Atomic Theory and Structure

Observations Supporting Atomic Theory

Mass is neither created nor destroyed in chemical reactions.
Example:
- Aqueous solutions of mercury nitrate and potassium iodide react to form a precipitate of mercury iodide and aqueous potassium iodide.
- HgI2(s) + 2KNO3(aq)
- 4.55g + 2.02g = 6.57g
Known amounts of solid KI and solid are weighed and dissolved in water.
The solutions are mixed to give solid which is removed by filtration.
The solution that remains is evaporated to give solid.
On weighing, the combined masses of the products equal the combined masses of the reactants.

Different samples of a pure chemical substance always contain the same proportion of elements by mass.
Example: Water
- By mass, water is 88.8% oxygen and 11.2% hydrogen.

Elements can combine in different ways to form different substances.
The mass ratios are small whole-number multiples of each other.
Examples:
- Nitrogen monoxide: 7 grams nitrogen per 8 grams oxygen
- Nitrogen dioxide: 7 grams nitrogen per 16 grams oxygen

Elements are made up of tiny particles called atoms.
Each element is characterized by the mass of its atoms.
Atoms of the same element have the same mass, but atoms of different elements have different masses.
The chemical combination of elements to make different chemical compounds occurs when atoms join in small whole-number ratios.
Chemical reactions only rearrange how atoms are combined in chemical compounds; the atoms themselves don’t change.

J. J. Thomson (1856–1940) proposed that cathode rays consist of tiny, negatively charged particles called electrons.
The electron beam ordinarily travels in a straight line.
The beam is deflected by either a magnetic field or an electric field.

Ernest Rutherford (1871–1937) bombarded gold foil with alpha particles.
Most alpha particles passed through the foil undeflected.
Approximately 1 in every 20,000 particles was deflected.
A fraction of those particles were deflected back at an extreme angle.
Rutherford proposed that the atom must consist mainly of empty space, with the mass concentrated in a tiny central core—the nucleus.

Particle	Mass (Grams)	Mass (u*)	Charge (Coulombs)	Charge (e)
Electron	9.109382 \times 10^{-28}	5.485799 \times 10^{-4}	-1.602176 \times 10^{-19}	-1
Proton	1.672622 \times 10^{-24}	1.007276	+1.602176 \times 10^{-19}	+1
Neutron	1.674927 \times 10^{-24}	1.008665	0	0

Atomic Number (Z): Number of protons in an atom’s nucleus, equivalent to the number of electrons around an atom’s nucleus.
Mass Number (A): The sum of the number of protons and the number of neutrons in an atom’s nucleus.
Isotope: Atoms with identical atomic numbers but different mass numbers.

Lithium-6: 3 protons, 3 electrons, 3 neutrons
Lithium-7: 3 protons, 3 electrons, 4 neutrons
Chlorine-37: atomic # = 17, mass # = 37, # of protons = 17, # of electrons = 17, # of neutrons = 20
Isotope symbol for Chlorine-37: ^{37}_{17}Cl
Another way to represent isotopes: 37Cl