1.
Chapter One Part II
Measurements in
Chemistry
Chem101: General & Consumer
Chemistry
Natural Sciences Department
College of Science & Information Technology
2.
2.1 Physical Quantities
Physical properties such as height, volume, and
temperature that can be measured are called
Physical quantity. A number and a unit of
defined size is required to describe physical
quantity.
3.
The number alone doesn’t say much. If
you say an average textbook weighs 1.
The question would then be asked 1
what? 1 pound? 1 kilogram? 1 ounce?
You have to mention the unit of mass
along with the number for your statement
to be more meaningful.
4.
u Physical quantities can be measured in
many different units. For example, mass of
an object can be measured in pounds,
kilograms, ounce, and many other units.
u To avoid confusion, scientists around the
world have agreed to use a standard units,
known as International System of Unit
abbreviated as SI units, for some common
physical quantities.
5.
u In SI Unit, mass is measured in kilogram
(kg), length is measured in meters (m),
volume is measured in cubic meters (m3),
and time is measured in second (s).
u Many other widely used units are derived
from these SI units. For instance, unit of
speed is meters per second (m/s), unit of
density is grams per cubic centimeters (g/
cm3).
6.
2.2 Measuring Mass
u Mass is a measure of amount of matter in
an object.
u Weight is a measure of gravitational pull
on an object.
At the same location, two objects with
identical masses have identical weights;
that is gravity pulls them equally. Thus
mass of an object can be determined by
comparing the weight of the object to the
weight of a known reference standard.
7.
Two types of balances used for measuring
mass in the laboratory are shown below.
8.
2.3 Measuring Length and Volume
u Meter (m) is the standard measure of length
or distance in both SI and metric system.
One meter is 39.37 inches.
u Centimeter (cm; 1/100m) and millimeter
(mm; 1/1000m) are commonly used for
most measurements in chemistry and
medicine.
9.
u Volume is the amount of space occupied by an
object.
- The SI unit for volume is the cubic meter (m3).
u Liter (L) is commonly used in chemistry as a unit
of volume. 1L =0.001m3 = 1 dm3. One liter has
the volume a cube 10 cm (1dm) on edge. One
liter is further divided into 1000 milliliters (mL).
1 mL has the volume of a cube with 1 cm on
edge.
u 1 milliliter is often called 1 cubic centimeter (1
mL = 1 cm3).
10.
The relationship between metric units are
shown in Fig 2.2 below.
11.
2.4 Measurement and Significant Figures
Every experimental
measurement, no
matter how precise,
has a degree of
uncertainty to it
because there is a
limit to the number
of digits that can be
determined.
12.
u To indicate the precision of the
measurement, the value recorded should
use all the digits known with certainty,
plus one additional estimated digit that
usually considered uncertain by plus or
minus 1 or + 1.
u The total number of digits used to express
such a measurement is called the number
of significant figures. The quantity 65.07
g has four significant figures.
13.
Rules for determining significant figures
1. All non-zero digits are significant.
789 g has 3 significant figures.
2. Zeroes in the middle of a number are
significant.
69.08 g has four significant figures, 6, 9, 0,
and 8.
3. Zeroes at the beginning of a number are not
significant.
0.0089 g has two significant figure, 8 and 9.
14.
4. Zeroes at the end of a number and after the
decimal points are significant. 2.50 g has
three significant figures 2, 5, and 0.
25.00 m has four significant figures 2, 5, 0,
and 0.
5. Zeroes at the end of a number and before an
implied decimal point may or may not be
significant. 1500 kg may have two, three, or
four significant figures. Zeroes here may be
part of the measurements or for simply to
locate the unwritten decimal point.
15.
2.5 Scientific Notation
Scientific Notation is a convenient way to
write a very small or a very large number.
Written as a product of a number between 1
and 10, times the number 10 raised to
power.
M x 10n
16.
Two examples of converting standard numbers
to scientific notations are shown below.
17.
Examples of converting scientific notations
back to the standard numbers.
18.
2.6 Rounding off Numbers
Often calculator produces large number as a
result of a calculation although the number of
significant figures is good only to a fewer
number than the calculator has produced – in
this case the large number may be rounded off
to a smaller number keeping only significant
figures.
19.
Rules for Rounding off Numbers:
Rule 1 (For multiplication and division): The answer
can’t have more significant figures than either of the
original numbers.
20.
Rule 2 (For addition and subtraction): The
number can’t have more digits after the
decimal point than either of the original
numbers.
21.
2.7 Problem Solving: Converting a
Quantity from One Unit to Another
Factor-Label-Method: A quantity in one unit
is converted to an equivalent quantity in a
different unit by using a conversion factor that
expresses the relationship between units.
22.
When solving a problem, set up an equation so that
all unwanted units cancel, leaving only the desired
unit. For example, we want to find out how many
kilometers are there in 26.22 mile distance. We will
get the correct answer if we multiply 26.22 mi by the
conversion factor km/mi.
23.
2.9 Measuring Temperature
Temperature, the measure of how hot or cold an
object is, is commonly reported either in
Fahrenheit (oF) or Celsius (oC). The SI unit of
temperature is, however, is the Kelvin (K).
Temperature in K = Temperature in oC + 273.15
Temperature in oC = Temperature in K - 273.15
24.
Freezing point of H2O Boiling point of H2O
32oF 212oF
0oC 100oC
212 – 32 = 180oF covers the same range of
temperature as 100oC covers. Therefore,
Celsius degree is exactly 180/100 = 1.8 times
as large as Fahrenheit degree. Fig 2.4 gives a
comparison of all three scales.
25.
Fig 2.4 Comparison of the Fahrenheit, Celsius, and
Kelvin temperature scales
26.
Converting between Fahrenheit and Celsius
scales is similar to converting between
different units of length or volume, but is a
little more complex. The following formulas
can be used for the conversion:
oF = (9oF/5oC x oC) + 32oF
oC = 5oC/9oF x (oF – 32oF)
27.
2.10 Energy and Heat
Energy: Capacity to do work or supply energy.
Classification of Energy:
1. Potential Energy: stored energy.
Example: a coiled spring have potential
energy waiting to be released.
2. Kinetic Energy: energy of motion. Example,
when the spring uncoil potential energy is
converted to the kinetic energy.
28.
u In chemical reactions, the potential energy
is often converted into heat. Reaction
products have less potential energy than the
reactants – the products are more stable
than the reactants.
u Stable products have very little potential
energy remaining as a result have very little
tendency to undergo further reaction.
u SI unit of energy is Joules (J) and the
metric unit of energy is calorie (cal).
29.
One Calorie is the amount of heat necessary to
raise the temperature of 1 g of water by 1oC.
1000 cal = 1 kcal (kilocalorie)
1000 J = 1 kJ
1 cal = 4.184 J 1 kcal = 4.184 kJ
Not all substances have their temperature raised
to the same extent when equal amounts of heat
energy is applied.
30.
The amount of heat needed to raise the
temperature of 1 g of a substance by 1oC is
called the Specific Heat of the substance.
Unit of specific heat is cal/g .oC
It is possible to calculate how much heat must
be added or removed to accomplish a given
temperature change of a given mass of a
substance.
Specific Heat =
Calorie
Grams X oC
31.
2.11 Density
Density relates the mass of an object with its
volume. Density is usually expressed in units as -
Gram per cubic centimeter (g/cm3) for solids,
and Gram per milliliter (g/mL) for liquids.
Density =
Mass (g)
Volume (mL or cm3)
32.
2. 12 Specific Gravity
Specific Gravity (sp gr): density of a substance
divided by the density of water at the same
temperature. Specific Gravity is unitless. At normal
temperature, the density of water is close to 1 g/mL.
Thus, specific gravity of a substance at normal
temperature is equal to the density.
Density of substance (g/ml)
Density of water at the same temperature (g/ml)
Specific gravity =
33.
The specific gravity of a liquid
can be measured using an
i n s t r u m e n t c a l l e d a
hydrometer, which consists of
a weighted bulb on the end of
a calibrated glass tube, as
shown in the following Fig
2.6. The depth to which the
hydrometer sinks when placed
in a fluid indicates the fluid’s
specific gravity.
34.
Chapter Summary
u Physical quantity, a measurable properties, is
described by both a number and a unit.
u Mass, an amount of matter an object contains, is
measured in kilograms (kg) or grams (g).
u Volume is measured in cubic meters (m3) or in
liter (L) or milliliters (mL).
u Temperature is measured in Kelvin (K) in SI
system and in degrees Celsius (oC) in the metric
system.
35.
Chapter Summary Contd.
u Measurement of small or large numbers are
usually written in scientific notation, a product
of a number between 1 and 10 and a power of 10.
u A measurement in one unit can be converted to
another unit by multiplying by a conversion
factor.
u Energy: the capacity to supply heat or to do
work.
Potential energy – stored energy.
kinetic energy – energy of moving particles.
36.
Chapter Summary Contd.
u Heat: kinetic energy of moving particles in a
chemical reaction.
u Temperature: is a measure of how hot or cold an
object is.
u Specific heat: amount of heat necessary to raise
the temperature of 1 g of the substance by 1oC.
u Density: grams per milliliters for a liquid or
gram per cubic centimeter for a solid.
u Specific gravity: density of a liquid divided by
the density of water.
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