Download Atomic Structure Test Review

Elements and Periodic
Trends
Chemistry
Mr. Treanor
Atomic Theories
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Democritus – first known theory of
atoms: “Matter is composed of very
small solid particles that are
indivisible.”
Dalton’s Theory – five points in his
theory explain:
– the law of conservation of mass
– the law of definite proportions
– the law of and multiple proportions
More Atomic Theories
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Experiments with cathode-ray
tubes showed that atoms were
composed of even smaller particles.
Millikan discovered the mass and
charge of the electron.
Rutherford’s gold foil experiment
led to the discovery of the nucleus.
These all led to the Bohr Model.
Quantum Model
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Electrons exist in orbitals, not orbits.
Orbitals are regions around the nucleus
where there is a 90% probability (as they
are usually shown) that electrons will be
found.
Heisenberg Uncertainty Principle: it is
impossible to know the speed and location
of an electron at the same time
Schrödinger Wave Equation: mathematical
equations for the shapes of orbitals
Shapes of orbitals
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l = 0: these are s orbitals (spherical)
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l = 1: these are p orbitals (pair)
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l = 2: these are d orbitals (drops/donut)
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l = 3: these are f orbitals (funky)
The f orbitals (funky shapes)
Electron Configurations
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Aufbau principle: electrons go
into the lowest energy orbital
that has room
Pauli exclusion principle: no two
electrons in the same atom have
the same set of quantum
numbers
Hund’s rule: electrons fill up
empty orbitals first
More about
configurations
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The maximum number of electrons in each
energy level is found using the quantum
number n – main energy level
The number of electrons can be up to 2n2
This gives 2,3,18, and 32 for n= 1,2,3, and 4
The highest occupied level is the main
energy level with the highest n
Inner-shell electrons are the ones that are
not in the highest occupied level
Electron Configuration Notation
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Used to show the locations of all electrons
This is the order electrons fill the orbitals:
1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d,7p
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The small numbers are the number of electrons in that
particular orbital
Example: Na is 1s22s22p63s1 – There are 2 electrons in
the 1s orbital, 2 in the 2s orbital, 6 in the 2p orbital, and
1 in the 3s orbital
The highest occupied level is 3
The inner shell electrons are in 1s,2s, and 2p
How to tell which element?
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If the outer electrons are only s, it is an sblock element. One s electron means group
1; 2 is 2.
If the outer electrons have p orbitals, it is a
p-block element. 1 p electron is group 13, 2
is 14, and so on until the p orbitals are full
with 6. These are group 18. The entire
energy level is full. That is why noble gases
are unreactive.
If there are s and d, it is a d-block metal.
Brief History
Mendeleev invented the periodic table
and arranged it by atomic mass and
group properties.
 Moseley, working in Rutherford’s lab
and playing with x-rays, arranged the
periodic table by atomic number.
 They did not know about the noble
gases yet.
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A look at the blocks
•s-block – metals that easily lose electrons
•d-block – metals that lose different numbers of
electrons depending on the other element
•p-block – mostly nonmetals that gain electrons
Periodic Group Properties
Group 1
Alkali metals, most reactive metals, only 1
outer s electron
Group 2
Alkaline earth metals, 2 outer s electrons
(no other outer electrons)
Transition metals, the ones we call metals
(gold, copper, iron, silver)
Mostly nonmetals & semiconductors
Halogens, most reactive nonmetals, 5
outer p electrons
Noble gases, all main energy levels full
D-block
P-block
Group 17
Group 18
Atomic Radius
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Gets bigger as you go down a group
(more shells)
Gets smaller as you go across a
period (more protons with same
number of shells)
Measured as half the distance
between 2 atoms in a pure sample of
a substance
Ionization Energy (IE or IP)
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The energy required to remove an electron from a neutral
atom is the 1st ionization energy, or IE1.
The energy required to remove a second electron after the
first ionization is the 2nd ionization energy, or IE2. There is
also an IE3, IE4, and so on.
Electron Affinity (EA)
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The ability of an atom to gain electrons
Negative numbers mean more energy
Electronegativity
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It is defined as the relative ability of an atom in a molecule
to attract electrons towards itself.
The atom with the higher electronegativity will dominate
the electrons.
Electronegativity (EN) Values
What causes
Electronegativity?
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Protons in the nucleus are attracted to
valence electrons in OTHER ATOMS
If there are a lot of electrons in the
way, the electronegativity is lower. This
happens when there are more energy
levels (shells)
If there are more protons, they will be a
stronger pull and higher
electronegativity. This happens when
the atomic number is higher.
Atomic Radii
1. Which atoms have less shells? lower
2. Of those, which ones have more protons? To the right
3. Explain why this agrees with the EN values. The
elements with less shells and more protons have the
highest electronegativity values.
The Tug-O-War
How electronegativity works:
The atom with the highest EN is
the strongest and wins.
Valence Electrons
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Outer s and p electrons
For group 1, it is 1; for group 2 it is 2.
For groups 13 – 18, it is the group number -10