Chapter 1: UNITS OFMEASUREMENTS
Chem100 - Lecture
By: Jeanette L. Gadiane
Objectives
At the end of the chapter, the students should be able to:
 Distinguish SI system of measurement from English system
 Determine figures which are significant
 Round-off quantities correctly
 Convert from one system to the another using factor-label method
 Interconvert temperatures in the Fahrenheit, Celsius and Kelvin
scales
Introduction
Almost everything we own –
clothes, house, food, vehicle- is
manufactured with measured parts,
sold in measured amounts, and paid for
with measured currency!
A. The International System of Units

It is a modified system of measurement based in metric units, that
is, using the decimals or by powers of 10.

There are seven fundamental units or base units in SI system, each of
which is identified with a physical quantity.These are the length,
mass, time, electrical current, temperature, amount of substance,
and luminous intensity.

All other units are derived units which are combinations of these
seven base units.
SI Base Units
BASE QUANTITY NAME OF UNIT SYMBOL
Length
Mass
Time
Electrical current
Temperature
Amount of
substance
Luminous intensity
Meter
Kilogram
Second
Ampere
Kelvin
Mole
m
kg
s
A
K
mol
Candela
cd
Prefixes Used with SI Units
PREFIXES SYMBOL
TeraT
GigaG
Mega- M
Kilok
Hecto
h
Deka
da
Decid
Centi- c
Millim
Micro- μ
Nano- n
Picop
MEANING
EXPONENTIAL
1,000,000,000,000,
1,000,000,000,
1,000,000,
1,000,
100
10
1012
109
106
103
102
101
1/10,
1/100,
1/1,000,
1/1,000,000,
1/1,000,000,000,
1/1,000,000,000,000
10-1
10-2
10-3
10-6
10-9
10-12
Commonly Used SI Units in the Study of
Chemistry
Length
 Meter (m) – SI base unit
 Standard meter is the distance light travels in a vacuum in
1/299,792,458 second
 Biological cells are often measured in micrometers ( 1 μm = 10-6
m)
 Atomic-sized scale, nanometer and picometer are used ( 1 nm =
10-9 m; 1pm = 10-12 m)
Volume



Cubic meter (m3) – SI unit
liter (L) and milliliter (mL) – the most important volumes used in
chemistry
cubic decimeter (dm3) – used in medical field
1 L = 1 dm3 = 10-3 m
Mass


Mass - refers to the quantity of matter it contains
Weight – the force exerted by a gravitational field
: therefore mass is constant while weight varies!
on an object

kilogram (kg) – is the SI unit
- Its standard is a physical object – the platinum-iridium
cylinder kept in france
Density

Density = mass/volumea derived unit
 kg/m3 - SI unit of density
 g/L (g/dm3) or g/mL (g/cm3) – commonly used in Chemistry
B. SIGNIFICANT FIGURES
Significant figures – are the figures that are both certain and
uncertain one.
Significant Figures
Figure TA 1.2
Copyr ight © 2001 T he McGr aw- Hill Com panies, Inc. Permission required for reproduction or display.
The more significant figures in a quantity, the
more accurate is the measurement.
Rules in DeterminingSignificant Figures
Determining which digits are significant:
1.
All non-zero digits are significant
2.
All zeros between a non-zero digits are significant
3.
All zeros to the right of a non-zero digit after the decimal point
are significant
4.
In the absence of decimal point, significant figures vary
Sample Problems
How many significant figure are there in:
1.
2.
3.
435 m
0.0120 L
27,075 g
Seat work
1.
0.0030 L
2.
0.1044 g
3.
53,069 mL
4.
2.00 x 103 cm
5.
57,600 s
Rules for Significant Figures in
Answers
For multiplication and division: The answer contains the same number of
significant figures with the measurements that contains the fewest number of
significant figures.
Ex.
Volume (cm3) = 9.2 cm x 6.8 cm x 0.3744 cm= 23
cm3

Least no .of significant figures

For addition and subtraction: The answer has the same number of decimal
places with the measurement with the fewest decimal places.
Ex.
Volume (mL) = 83.5 mL + 23.28 mL
= 106.8 mL
least no. of decimal places
Rules for Rounding Off
1.
If the digit removed is more than 5, the preceding number is
increased by 1.
Ex. 5.379 rounds to 5.38 for three s.f., and 5.4 for four s.f.
2.
If the digit removed is less than 5, the preceding number is
unchanged.
Ex. 0.2413 round to 0.241 for three s.f., 0.24 for two s.f.
3.
If the digit removed is 5, the preceding number is increased by 1.
4.
Always carry one or two additional significant figures through a multi-step
calculation and round off the final answer only.
Round-off the following:
1.
The product of 12.2 x 4.0230
2.
The sum of 12.2 + 4.0230
3.
The sum of 1.2 x 10-3 + 1.4 x 10-2
Accuracy versus Precision
 Accuracy refers to the proximity of a measurement

to the true value of a quantity.
Precision refers to the proximity of several
measurements to each other.
Calculations...



Factor Label method
Metric-Metric Conversion
English- SI Conversion
Factors
Common SI-English Equivalents
Examples
Convert the following quantities:
1)
1500 mL = L
2)
4.67 x 104 μg = g
3)
4)
5)
6)
7)
8)
9)
10)
3.15 x 10-5 Mm = km
2.10 x 106 pm = μm
0.89 g = mg
3.45 cm = ft
0.0678 fl oz = mL
45 g = lb
1.5 qt = mL
0.075 mi = m
C. Temperatures
Heat – the energy in transfer
Temperature – the hotness or coldness of a body.
Thermometer – a device used to measure the temperature of a
certain body.
Temperature Scales
Celsius (oC)
Fahrenheit (oF)
Kelvin (K)
Temperature Scales

Fahrenheit (F): freezing point of water is 32°F and the boiling point of
water is 212°F

Celsius (C): freezing point of water is 0°C and boiling point of water is
100°C

Kelvin (K): zero is the lowest possible temperature; also called the
absolute scale
◦ degree is the same size as Celsius degree
◦ K = °C + 273
Fig. 1.6, p.11
Formulas
9
°F = _ °C + 32
5
K = oC + 273
5
°C = _ (°F - 32)
9
Examples
Convert the following temperatures:
1.
40.5 oC to oF
2.
124 oF to oC
3.
205.9 oF to K
4.
A child has a body temperature of 38.7oC.
a. If the normal body temperature is 98.6oF, does a child has a
fever?
b. What is the child’s temperature in Kelvin?
Density and Specific Gravity
Calculations
Density – mass of the object divided by its volume.
D = M/V
 It is used to characterize or identify a substance since its substance
has a corresponding unique density.
example: during gold rush era, density is used to distinguish a real
gold from a “fool’s gold” (pyrite).
Examples
1.
2.00 cm3 of aluminum are found to weigh 5.40 g. Calculate the density
of aluminum in units of g/cm3.
2.
Air has a density of 0.0013 g/mL. What is the mass of 6.0-L sample of
air?
3.
Calculate the mass in grams of 10.0 mL of mercury if the density of
mercury is 13.6 g/mL.
4.
Calculate the volume in mL of a liquid that has a density of 1.20 g/mL
and a mass of 5.00 g.
Exercise
1.
The density of ethyl alcohol is 0.789 g/mL at 20oC. Calculate the
mass of a 30.0-mL sample.
2.
Calculate the volume in mL of 10.0 g of a saline solution that has
a density of 1.05 g/mL.
Specific Gravity
A reference density based on the density of water (1.00 g/mL) at 4oC.
 Unitless
 It is measured by a hydrometer (consists of a weighted glass bulb that is
inserted into a liquid and allowed to float).
 Urinometer – a hydrometer used to measure urine samples.
density of obj. (g/mL)
Specific gravity=
density of H2O (g/mL)

Application: routine hospital tests including measurement of the specific
gravity of urine and blood samples are frequently used as diagnostic
tools.
Examples
1.
The density of ethanol at 20oC is 0.789 g/mL. What is its specific gravity?
2.
The specific gravity of a urine sample is 1.016. What is its density in
g/mL?
3.
The normal range for normal urine specific gravity is 1.010 – 1.030.
Given that the a 335.0 cc sample of urine has a mass of 342.6 g, is the
urine sample fall within the normal range?
4.
The density of dichloromethane (an organic substance), a liquid
insoluble water, is 1.33 g/cc. If dichloromethane and water are placed in
a separatory funnel, which will be the upper layer?