Unit 1 - Principles of Chemistry: Chapter 3: Atomic Structure

ATOMIC STRUCTURE

This chapter explores atoms and their differences. The 118 elements are the building blocks of everything.

LEARNING OBJECTIVES

• Define atom and molecule.

• Describe atom structure: positions, relative masses, and charges of sub-atomic particles.

• Define atomic number, mass number, isotopes, and relative atomic mass ($A_r$).

• Calculate element's relative atomic mass ($A_r$) from isotopic abundances.

ATOMS AND MOLECULES

• Atoms join to make molecules.

• Molecule: two or more atoms chemically bonded (covalent bonds).

• Atoms in a molecule can be from the same or different elements.

• Hydrogen ($H_2$) molecule: 2 hydrogen atoms bonded.

• Water ($H_2O$) molecule: 2 hydrogen atoms and 1 oxygen atom bonded.

THE STRUCTURE OF THE ATOM

• Atoms are made of protons, neutrons, and electrons (sub-atomic particles).

• Nucleus: protons and neutrons.

• Electrons are far from the nucleus.

• Table 3.1 shows properties of protons, neutrons, and electrons.

Table 3.1 The properties of protons, neutrons and electrons

• Atom's mass is mostly in the nucleus because electrons are much lighter.

• Masses and charges are relative because actual values are very small.

• Example: 6 x $10^{23}$ protons weigh 1 g.

ATOMIC NUMBER AND MASS NUMBER

• Atomic number: number of protons in an atom's nucleus.

• Each element has a different number of protons.

• Example: 8 protons = oxygen atom.

• atomic number = number of protons

• Mass number: total number of protons and neutrons in the nucleus.

• mass number = number of protons + number of neutrons

• Notation: $_{27}^{59}Co$ (27 protons, 32 neutrons).

• number of neutrons = mass number - atomic number

ISOTOPES

• Number of neutrons can vary.

• Isotopes: atoms of the same element with the same atomic number but different mass numbers.

• Same number of protons, different numbers of neutrons.

• Varying neutrons don't affect chemical reactions.

• Chemical properties are controlled by electrons, which are identical for isotopes.

RELATIVE ATOMIC MASS

• Relative atomic mass: average mass of an atom, considering isotope abundance.

• Naturally occurring chlorine has more 35Cl than 37Cl (relative atomic mass is closer to 35).

• Calculate relative atomic mass using isotopic abundances to find the average mass of an atom.

EXAMPLE 1

Boron: 20% $^{10}B$ and 80% $^{11}B$. Calculate the relative atomic mass.
Imagine 100 atoms: 20 have mass 10, 80 have mass 11.
Total mass = (20 x 10) + (80 x 11).
Average mass = \frac{20 \imes 10 + 80 \imes 11}{100} = 10.8
Relative atomic mass of boron = 10.8.

COUNTING THE NUMBER OF ELECTRONS IN AN ATOM

Atoms are electrically neutral.
Proton charge (+1) = electron charge (-1).
number of electrons = number of protons
Oxygen (atomic number 8): 8 protons, 8 electrons.
Key: atomic number (from Periodic Table).
Atomic number is defined by protons because electron number changes in reactions.

THE PERIODIC TABLE

Atoms are arranged by increasing atomic number.
Use the Periodic Table to find protons, neutrons,