10. Mixture  Components: The substances making up a mixture
Homogeneous Mixture (solution) : Uniformly distributed
throughout. (air, salt solution, sugar solution …)
Heterogeneous Mixture : Do not have the same
composition, properties and appearance throughout.
(rock, wood …)
Oil on water
Air

11. Classification of Matter

12. Classification of Matter
 Example
14 K gold
(2) Orange Juice
(3) A cup of coffee
(4) Mud

13. Separation of Mixture
Separate a mixture into its components by taking advantage
of the difference in their properties
Filtration : Separation is based
on the size of particles in the
mixture. Filtration is used with
heterogeneous mixtures

14. Separation of Mixture  Distillation : Separation is based
on the boiling points of the
components in the mixture.
Distillation is typically used
with homogeneous solutions.
Water changes its
states from gas to
liquid

15. Separation of Mixture
 Chromatography : Separation is based on the solubilities
of the components in the mixture. It is normally used with
homogeneous mixture.
Paper chromatography

16. Classification of Matter:
states of matter
States of matter: A sample of matter can have three
physically different states
 Gas : Indefinite volume and indefinite shape
(depends on the volume and shape of its container)
 Liquid : definite volume, but indefinite shape.
 Solid : definite volume and definite shape
Pure substance can have any state depending
on the temperature and pressure

17. Three States of Water

18. Properties of Matter
Physical properties : They can be measured without
changing the identity and composition of the substance
Ex. color, order, density, boiling point…
 Chemical properties : They describe the way a substance
can change or react
Ex. flammability, solubility, …

19. Physical vs. Chemical Properties
 Example : Zinc (Zn)
silver-grey metal
melting point: 420oC
reacts with oxygen to
form Zinc oxide (ZnO)
density (25oC) = 7.13 g/cm3
generates hydrogen when
dissolved in sulfuric acid

20. Properties of Matter Extensive properties : Properties that depend on
the quantity of a sample.
Ex. Volume :  +  = 
V1 + V2 = V1 + V2
 Intensive properties : Properties that are independent
on the quantity of a sample
Ex. Temperature :  +  = 
T T T

21. Extensive vs. Intensive Properties
 Example :
Boiling/melting point (bp/mp)
Mass
Density
Pressure

22. Changes of Matter Physical changes : Phase changes, but it is
still H2O (no change in its
composition)
Chemical changes :
Aluminum (Al) reacts with
Bromine (Br2). (A substance
is transformed into a chemically
different substance: AlBr3)

23. Units of Measurement : SI Unit
Système International (SI) d’Unités
International agreement on the metric units for the
uses in science (1960)

24. Units of Measurement : Prefixes
 Prefixes : They are used to indicate decimal fractions
or multiples of various units.
A Megabyte of memory : 106 bytes of memory
Femtochemistry : chemistry that occurs on the time scale of second
check out (prof. Zewail’s homepage)

25. Length and Mass Length : 1 meter (m) = 100 cm
Mass : 1 kilogram (kg) = 1000 g
Metric to English conversion
1 m = yard
1 cm = inch
1 kg = lb
Check out
NOTE: Mass and weight are not the same thing. Mass is an intrinsic
property of matter, but weight depends on the gravity.

26. Temperature Water freezing Water boiling Celsius scale (oC) 0 100
Fahrenheit scale (oF)
oC = 5/9 (oF  32) oF = 9/5(oC) + 32
Kelvin : K = oC (exact)
Absolute zero temperature : 0 K =  oC
The lowest attainable temperature in our universe

27. Temperature (98.6 oF  32)5/9 = 37 oC 37 oC + 273.15 = 310.15 K
William Thomson Kelvin
( )
“On an Absolute Thermometric Scale”
Philosophical Magazine, vol. 1
pp (1848)
(98.6 oF  32)5/9 = 37 oC
37 oC = K

28. Derived Units Use the defining equation for the quantity of interest
and substitute the appropriate SI units
 Volume: abc = (length)3 = m3 a b c
In chemistry, we normally use
smaller units.
Liter : (10 cm)3 = 1 L = 1 dm3 = 10-3 m3
1 gal = 3.8 L
(2) Milliliter = 1 mL = 10-3 L = 1 cm3 = 1 cc

29. Derived Units Density : The amount of mass in a unit volume of
substance
SI unit of
density
 In chemistry, we typically
use g/mL = g/cm3 = g/cc
 Density depends on
temperature
 Don’t be confused about
density and weight

30. Density, Volume and Mass
(1) 1.00  102 g of mercury occupies a volume of 7.36 cm3. What is
the density of mercury?
(2) The density of liquid methanol is g/mL. What is the volume
of 65.0 g of liquid methanol?
(3) The density of gold is g/cm3. What is the mass in gram of a
cube of gold if the length of the cube is 2.00 cm?

31. Uncertainty in Measurement
We need to distinguish two different types of
number in science
 Exact Number : Defined number
1 dozen = 12, 1 m = 100 cm
Counted number
There are 120 students in the class.
 Inexact Number : Numbers from measurement
(human errors, machine errors..)

32. Precision and Accuracy
 Precision : How closely individual measurements agree
with one another.
 Accuracy : How closely individual measurements agree
with the correct or “true” value.
good precision
poor accuracy
good precision
good accuracy
poor precision
poor accuracy

33. Significant Figures
Measured quantities are generally reported in such a
way that only the last digit is uncertain.
mass of a dime = g
Uncertain. Could be 6 or 4…
(2) Sometimes,  sign is used to specify the uncertainty.
mass of a dime =  g
Significant Figures : All digits of a measured quantity,
including the uncertain one.
g  5 significant figures

34. Rules for Significant Figures
All non-zero digits are significant
(2) Zeros at the beginning of a number are never significant
 count the digits starting with the first non-zero digit
has TWO significant figures
(3) Zeros between non-zero digits are significant
 has THREE significant figures
(4) Zeros at the end of a number are significant.
 has FOUR significant figures
2060 has FOUR significant figures
2.06 x 103 has THREE significant figures

35. Significant Figures in Calculation
The number with the fewest number of significant figures
limits the certainty of the calculated quantity.
Multiplication & Division : The final answer can have
no more significant figures than the fewest number of
significant figures in any number in the problem.
 Addition & Subtraction : The final answer can have
no more decimal places than the fewest number of
decimal places in any number in the problem

36. Significant Figures in Calculation
Example 1: Area of a rectangle whose measured edge
lengths are cm and 5.2 cm
Area = (6.221 cm) x (5.2 cm) = cm2 =
Only 2 significant
figures
Include only 2
significant figures
Example 2 : Addition of three measured numbers
20.42
1.322
+ 83.1 

37. Significant Figures in Calculation
When calculation involves multiple steps…
Retain at least one more extra digit (past the number
of significant figure) in each step
 When you use a calculator…
Enter the numbers one after another (without
worrying about significant figures) and rounding
only the final answer

38. Significant Figures in Calculation
Example 3: [1255  (3.45  108)]
= 863  [1255  372.6]
= 863  882.4 = =
Example 4: (  ) + (2813  12) =
= 33890 =
From calculator = =

39. Dimensional Analysis
We carry units through all calculations. Units behave
like numbers: they are multiplied together, divided
into each other, or canceled.
Example: How many inches are in 10 cm?
Correct Wrong
Advantages of dimensional analysis
(1) It ensures that your answer has the correct unit
(2) It makes it easier to find out possible errors

40. Unit Conversion Example: The speed of N2 in air at 25 oC is 515 m/s.
Conversion factor
Example: The speed of N2 in air at 25 oC is 515 m/s.
Convert the speed into mile/hour

41. Unit Conversion Example: The density of water is 1.00 g/mL.
What is the mass 1.00 gal of water in grams?

42. An example
The density of gold is g/cm3. If 2.00 g of gold wire has 0.12 mm
radius, how long the wire is?