1. Lecture 2: Atoms, isotopes, ions & molecules
Topics Brown, chapter 2
1. Dalton’s Atomic Theory
Four postulates
2. Molecules & formulas
Molecular vs. empirical
3. Discovery of atomic structure
Subatomic particles
Radioactivity
Nuclear atom
4. Sub-atomic particles & the periodic table
Atomic number & mass
Isotopes
5. Atomic weights
Average atomic mass
Mass spectroscopy
6. Periodic table
7. Ions & ionic compounds
Ionic charges
8. Naming chemical compounds
Ionic compounds (salts), molecular compounds, & acids
“Grey” topics will be covered next week. Labor day has made this week bit short for a full chapter.

2. Ions are atoms that have lost or gained electrons.
Metals form cations.
Non-metals form anions.
Ionic compounds form via simultaneous ionization of metals & non-metals.
The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws.
It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made.
This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory.
While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

3. Ions and Charges
While the nuclei of atoms are pretty stable, atoms can easily gain or lose electrons, particularly from their outermost (or valence) shell.
Ions = atoms with missing or extra electrons
cations = atoms with missing electrons and a net positive charge; metals
anions = atoms with extra electrons and a net negative charge; non-metals
11p+
lose 1 electron
10e-
Na+ ion
Na atom
11e-
11p+
Let’s look at how different ions are from the atoms they are derived from: SODIUM example
Loss of an electron leaves results in an ‘extra’ or ‘excess’ proton, so the ion of sodium has a positive charge of +1.
p.54

4. Ions “collaborate” to form salts
Atoms interact to increase their own stability & to achieve full valence shells
1 ve-
1 valence electron
7 valence electrons
K atom
Br atom
Let’s look at how different ions are from the atoms they are derived from: SODIUM example

5. Electrostatic attraction pulls the cation & anion towards one another
Ions “collaborate” to form salts
Ions
8 valence electrons
8 valence electrons
K+1 ion
Br-1 ion
Let’s look at how different ions are from the atoms they are derived from: SODIUM example
Electrostatic attraction pulls the cation & anion towards one another

6. The ionic “compound” KBr
Ions “collaborate” to form salts
Ions
K+1 ion
Br-1 ion
Let’s look at how different ions are from the atoms they are derived from: SODIUM example
The ionic “compound” KBr

7. How different are atoms from ions?
Chemical and physical properties of ions are VERY different from the atoms they are derived from.
Sodium atoms lose an entire shell of electrons, the 3s valence shell, when they ionize. The remaining electrons are closer to the nucleus and held very tightly, so the ion is LESS chemically reactive than the atom (since electrons are responsible for chemical change).
PROPERTY
Na ATOM
Na+1 ION
Charge
Neutral +1
e- configuration
1s22s22p63s1
1s22s22p6
Pure form
Soft, shiny metal
Not found alone
Rxn with H2O
Violent, exothermic
Simply dissolves
Rxn with H2 gas
Makes NaH
No rxn
Rxn with O2 gas
Makes Na2O2
Rxn with alcohol
Makes alkoxide
Conductivity
Heat & electricity
Poor unless in H2O
Let’s look at how different ions are from the atoms they are derived from: SODIUM example

8. Ionic Compounds How and why do ionic compounds, or “salts”, form?
It’s basically a story of greed, desire & ability. All atoms want to have 8 electrons in their outermost (or valence) shell. Metals typically have less than 4 atoms in this shell, & so lose these e- to get down to the next full shell. Non-metals usually have more than 4 ve- & thus gain e- to fill the outermost shell.
1. Transfer of electrons from Na to Cl produces charged ions Na+ and Cl-.
2. Ions are attracted by opposite charges and bind together.
Notice that the ions are attracted by electrostatic force & that this causes them to form a lattice in which each + ion is surrounded by + ions, & vice versa.
The size of the crystal lattice varies with number of ions present.
What “law” dictates that the formation of oppositely charged ions is mutually “codependent”?
The Law of Conservation of Mass says that e- lost by one atom must be gained by another; formation of anions & cations is complementary.
p.56

9. Knowing Ionic Charge
Use group numbers to predict ionic charge as shown below. Transition metals have varying charges & cannot be predicted this way.
For now, you may find it useful to mark up your periodic table so that you’ll know ionic charges of elements in these compounds.
Transition metals are the beige squares in the middle (“Midwest”) of the periodic table. -3 -1 +1 +3 -2 +2
p.55