Download 4. Atomic Structure

Atomic Structure
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Subatomic Particles
• An atom is the smallest unit of an
element. It consists of three major
particles:
• Note: amu = atomic mass unit
Particle
Mass
Location
Charge
Proton
1 amu
Nucleus
+1
Neutron
1 amu
Nucleus
0
Electron
1/1836 amu Orbitals
-1
Atomic Theory
explains the structure of matter in
terms of different combinations of
very small particles called atoms.
Dalton’s Theory
• All elements are composed of
indivisible atoms.
• All atoms of a given element are
identical.
Atoms of different elements are
different; that is, they have different
masses.
Compounds are formed by the
combination of atoms of different
elements.
JJ Thomson
• used a cathode ray tube to show
smaller units that make up an atom.
The ray was deflected a certain way
by a magnetic field, so he concluded
that the ray was formed by particles
and that the particles were negatively
charged. The only source available
for the particles was the atoms
present.
• Therefore, Thomson theorized that an
atom contains small, negatively
charged particles called electrons.
This theory is referred to as the Plum
Pudding Model. In this model, the
mass of the rest of the atom was
evenly distributed and positively
charged, taking up all of the space
not occupied by the electrons.
Rutherford’s Gold Foil Experiment
• Alpha particles directed at a piece
of gold foil.
• Some alpha particles went
through, some were deflected,
some were returned.
Conclusion:
Majority of the volume of an atom is
empty space.
Atoms have a dense positively
charged central core.
Do Now
• How many p+, n0, e- are in:
35
17 Cl
• Draw the model (where are the
+
0
p , n , e ) *look at periodic
table
Quick Review
• Who are the 3 theorists we went
over yesterday? What are the
“nicknames” of their theories and
why?
The Bohr Model
•
•
•
•
1 or K-shell = max 2 e2 or L-shell = max 8 e3 or M-shell = max 18 e4 or N-shell = max 32 eNiels Bohr
Also called the Planetary Model
Use Periodic Table to determine how electrons are
arranged.
Basic formula for determining the amount
of electrons a shell can hold is: 2n2 where
n is the principle energy level (level 4 =
2(42) = 32 electrons.
The Orbital Model
• Electron Cloud model
• Wave Mechanical Model
• *Important Definitions:
–Principle Energy Level: Region
around the nucleus (the dense
positively charged central core
of an atom) in which electrons
can be found. (The closer to the
nucleus, the lower the energy).
–Quanta: Small amount of energy
that a(n) electron can absorb or
release as it moves through
principle energy levels.
–Ground State: all electrons fill
lowest energy levels before higher
energy levels begin to fill.
–Excited State: one or more
electrons fills a higher energy
level before the lower ones are
filled.
–Spectral Lines: As electrons at
higher energy levels (excited
electrons) fall back to their normal
energy levels (ground state) they
release energy in the form of the
spectrum. ROYGBIV
Electron Configurations
• Looking at the periodic table of elements,
you will notice numbers at the bottom of
each element. These numbers represent
the electron configuration of the element
(the address of the electrons). The
Period represents the number of
principle energy levels (orbitals) present.
The Group represents the sublevel for
each principle energy level.
Principle Energy Levels
•
•
•
•
•
•
Period 1 = 1 shell
Period 2 = 2 shells
Period 3 = 3 shells
Period 4 = 4 shells
Period 5 = 5 shells
Period 6 = 6 shells
SPDF Sublevels
• Group 1 & Group 2 = ‘s’ sublevel (max – 2
electrons)
• Group 13 – Group 18 = ‘p’ sublevel (max – 6
electrons)
• Group 3 – Group 12 = ‘d’ sublevel (max – 10
electrons)
• Lanthanum & Actinum Series = ‘f’ sublevel
(max – 14 e)
• See examples
Valence Electrons
• Electrons that fill the outermost
principle energy level of an atom.
• Ex: Mg 2-8-2 has 2 valence
electrons
• Ex: Ne 2-8 has 8 valence electrons
• Valence electrons are largely
responsible for an element’s
chemical and physical properties.
Electron Dot Diagrams
• The term Kernel refers to all of the
non-valence electrons and the
nucleus of an atom.
• The Kernel is represented by the
element’s symbol, valence
electrons are represented by dots.
• Example: Oxygen
• Atomic Number: the number of
Protons in the nucleus of an atom.
(Atom by definition is an electrically
neutral particle, so this must also be
equal to the number of Electrons).
• Mass Number: the number of
protons plus neutrons.
• Question: Why are there fractional
mass numbers (decimals) on the
periodic table? (ex: Na, O2, …)
Answer: Due to the existence of
isotopes.
• Note: Atomic Symbols: One or two
letters, 1st is always capital, the 2nd
is always lower case.
• Isotope: Atoms if the same element
having the same number of protons,
but different number of neutrons
• Example: Isotopes of Hydrogen
Particle
Protons Neutrons Mass #
Symbol
Protium
1
0
1
1 H
1
Deuterium
1
1
2
2 H
1
Tritium
1
2
3
3 H
1
Calculating Isotopes (weighted
atomic mass)
1) Take the percent of each isotope and
convert it back to a decimal ( 100)
2) Multiply the decimal by the mass number
3) Add the numbers together to get the
Weighted Atomic Mass
C-12 99%
( .99 x 12) +
C-14 1%
(.01 x 14) = 12.02
Mg-26 1.75 % Mg-24 98.25%
Cl-35
75%
Cl-37 25%
Ions
• Atoms of the same elements
having the same number of
protons, but different number of
electrons
Ca+ions: positive charge formed
by losing electrons
Anions: negative charge formed
by gaining electrons