Download topic 03 outline YT 2010 test

Honors Chemistry - Topic 3 Test (Sections II and III) Outline
Atoms - The Building Blocks of Matter
II. Bohr Model of The Atom
pp.
303-312
A. The Nature of Light [5.2a, 4.1a end “Bohr”; 5.2b]
1. Types of spectra
 Continuous spectrum – entire or most of
electromagnetic spectrum
 Line Spectrum – pieces/little bits of color seen for
element
2. Concepts of light
a. light as a wave
 See SG from short quiz
b. light as a particle - photons
 Atoms that have gained extra energy release that
energy in the form of light.
Line spectrum: very specific wavelengths of light that
atoms give off or gain
o
Each element has its own line spectrum, which can be
used to identify that element.
o
The line spectrum must be related to energy transitions in
the atom
o
•The atom is quantized, i.e. only certain energies are
allowed.
c. electromagnetic spectrum
3. Quantizing line spectra: Rydberg Equation
o used in atomic physics to describe the wavelengths of spectral
lines of many chemical elements.

B. The Bohr Atomic Model
1. Explaining the existence of line spectra
 The Bohr model was based on a simple postulate, Bohr
applied to the hydrogen atom the concept that the
electron can exist only in certain energy levels without
an energy change but that, when the electron changes
its state, it must absorb or emit the exact amount of
energy that will bring it from the initial state to the
final state.
 (The ground state is the lowest energy state available to
the electron. The excuted state is any level higher than
the ground state. )
 The formula for a change in energy (∆E) is:
o ∆Eelectron = Efinal – Einitial
2. Bohr’s Electron configurations
 Energy of atom is related to distance of electron from nucleus
3. Bohr’s Energy-level diagram and quantum numbers
o Energy of the atom is quantized
 Atom can only have certain specific energy states
called quantum levels or energy levels.
 When atom gains energy, electron “moves” to a
higher quantum level
 When atom loses energy, electron “moves” to a
lower energy level
 Lines in spectrum correspond to the
difference in energy between levels
o
 Ground state: minimum energy of an atom
 Therefore electrons do not crash into the
nucleus
 The ground state of hydrogen corresponds to having
its one electron in the n=1 level
 Excited states: energy levels higher than the
ground state
o Distances between energy levels decrease as the energy
increases
 1st energy level can hold 2e-1, the 2nd 8e-1, the 3rd
18e-1, etc.
 Further from nucleus = more space = less repulsion
o Valence shell: the highest-energy occupied ground state
orbit
Problems with Bohr Model
o Only explains hydrogen atom spectrum (and other 1-electron
systems)
o Neglects interactions between electrons
o Assumes circular or elliptical orbits for electrons (which is not
true)
III. Modern Atomic Structure – Wave Mechanical Model
Wave mechanical - a new model of hydrogen atom that seemed
to apply equally well to all other atoms, which Bohr’s model
failed to do.
A. Changing Bohr’s Model [4.1b]
p.312-318; 319327; 352
 Orbital Theory – Electrons are found in orbitals around the
nucleus. The orbitals describe regions of space where an
electron is most likely to be found at any given moment.
o s-sublevel – shape of sphere around nucleus. Electron is
located somewhere within this space.
 Quantum Mechanical Model (aka Charge Cloud Model) - treats
electrons as waves and uses wave mathematics to calculate
probability densities of finding the electron in a particular
region in the atom
o –Schrödinger Wave Equation
o –Can only be solved for simple systems, but
approximated for others

1. Duality of matter: DeBroglie (Wave Mechanical Model of the Atom)
o Suggested matter consist of waves
o Experiments later showed that electrons could be treated as
waves
o Just as light energy could be treated as particles
2. Heisenberg’s Uncertainty Principle
 Heisenberg- Says you can’t know where electron is and its is
going at the same time (b/c any effort to determine one is going
to change it).
o Heisenberg Uncertainty Principle – can’t actually
describe where an electron is in an atom. All you
can do is define an area of space around the
nucleus where you are most likely to find an
electron (Led to orbital theory)
3. Matter-Waves: Schrödinger’s wave concept
 Schrödinger – Came up with equation that describes the motion
of a single electron around a single proton of hydrogen
B. Describing The Modern Atom
As we mentioned earlier, each principal energy level, n, has n sublevels. This
means the first has one sublevel, the second has two, the third has three,
etc. The sublevels are named s, p, d, and f.
Energy level principal quantum
Number of
Names of
number, n
sublevels
sublevels
1
1
s
2
2
s, p
3
3
s, p, d
4
4
s, p, d, f
At each additional sublevel, the number of available orbitals is increased by
two: s = 1, p = 3, d = 5, f = 7, and as we stated above, each orbital can hold
only two electrons, which must be of opposite spin. So s holds 2, p holds 6
(2 electrons times the number of orbitals, which for the p sublevel is equal
to 3), d holds 10, and f holds 14.
GOOD TABLE BELOW
Sublevel
Number of orbitals
Maximum number of
electrons
Quantum number, l
s
1
2
p
3
6
d
5
10
f
7
14
0
1
2
3
1. Electron configurations [5.3a] – pg.75
It is important to remember that, when there is more than
one orbital at a particular energy level, such as three p
orbitals or five d orbitals, only one electron will fill each
orbital until each had one electron.
 (Each electron has a charge of -1)
 Each period represents the principle energy level
 Principle Energy levels are made of Sublevels
o There are 4 basic sublevel types
 s-sublevel – can hold up to 2 electrons
 p-sublevel - can hold up to 6 electrons
 d-sublevel – can hold up to 10 electrons
 f-sublevel - can hold up to 14 electrons
o (# sublevels = principle energy level #)
a. Aufbau Principle – the principle that an electron occupies
the lowest energy orbital that can receive it.
b. Hund’s Rule(of Maxium Muliplicity) – principle that, after
this(above/Augbau), pairing will occur with the addition of
one more electron to each orbital.
c. electron blocks and the periodic table
 s, p, d, and f blocks
 Box Diagrams – show from most general (energy levels) to
most specific (orbitals)
o each sublevel is made of orbitals
o 1 box=1 orbital
o Orbital – region of space where you are likely to
find 2 electrons, which have opposite spins
 s-sublevel – 1 orbital (box)
 p-sublevel - 3 orbital (boxes)
 d-sublevel – 5 orbitals (boxes)
 f-sublevel – 7 orbitals (boxes)
o Spins are represented by arrows (up and down) – up first,
followed by down
This matches the size of the s-block.
The p subshell can hold up to six electrons. This matches
the size of the p-block.
The d subshell can hold up to 10 electrons. This matches
the size of the d-block with 10 columns.
The f subshell can hold up to 14 electrons. The 14 columns
of the f-block match the filling of the f subshell.

d. Lewis dot structures [5.3b] –Page 78
 In 1916, G.N. Lewis devised the electron dot notation , which
may be used in place of the electron configuration notation. The
electron dot noation shows only the chemical symbol
surrounded by dots to represent the electrons in the incomplete
outer level (valence electrons).
 Lewis Dot Structures – 2 dots on each side of letter
(maximum amount is 8)
o Valence electrons – electrons in outermost energy level
(most possible is 8 = stable octet)
 Family 1A 1 valence electron
 Family 2A - 2 valence electrons
 Family 3A - 3 valence electrons
 Family 4A - 4 valence electrons
 Family 5A- 5 valence electrons
 Family 6A- 6 valence electrons
 Family 7A - 7 valence electrons
 Family 8A – 8 valence electrons
o Examples – # dots=# of valence electrons
 Sodium (3s) has 1 valence electron and dot diagram
would have a single dot
 Magnesium (3s2) has 2 valence electrons and has
two dots
 Aluminum has 3 valence electrons has 3 dots
 Tin (Sn) has 4 valence electrons and 4 dots
 Nitrogen has 5 valence electrons and 5 dots
 Oxygen has 6 valence electrons and 6 dots
 Bromine has 7 valence electrons and 7 dots
 Argon has 8 valence electrons and 8 dots (stable
octet of valence electrons)
2. Quantum numbers [5.4a] (Electron Zipcode/General Location
is space)
 There are 4 quantum numbers that recognize a specific electron
in orbital notation (See more info below)

1. n  Ex. 1st principle energy level has a quantum number
of 1, 2nd PEL has quantum number of 2, 3, 4, etc.
2. ℓ - Sublevel (ℓ) – different sublevels are given dif.
quantum #’s
ℓ Sublevel
0 s
1 p
2 d
3 f
3. m ℓ - Orbital (m ℓ) – which orbital the electron is located
in
 0 is central orbital
4. ms - Spin (ms) – Box diagram: Up arrow (+ ½), down
arrow (- ½)

You can name an electron with 4 quantum numbers, and find a
specific electron if given the 4 numbers.
n
The principal quantum number – only know this name
l
The subsidiary or azimuthal or angular momentum or orbital
shape quantum number
ml The magnetic quantum number
ms The electron spin quantum number
3. Atomic Orbitals
a. shapes of orbitals = s, p, d and f
 Another method is commonly used to designate
sublevels. Instead of giving the quantum number ℓ, we
use a letter (s,p, d, or f) to indicate the sublevel. A
sublevel for which ℓ=0 is refered to as an s sublevel. If ℓ=1,
we are dealing with a p sublevel. A d sublevel is one for which
ℓ=2, in an f sublevel, ℓ=3
o A sublevel __________
b. Pauli Exclusion Principle
 States that in a given atom no two electrons can have
the same set of four quantum numbers (n, l, m1 and
ms)
o No orbital may have more than 2 electrons.
o Electrons in the same orbital must have opposite spins.
o s sublevel holds 2 electrons (1 orbital)
o p sublevel holds 6 electrons (3 orbitals)
o d sublevel holds 10 electrons (5 orbitals)
o f sublevel holds 14 electrons (7 orbitals)
4. Excited and ground state [5.4b; 5.4c]
 The ground state is the lowest energy state available to
the electron.
 The excited state is any level higher than the ground
state.
o Ex. 1s22s233s1 is excited state of Boron b/c skipped over 2p
o If matches periodic table, it is the ground state
o If electrons are in a level they are occupied, if they reach
maximum amount, then they are
full
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