unit 1: atomic structure & electron configuration
TRANSCRIPT
Unit 1: Atomic Structure & Electron Configuration
I. Theories and Models Scientific Model – A pattern, plan,
representation or description designed to show the structure or workings of an object, system or concept.
A. Greeks • 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• 1808• English school teacher• Established first atomic theory:
1. Matter is composed of atoms.2. Atoms of a given element are identical to each other, but different
from other elements.3. Atoms cannot be divided nor destroyed.4. Atoms of different elements combine in simple whole-number ratios
to form compounds.5. In chemical reactions, atoms are combined, separated or rearranged.
• Model: tiny, hard, solid sphere
C. JJ Thomson• 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• 1911• Gold Foil (Alpha Scattering) Experiment
• Conclusions: atom is mostly empty space most of mass of atom is in the nucleus nucleus is positively charged
• Model:
E. Niels Bohr• 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• 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
nucleus center of atom p+ & n0 located here positive charge most of mass of atom, tiny
volume very dense
e- cloud surrounds nucleus e- located here negative charge most of volume of atom,
negligible mass low density
III. Quantification of the AtomA. Atomic Number - the number of p+ in nucleus
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.
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
atoms form ions in order to be more stable like the noble gases anion – ion with negative charge (gained e-)
• 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-) • 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.
• Two ways to represent isotope symbols:
or C-14
• Write mass # after the element name:
carbon-14
C146
mass #
atomic #mass #
Isotopes of Hydrogen
Name Symbol e- n0 p+ Mass # Atomic #
Hydrogen-1 (protium) 1 0 1 1 1
Hydrogen-2 (deuterium) 1 1 1 2 1
Hydrogen-3 (tritium) 1 2 1 3 1
H11
H21
H31
E. Average Atomic Mass – weighted average of all natural isotopes of an element expressed in amu* (atomic mass units).
based on % abundance of isotopes steps for calculating:
1. change % to decimal
2. multiply decimal and mass number
3. add all results
4. place amu unit with answer
*amu=1/12 mass of C-12 isotope
IV. Electromagnetic RadiationA. Properties
1. Form of energy which requires no substrate to travel through.
2. Exhibits properties of a sine wave
amplitude wavelength (λ)crest
trough
line of origin
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 ConfigurationA. 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
shape
s 1
p 3
d 5
f 7
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 = counterclockwise spin
*exceptions