Download File - Miss Cummings

Standards: SCSh1-9
SC3 Students will use the modern atomic theory to explain the characteristics of atoms.
a. Discriminate between the relative size, charge, and position of protons, neutrons, and electrons in the atom.
b. Use the orbital configuration of neutral atoms to explain its effect on the atom’s chemical properties.
c. Explain the relationship of the proton number to the element’s identity.
d. Explain the relationship of isotopes to the relative abundance of atoms of a particular element.
f. Relate light emission and the movement of electrons to element identification.
Unit Three: Atomic Theory
Theory- explanation of observable facts and phenomena
Development of the Atomic Theory
Democritus believed that matter could not be continually cut and called the tiniest piece the “atom,”
which means, “can not be cut.”
Aristotle rejected Democritus’ idea of the atom and said that matter could be continually cut.
Dalton proposed that atoms are the building blocks of matter, are indivisible, are identical for the same
element and different for different elements, and that atoms unite to form compounds.
Thomson is credited with the discovery of the electron, which was a blow to Dalton’s indivisible atom.
He proposed the “plum pudding” model of the atom: negatively charged electrons are embedded in a ball
of positive charges.
Rutherford concluded that most of the atom is empty space with a dense positively charged core, after
he completed the “gold foil experiment,” in which he aimed particles at gold foil and most passed straight
through, while only a few were deflected. He proposed the “planetary model” of the atom: electrons orbit
the positively charged nucleus like planets orbit the sun.
Bohr believed that electrons have a certain amount of energy and are divided into different
levels/orbitals. Typically, electrons are found in the ground level state, the lowest energy level they can
occupy (closer to the nucleus). When electrons absorb energy they become excited, move to higher levels
(farther from the nucleus), are unstable, and soon release photons (light) as they fall back down to the
ground state.
The Modern View is based off of the “charged cloud” model. Although we still can’t see exactly where
electrons are, this model suggests that electrons are confined to electron clouds or regions of space
where electrons are likely to found based off of probability.
Structure of the Atom
Particle
Proton
Electron
Charge
+
-
Mass*
1 amu
0**
Location
Nucleus
Electron Cloud
Neutron
0
1 amu
Nucleus
Calculation
= atomic number
= atomic number
in neutral atoms
= mass number –
atomic number
*amu: atomic mass unit
**2,000 electrons = 1 proton (the mass is so small that it doesn’t even factor in)
Miss Cummings – Chemistry
1
Atomic Mass vs. Mass Number
Mass Number – total number of protons and neutrons in an atom; always a whole number because you
can’t have part of a proton or part of a neutron (can round the atomic mass to get the mass number)
Atomic Mass- average of the masses of all the element’s isotopes; typically printed on the periodic table
and not a whole number
Isotope- atoms of elements that have different numbers of neutrons; typically written in hyphen
notation (Element – Mass Number)
Example:
Isotopes of Carbon
Carbon - 12
Has six protons and six neutrons
Abundance = 98.9%
Carbon - 13
Has six protons and seven neutrons
Abundance = 1.1%
Atomic Mass = 12 x (0.989) + 13 (0.011) = 12.011 amu
Quantum Mechanics
Video: https://www.youtube.com/watch?v=7u_UQG1La1o
Albert Einstein said that light has a dual nature, it may behave as a wave or as a stream of particles called
quanta or photons. (Think Schrodinger’s Cat and Super Positioning)
Spectroscopy- spectral lines (similar to a prism) are unique for each element and represent the energy
released as an electron returns to a lower energy state
Quantum Numbers
letters (n, l, m, and s) that represent numbers and are used to describe an electron in an atom
“Never Leave Me Sweetie”
 n: Energy Level
o There are at least seven energy levels (1, 2, 3, 4, 5, 6, and 7), with #1 being closest to
the nucleus
o Maximum number of electrons in an energy level = 2n2
Example: 5th Energy Level = [2 x 52] = 50 electrons

l: Sublevel
o s, p, d, and f : “Some People Don’t Fart”
o Number of sublevels = number of levels
 level 1 = s
 level 2 = sp
 level 3 = spd
 level 4 = spdf

m: Number of Orbitals in Each Sublevel
o s = 1 orbital
o p = 3 orbitals
o d = 5 orbitals
o f = 7 orbitals
*skips every other number
Miss Cummings – Chemistry
2

s: Number of Electrons in Each Orbital
o each orbital can hold only two electrons that have opposite spins
o s = 1 orbital = 2 electrons
o p = 3 orbitals = 6 electrons
o d = 5 orbitals = 10 electrons
o f = 7 orbitals = 14 electrons
* # of orbitals x 2 = # of electrons
Hund’s: “empty bus seat rule,” each orbital in a sublevel is occupied by one electron before any
orbital is occupied by a second electron
Pauli’s: two electrons in the same orbital must have opposite spins
Aufbau: “diagonal rule,” there is a specific order in which electrons fill energy levels/sublevels
2
l
2
6
2
6
2
6
2
6
2
6
m
s
10
10
14
14
10
RULES
2
6
10
14
10
14
Miss Cummings – Chemistry
3