Warm-up
For sulfur:
1. How many valence electrons does it
have?
2. What ion does this typically form?
3. Write the electron configuration for the
ion.
Nucleus
• Contains 99.9% of the mass of an atom
• Found at the center of the atom
• Diameter is only 1/10,000 of the entire atom
Protons and Neutrons
• Proton – positively charged subatomic particle, number
determines element’s identity (atomic number)
• Neutron – neutral subatomic particle
• Protons and neutrons have almost the same mass
Electron
• Negatively charged subatomic particle
• Electrons have almost no mass
• Move around the nucleus in a “cloud”
Changes to atoms
• Ions – different number of protons and electrons
• Calculate the difference in charge
• Example: Chloride ion (Cl-) has 17 protons (+) and 18 electrons (-)
• Cations are positively charged
• Anions are negatively charged
• Isotopes – different number of protons and neutrons
• Add number of protons and neutrons to find the mass number of
that specific isotope
• Example: Carbon-14 has 6 protons and 8 neutrons
• Atomic mass on the periodic table is an average of all of the
isotopes of that element that exist
*Worksheet*
History of Atomic Theory
• http://hi.fi.tripod.com/timeline/timeline.htm
Figure 11.17: The Bohr model of the hydrogen atom.
Copyright© by Houghton Mifflin Company. All rights reserved.
7
Figure 11.5: Electromagnetic radiation.
Copyright© by Houghton Mifflin Company. All rights reserved.
8
Figure 11.6: Photons of red and blue light.
Copyright© by Houghton Mifflin Company. All rights reserved.
9
Figure 11.10: An excited H atom returns to a lower energy level.
Copyright© by Houghton Mifflin Company. All rights reserved.
10
Figure 11.11: Colors and wavelengths of photons in the visible region.
Because each element has a different electron
configuration, each has a unique emission spectrum
Copyright© by Houghton Mifflin Company. All rights reserved.
11
Determining electron location
• Principle energy levels
• Represented by the letter “n”
• A whole number between 1 and 7
• As n increases, energy increases and the electron is farther away
from the nucleus
• Sublevel
• Shape of electron cloud
• s = spherical
• p = dumbbell
•
•
•
•
1st energy level has 1 sublevel -- s
2nd energy level has 2 sublevels – s, p
3rd energy level has 3 sublevels -- s, p, d
4th energy level has 4 sublevels -- s, p, d, f
Determining electron location
• Orbitals
• Describes the orientation in space within a sublevel
•
•
•
•
s = 1 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals
Electron Configuration Rules
• Aufbau Principle - electrons enter orbitals of lowest energy
first.
• Pauli Exclusion Principle - an atomic orbital can hold a
maximum of 2 electrons and those 2 electrons must have
opposite spins
• Hund’s Rule - When electrons occupy orbitals of equal energy
(same sublevel), one electron enters each orbital with parallel
spin before pairing oppositely.
Valence Electrons
• Electrons in the outermost occupied principle energy level of
an atom
• Electrons that are NOT in the outermost principle energy level
are known as “core electrons”
Periodic trends
• Atomic size/radius
- Tends to increase down a column (group)
- Tends to decrease across a row (period)
• Ionization energy - energy required to remove an electron
from an individual atom (gas)
– Tends to decrease down a column (group)
– Tends to increase across a row (period)
• Electronegativity - tendency for an atom to attract electrons
to itself when bonded to another element
– Tends to decrease down a column (group)
– Tends to increase across a row (period)
Periodic table
• Write principle energy level for each period (1-7)
• Label the representative groups (1-8)
• For each group list:
•
•
•
•
•
•
Metal, nonmetal, or metalloid?
Special name (if it has one, such as alkali metals)
Number of valence electrons
General valence electron configuration (s1, s2p3)
Do the elements typically form a cation or anion?
What is the usual charge on the ion?