Download Unit 1: Atomic Structure & Electron Configuration

Unit 1: Atomic Structure
& Electron Configuration
I. Theories and Models
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Scientific Model – A pattern, plan,
representation or description designed to
show the structure or workings of an object,
system or concept.
A. Greeks
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400 B.C.
Democritus
particle theory- matter could not be divided into
smaller and smaller pieces forever, eventually the
smallest possible piece would be obtained and would
be indivisible.
called nature’s basic particle atomos-indivisible
no experimental evidence to support theory
B. John Dalton
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1808
English school teacher
Established first atomic theory:
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Matter is composed of atoms.
Atoms of a given element are identical to each other, but different
from other elements.
Atoms cannot be divided nor destroyed.
Atoms of different elements combine in simple whole-number ratios
to form compounds.
In chemical reactions, atoms are combined, separated or rearranged.
Model: tiny, hard, solid sphere
C. JJ Thomson
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1897
cathode ray tube experiment
given credit for discovering electrons,
resulting in the electrical nature of an atom
Plum pudding model – sea of positive charges
with negative charges embedded evenly
throughout.
D. Ernest Rutherford
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1911
Gold Foil (Alpha Scattering) Experiment
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Conclusions:
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atom is mostly empty space
most of mass of atom is in the nucleus
nucleus is positively charged
Model:
E. Niels Bohr
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1913
Rutherford’s student
electrons arranged in energy levels (orbits)
around the nucleus due to variation in
energies of electrons
higher energy electrons are farther from
nucleus
Planetary Model:
F. Quantum Model
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1924-current
Collaboration of many scientists
Better than Bohr’s model because it describes
the arrangement of e- in atoms other than H
Based on the probability (95% of time) of
finding and e- or an e- pair in a 3D region
around the nucleus known as an orbital
Model (on board)
II. General Structure of Atom
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nucleus
center of atom
p+ & n0 located here
positive charge
most of mass of atom, tiny
volume
very dense
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e- cloud
surrounds nucleus
e- located here
negative charge
most of volume of atom,
negligible mass
low density
III. Quantification of the Atom
A. Atomic Number - the number of p+ in nucleus
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All atoms of the same element have the same
atomic number.
Periodic table is arranged by increasing atomic
number.
if atom is electrically neutral, then the
#p+ = #e-
B. Mass Number - the total number of p+ & n0 in
nucleus of an atom.
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Round the atomic weight to a whole number
n0 = mass number - atomic number
C. Ions – atoms of an element with the same number of
p+ that have gained or lost e-, therefore having a – or
+ charge
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atoms form ions in order to be more stable like the noble gases
anion – ion with negative charge (gained e-)
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non-metal elements tend to form anions (ex. S2-)
change the end of the element name to –ide (sulfide ion)
cation – ion with a positive charge (lost e-)
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metal elements & H tend to form cations (ex. Sr2+)
Roman numerals may be used in the name of some metal ions that
can lose various numbers of e- (ex. Tin (IV) ion)
D. Isotopes – atoms of an element having the
same number of p+, but a different number of n0,
resulting in a different mass number.
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Two ways to represent isotope symbols:
mass #
atomic #
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14
6
C
or
C-14
mass #
Write mass # after the element name:
carbon-14
Isotopes of Hydrogen
Name
Symbol
e-
n0
p+
Mass #
Atomic #
1
0
1
1
1
Hydrogen-1 (protium)
1
1
Hydrogen-2 (deuterium)
2
1
H
1
1
1
2
1
Hydrogen-3 (tritium)
3
1
H
1
2
1
3
1
H
E. Average Atomic Mass – weighted average of
all natural isotopes of an element expressed
in amu* (atomic mass units).
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based on % abundance of isotopes
steps for calculating:
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change % to decimal
multiply decimal and mass number
add all results
place amu unit with answer
*amu=1/12 mass of C-12 isotope
IV. Electromagnetic Radiation
A. Properties
1. Form of energy which requires no substrate
to travel through.
2. Exhibits properties of a sine wave
amplitude
crest
line of origin
trough
wavelength (λ)
a. wavelength = distance between consecutive
crests (Greek letter lambda = λ)
b. frequency = # wave cycles passing a given
point over time (seconds); (Greek letter nu = ν )
*measured in Hertz (Hz)= 1/s, s-1, or
per second
c. all types of ER travel in a vacuum at the
speed of light (c) = 3.00 x 108 m/s
3. light equation
c=λν
* λ & ν are inversely (indirectly)
proportional (as one increases, the
other decreases)
λ
ν
*energy & ν are directly related (as one
increases/decreases, so does the other
*energy equation: E=hν
h= Plank’s constant = 6.63 x 10-34 J·s
V. Emission/Absorption Spectra
*The e- is the only SAP that absorbs/emits
energy.
A. Absorption Spectrum –when an e- absorbs
energy, it moves from the ground state (most
stable arrangement of e-) to an excited state
(which is not stable)
B. Emission Spectrum - when an e- emits energy, it
falls from the excited state back to ground state,
releasing energy in the form of electromagnetic
radiation, which may be visible
*unique to each atom
http://chemistry.bd.psu.edu/jircitano/periodic4.html
VI. Electron Configuration
A. Describes the arrangement of e- in an atom
1. each main energy level is divided into
sublevels
2. each sublevel is made up of orbitals, each of
which can hold up to 2 e*chart
Sublevel
# of
orbitals
s
1
p
3
d
5
f
7
shape
3. due to main energy levels getting closer
together, sublevels overlap
4. Aufbau principle – states that e- fill
orbitals of lower energy sublevels first
5. Abbreviated Configurations – use the
preceding noble gas symbol (in
brackets) to represent the filled inner
core of e-. Then write the remaining
configuration for the atom.
6. Orbital Configurations- arrangement of e- within
sublevels
2 rules determine arrangement:
a. Hund’s Rule – each orbital within
a sublevel receives 1 e- before it gets 2
* orbitals in the same energy sublevel are
degenerate (of equal energy)
b. Pauli Exclusion Principle – no 2 e- in an
orbital can have the same spin.
= clockwise spin
*exceptions
= counterclockwise spin