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Chemistry Lecture Notes — Atomic Theory, the Mole, and Nomenclature
Chemistry Lecture Notes
Dalton’s Atomic Theory and Chemical Laws
Historical development
Ancient Greek philosophers proposed that matter was made of tiny, indivisible particles called atomos.
Aristotle later described matter as a combination of four elements: earth, air, fire, and water.
In 1807, John Dalton proposed a modern atomic theory based on experimental measurements.
Dalton’s five postulates
Matter is composed of extremely small particles called atoms.
Each element contains one type of atom with a characteristic mass.
Atoms of different elements have different properties and masses.
Compounds form when atoms of different elements combine in small, whole-number ratios.
Atoms are not created or destroyed during a chemical change; they are rearranged.
The fifth postulate connects to the law of conservation of matter. Although later discoveries showed that atoms contain smaller particles, Dalton’s theory established the foundation for modern chemistry.
Law of definite proportions
A pure compound always contains the same elements in the same proportion by mass, regardless of the sample’s source or amount. For example, pure water always has a fixed hydrogen-to-oxygen mass ratio.
Law of multiple proportions
If two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are related by small whole-number ratios. This law supports the idea that atoms combine in discrete units.
Subatomic Particles, Isotopes, and Atomic Symbols
Subatomic particles
Atoms contain:
Protons: positively charged particles in the nucleus.
Neutrons: neutral particles in the nucleus.
Electrons: negatively charged particles surrounding the nucleus.
Isotopes
Isotopes are atoms of the same element that have:
the same number of protons, but
different numbers of neutrons.
Therefore, isotopes have the same atomic number but different mass numbers.
Hydrogen has three important isotopes:
Protium: no neutrons
Deuterium: one neutron
Tritium: two neutrons; radioactive
Magnesium also occurs as isotopes such as magnesium-24, magnesium-25, and magnesium-26.
Atomic notation
An atom or ion can be represented as:
\[
{}^{A}_{Z}X^{q}
\]
where:
\(X\) is the element symbol.
\(Z\) is the atomic number, equal to the number of protons.
\(A\) is the mass number, equal to protons plus neutrons.
\(q\) is the charge, shown when the particle is an ion.
Useful relationships:
\[
\text{neutrons} = A - Z
\]
For a neutral atom:
\[
\text{electrons} = \text{protons}
\]
For an ion, the number of electrons changes while the number of protons remains the same.
Atomic mass and mass spectrometry
The atomic mass shown on the periodic table is a weighted average of the masses of an element’s naturally occurring isotopes. Mass spectrometry helps determine isotope masses and abundances by separating charged particles according to their mass-to-charge ratios.
Chemical Formulas and the Mole Concept
Types of chemical formulas
Molecular formula: gives the actual number of atoms of each element in a molecule. Example: benzene is \(C{6}H{6}\).
Empirical formula: gives the simplest whole-number ratio of the elements. Benzene’s empirical formula is \(CH\).
Structural formula: shows how atoms are connected to one another.
Isomers
Isomers are compounds with the same molecular formula but different arrangements of atoms. Acetic acid and methyl formate are examples: both have the formula \(C{2}H{4}O_{2}\), but their structures differ.
The mole
A mole is a counting unit used for atoms, molecules, ions, and other particles.
\[
1\text{ mol} = 6.022 \times 10^{23}\text{ particles}
\]
This number is called Avogadro’s number.
Molar mass
Molar mass is the mass of one mole of a substance, measured in \(g/mol\). Its numerical value matches the substance’s atomic mass or formula mass in atomic mass units.
The central conversion pathway
Chemistry calculations commonly move between three quantities:
\[
\text{grams} \leftrightarrow \text{moles} \leftrightarrow \text{particles}
\]
Use molar mass to convert between grams and moles, and Avogadro’s number to convert between moles and particles. Dimensional analysis helps keep the units organized and ensures that unwanted units cancel.
The Periodic Table
Organization
The periodic table is arranged by increasing atomic number.
Periods are the horizontal rows.
Groups are the vertical columns.
Elements in the same group often have similar chemical properties because they have related valence-electron arrangements.
Major classifications
Metals: generally shiny, malleable, and good conductors; found mainly on the left and in the center.
Nonmetals: generally duller and poorer conductors; found mainly on the upper right.
Metalloids: have intermediate properties and lie along the staircase boundary.
Important group names
Group 1: alkali metals
Group 2: alkaline earth metals
Group 15: pnictogens
Group 16: chalcogens
Group 17: halogens
Group 18: noble gases
Chemical Nomenclature
Ionic compounds
Name the cation first and the anion second.
For a monatomic anion, change the element ending to -ide. For example, sodium and chloride form sodium chloride.
Transition metals
Transition metals can form ions with different charges. A Roman numeral in parentheses identifies the metal’s charge.
Example:
\[
FeCl_{2} = \text{iron(II) chloride}
\]
Hydrates
Hydrates are ionic compounds that contain water molecules bound within their crystal structure. Name them using:
the ionic compound’s name,
a numerical prefix, and
the word hydrate.
Example: copper(II) sulfate pentahydrate.
Molecular compounds
Use prefixes to show how many atoms of each element are present.
Common prefixes include:
mono- = one
di- = two
tri- = three
tetra- = four
penta- = five
Examples:
\(CO\): carbon monoxide
\(N{2}O{4}\): dinitrogen tetroxide
The prefix mono- is not used for the first element in a molecular compound name.
Acids
Binary acids
Binary acids contain hydrogen and one nonmetal. Their names follow this pattern:
\[
\text{hydro-} + \text{root} + \text{-ic acid}
\]
Example: hydrochloric acid.
Oxyacids
Oxyacids contain hydrogen and an oxygen-containing polyatomic ion.
An ion ending in -ate becomes an acid ending in -ic.
An ion ending in -ite becomes an acid ending in -ous.
For example, nitrate becomes nitric acid.
High-yield connections
Atomic theory explains chemical formulas. Atoms combine in whole-number ratios, which is why compounds have definite formulas.
Isotopes explain average atomic masses. The periodic table reports a weighted average rather than the mass of only one isotope.
The mole connects microscopic and macroscopic chemistry. It lets you convert between individual particles and measurable masses in grams.
The periodic table supports nomenclature. Element location and classification help identify likely ion charges and naming patterns.
Dimensional analysis prevents conversion errors. Write units at every step and cancel them systematically.
Quick review questions
What are the five main ideas in Dalton’s atomic theory?
How do the laws of definite proportions and multiple proportions differ?
What is the difference between atomic number and mass number?
How are isotopes of the same element alike and different?
What does the atomic mass on the periodic table represent?
How do molecular, empirical, and structural formulas differ?
What is Avogadro’s number, and why is it useful?
What is the conversion pathway between grams, moles, and particles?
How are metals, nonmetals, and metalloids distinguished?
When are Roman numerals used in ionic-compound names?
How do the endings -ate and -ite affect oxyacid names?
One-minute summary
Matter is composed of atoms, and chemical compounds form when atoms combine in fixed whole-number ratios. Isotopes differ in their neutron counts, while atomic number identifies the element through its proton count. The mole and Avogadro’s number connect individual particles to measurable amounts. The periodic table organizes elements by atomic number and properties, while nomenclature rules allow chemical compounds and acids to be named systematically.