Some Basic Concepts of Chemistry

Lesson

Part 1 — how chemists measure, why matter combines in fixed ratios, and how the mole lets us count atoms by weighing.

In this lesson you will

Use SI units, scientific notation and significant figures when reporting measurements
State and test the laws of chemical combination against experimental data
Explain Dalton’s atomic theory, atomic mass and molecular mass
Convert between mass, amount in moles and number of particles

Table of very large and very small quantities met in chemistry
Quantity	Typical size
Atoms in a grain of salt	about $10^{18}$ formula units
Mass of one carbon atom	$2.0 \times 10^{-23}$ g
Atoms in 12 g of carbon-12	$6.022 \times 10^{23}$
Diameter of an atom	about $10^{-10}$ m

Concept

Measurement and SI units

A measurement is a number with a unit. Chemists use the SI system: seven base units, with prefixes to scale them. Volume is often given in litres: $1\ \text{L} = 1\ \text{dm}^3 = 10^{-3}\ \text{m}^3$ .

Table of SI base units used in chemistry and common prefixes
Quantity	SI unit	Symbol
Length	metre	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol
Prefixes	milli, micro, nano, pico	$10^{-3}, 10^{-6}, 10^{-9}, 10^{-12}$

Concept

Scientific notation, accuracy and precision

Write numbers as $N \times 10^{n}$ with $1 \le N < 10$ . Precision is closeness of repeated readings to each other; accuracy is closeness to the true value.

Table comparing three sets of readings for precision and accuracy
Readings of a 10.00 g standard	Precise?	Accurate?
10.01, 10.02, 10.01	Yes	Yes
9.52, 9.51, 9.52	Yes	No
9.8, 10.3, 10.1	No	Roughly on average

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