Honors Chemistry
Mr. Fedell
Chapter 5 & 6 Notes – Electrons in Atoms & The Periodic Table
(Student edition)
Chapter 5 problem set:
Chapter 6 problem set:
5.1
27, 31, 34-37, 50, and 57
24, 32, 35, 43, 45, and 49
Models of the Atom
The Development of Atomic Models
BIG PROBLEM:
**Rutherford’s atomic model could not explain the chemical
of elements.**
BIG QUESTION:
Why do elements produce different colors when heated?
(a.k.a why do metals turn red when hot)
BIG SOLUTION:
In 1913, Neils Bohr (Danish), stated that electrons could occupy fixed
without giving off energy.
NEED TO KNOW Conclusions:
-
energy, thus…
energy
Electrons FARTHER from the nucleus have
electrons CLOSER to the nucleus have
Diagram and staircase analogy:
This theory was deduced from
and
.
Describe what happens in the flame test:
______________________________________________________________________________
______________________________________________________________________________
1
Draw a gas tube:
5.3
Physics and the Quantum Mechanical Model
So… Why DO elements produce
when
?
NEED TO KNOW: Bohr’s answer
**Since the amount of energy an electron has determines the distance it orbits from the
nucleus, the electrons:**
1.
energy (heat, electricity, or light) and move to higher energy levels
(
from the nucleus).
2.
the extra energy as
and
to a
energy level.
WE SAY:
“Electrons take in energy, jump up, fall back down, and release energy in the form of light.”
Examples: ___________________________________________________________________.
5.1 Continued . . . Models of the Atom
Eventually, we learn that the Bohr model is a lie.
Bohr model
really is
wave mechanical model
Quantum mechanics (equations) allows scientists to determine the probability of
finding particles in certain places. This leads to the quantum (wave) mechanical model.
This model does not show the path of electrons - just the most
location
2
6.1 Organizing the Elements and Classifying the Elements
NEED TO KNOW HISTORY:
Origin of the periodic table
Dimitri Mendeleev - publish first real periodic table in 1869
- based on
- listed elements in order of
- left spaces for
Property
Eka - Aluminum
Atomic Mass
Density
Melting Point
Oxide Formula
68 amu
5.9 g/cm3
low
E2O3
Ga (1875)
NEED TO KNOW SKILL:
Predicting properties using other elements data:
Example: Predict the density of Aluminum given:
Density: Ga = 5.9g/cm3 & B = 2.3 g/cm3
**FIND THE AVERAGE!!!
NEED TO KNOW Modern Periodic Law:
** Properties of elements are a periodic function of atomic number. **
NEED TO KNOW INFO: Reading the periodic table:
Periods:
left to right on the periodic table
elements have the same # of
elements
have similar properties
a.k.a rows, shells, and energy levels
Groups:
up and down on the periodic table
elements have the same # valence electrons
elements have similar properties
a.k.a families (columns)
3
NEED TO KNOW CLASSIFICATION:
General Drawing of the Periodic Table:
Metals, Metalloids, Nonmetals)
NEED TO KNOW FAMILIES:
Metalloids - B, Si, As, Te, At, Ge, Sb
Group 18 – Noble Gases – Very UNreactive
5.2 & 6.3 Electron Arrangement in Atoms & Periodic Trends
NEED TO KNOW PERIODIC TREND:
Valence Electrons – electrons that occupy the
H
Be


1 2
Transition Metals
(usually 2 VE)
B

3
shell
C

4
N

5
O

6
He
F


7 8
NEED TO KNOW PERIODIC TREND:
Electronegativity – the ability of an atom to attract electrons
(
has highest value)
Increasing trend looks like….
4
Organizing Electrons on the PT: By Row and Sublevel
NEED TO KNOW SKILL: Writing Orbital Notation
Type
S
P
D
F
number of sublevels
__
__
__
__
total # of electrons
2
6
10
14
Notation for orbital notation
Orbital Notation: a representation of the
Example:
_
1s
example - oxygen 
1s
1=

2s
of each electron
=
s=
 
2p
NEED TO KNOW SKILL: Writing Electron Configuration
Notation for electron configurations
Electron Configuration: a representation of the arrangement of electrons in an atom.
Example: 2p2
2 = Row
p = sublevel
2
= # electrons in sublevel
example - oxygen 1s22s22p4
5
Electron Notation: Electrons enter orbitals in a set pattern. For the most part, they follow the
Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule:
NEED TO KNOW PRINCIPLES:
Principle: electrons must fill
energy
levels before entering higher levels.
Example: Why we fill s-orbitals (low energy) before p-orbitals (high energy)
Pauli Exclusion Principle - electrons occupying the same orbital must have opposite spin.
Example: Why we fill orbitals with an UP-arrow and DOWN-arrow
Hund’s rule ( better known as the
): before any second electron
can be placed in a sub level, all the orbitals of that sub level must contain at least one electron.
Example: Why we do one UP-arrow in each orbital before placing a DOWN-arrow
Practice:
Orbital Notation: Draw the first 18
H
He
Electron Configuration: Draw the first 18
H
He
Li
Be
Li
Be
B
C
B
C
N
N
O
O
F
F
Ne
Ne
Na
Na
Mg
Mg
Al
Al
Si
Si
P
P
S
S
Cl
Cl
Ar
Ar
6