# Importance of Scientific Measurements in Chemistry

An essential
element of all sciences is obtaining proper measurements. The International
System of Units, known as SI Units, was developed by scientists to

standardize
measurements across all sciences. Even with a standardized system, though,
there is plenty of uncertainty that can come into play. The uncertainty must be
minimized to ensure proper understanding of a process or experiment.**SI Units**

Scientific
measurements quantify the magnitude of something, described with a unit. For
example, length is quantified in meters. However, because there are many
different units -- e.g., inches, feet, centimeters -- scientists developed SI
units to avoid confusion. Using common units, scientists from different
countries and cultures can easily interpret each others' results. SI units
include meters (m) for length, liters (L) for volume, kilograms (kg) for mass,
seconds (s) for time, Kelvin (K) for temperature, ampere (A) for electrical
current, mole (mol) for amount and candela (cd) for luminous intensity.

**Accuracy and Precision**

When taking
scientific measurements, it is important to be both accurate and precise.
Accuracy represents how close a measurement comes to a true value. This is
important because bad equipment, bad data processing or human error can lead to
inaccurate results, meaning they are not very close to the truth. Precision is
how close a series of measurements of the same thing are to each other.
Measurements that are not precise do not properly identify random errors and
can yield a widespread result that is not helpful.

**Significant Figures**

Measurements
can be only as accurate as the limitations of the measuring instrument allow.
For example, a ruler marked in millimeters can be accurate only up to the
millimeter, because that is the smallest unit available. When a measurement is
made, its accuracy must be preserved. This is achieved through
"significant figures." The significant figures in a measurement are
all the digits that are known for certain, plus the first one that is uncertain.
For example, a meter stick delineated in millimeters can measure something to
be accurate to the fourth decimal place. If the measurement were 0.4325 meters,
there would be four significant figures.

**Significant Figures Limits**

Any non-zero
digit in a measurement is a significant figure. Zeroes that occur before a
decimal point are also significant. Also, zeroes that appear after a non-zero
digit in a decimal value are considered significant. Whole number values, such
as a full known quantity like five apples, is considered to have no impact on
the significant digits of a calculation.

**Multiplying and Dividing Significant Figures**

When
multiplying or dividing measurements, count the significant figures in the
numbers. Your answer should have the same number of significant figures as the
original number with the lowest number of significant digits. For example, the
answer to the problem 2.43 x 9.4 = 22.842 should be converted to 23, rounding
up from the partial number.

**Adding and Subtracting Significant Figures**

When adding or
subtracting measurements, the number of significant figures is determined by
the placement of the largest uncertain digit. For example, the answer to the
problem 212.7 + 23.84565 + 1.08 = 237.62565 should be converted to 237.6,
because the largest uncertain digit is the .7 in the tenths place in 212.7. No
rounding should take place because the 2 that follows the .6 is smaller than 5.

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