Download HL Chemistry: Notes Atomic Theory

HL Chemistry: Notes Atomic Theory
A. The Atom
1. Atoms composed of other parts protons, neutrons and electrons – smaller pieces
have been discovered - quarks, muon, gluon... but our study will be limited to p+, no,
e-.
2. protons - within nucleus, relative mass 1, identify atom, positive charge, atomic
number (Z)
3. neutrons - within nucleus, neutral charge, relative mass 1, equal to mass number
(A) - atomic # (Z)
4. electrons - outside nucleus, negative charge, relative mass 1/1840 (5 x 10-4),
determine reactivity of element.
anions: - ions , extra electrons, gained electrons
cations: + ions fewer electrons, lost electrons
5. atoms vs. molecules
a. atoms= smallest piece of element
b. molecule = smallest piece of compound
(covalent bonds= molecule) ionic bonds: not molecule – formula unit
6. isotopes and notation
a. isotope = atoms of same element with different masses
b. atomic mass gives a clue about the relative abundance of isotopes
C has a molar mass = 12.011 g and three isotopes are known to exist:
12
C, 13C, and 14C. Since the mole mass is close to 12 the 12C isotope is
more abundant than the other isotopes.
7. isotopes and their different physical properties (same chemical properties),
1
H,2H,3H are isotopes of hydrogen
8. averages of isotopes give atomic masses - non integer values
element “X” has 2 known isotopes
10.012 amu = 19.91 % and 11.009 amu = 80.09 % (assume 100 atoms)
(10.012amu x 19.91 + 11.009 amu x 80.09)/100 =
(1.99.3 amu + 881.7 amu)/100 =
10.810 amu = atomic mass
9. Periodic Table - metals, nonmetals, metalloids
Families: alkali metal, alkaline earths, halogens, noble gases.
(period ----- across table)
(family or group vertical columns)
B. The mass spectrometer
1. purpose -instrument used to compare masses of atoms
2. operation - atoms or molecules are passed into a beam of high speed electrons.
Electrons knock off electrons from atom (forming cations). An electric field
accelerates the ions into a magnetic field. Since a moving charged particle makes
its own magnetic field, the ion has an interaction with the applied electric field. The
path of the ion is altered. The amount of path deflection depends on its mass (heavier ions move less). Ions of different masses will separate and follow
different paths. The positions where they hit the detector give accurate values of
atomic masses.
3. parts of the spectrometer
a. ionizer : shoots electrons to form cations
b. accelerator : electrically charged plates repel then attract the cations
c. deflector: magnetic field separates ions by mass
d. detector: abundance and masses are shown
See the following we site for additional info:
http://www.chemguide.co.uk/analysis/masspec/howitworks.html
C. Electron Arrangement
We know the nucleus contains protons and neutrons which identify the atom, but
the electrons are the reason that atoms can and do combine in the ways we have
already discussed. We will now see how those electrons are arranged in atoms and
how this relates to chemical reactions.
1. Evidence for electron arrangement is the line spectra of different atoms
2. EM radiation
Since the electrons carry energy we need to know about the types of energy.
EM radiation is a way that energy travels.
The EM spectrum is a list of the types of EM radiation and their corresponding
wavelengths including: gamma rays, X-ray, UV, visible, IR, microwaves, and radio
waves.
Spectroscopes separate white light into its’ component colors to form the
continuous color spectrum (visible spectrum).
3. quantized energy
Energy occurs in discrete amounts (bundles) called quanta.
4. wave and particle behavior
Energy can be thought of as waves or particles - photons represent the dual nature
of light.
All matter can be discussed this way, but since large pieces of matter have such
short wavelengths and since very small photons have such small mass, it is easier to
discuss them as either matter or energy. However, electrons can be seen to exhibit
both wave and particle natures.
5. Atomic spectra of hydrogen - line vs. continuous
a. When atoms are Excited (given extra energy) they can release this energy in the
form of light. Electrons themselves absorb and then release energy which is shown
in the “emission spectra” of atoms.
b. The emission spectrum is a line spectra of colored light emitted by the excited
electrons of an atom.
c. A continuous spectrum is the pattern of light which is formed when white light is
separated into its component wavelengths.
d. A line spectra is a pattern of lines (either black lines on a continuous background
{absorption spectra}, or colored lines on a black background { emission spectra})
e. Hydrogen shows specific line emission spectra. Specific energies of light are
emitted, therefore energy levels are quantized (are of a specific amount of energy)
otherwise a continuous spectrum would be given for H emission spectra.
Refer Fig 7.6, p. 275
f. Note that the line spectra has converging lines (the lines are closer together as
they move farther from the nucleus). The changing type of EM radiation
determines the “jump” between energy levels. Transitions from higher to n=2 are
wavelengths of visible light. Transitions from higher to n= 1 are in the UV range.
Transition from higher to n=3 or 4 are in the IR range.
6. Bohr model of atoms other than H
a. Bohr’s was the “solar system” model of the atom. Electrons orbited the nucleus
like planets orbit the sun.
b. Bohr stated that electrons occupy energy levels in specific locations around the
nucleus of the atom.
c. The ground state is the location an unexcited electron occupies - it is the lowest
possible energy state for that electron.
d. Excited electrons (those that have gained energy) would occupy energy levels
higher (farther away) from the nucleus. These levels where labeled n =1, 2, 3,4...
(n is the principal quantum number)
e. When electrons move from a higher (excited) level to another lower level,
discrete (specific) wavelengths of energy are emitted. If these are in the visible
spectrum we see the associated color of light.
f. show energy level diagram and emissions... see Fig 7.11, p. 278
g. problems with Bohr model
Limited to H atom - doesn't work with >1 electron (orbits)
h. While Bohr’s research helped to develop our other models, we do not believe that
the solar system model is the correct picture of the atom and electron.
7. Quantum mechanical model of the atom - more complex than Bohr’s solar system
model
a. Quantum mechanics- electrons behave both like waves and particles
b. Orbitals - electrons are thought to occupy a space surrounding the atom,
different electrons will occupy different types of orbitals - based on a
mathematical theory of probabilities.
c. Heisenberg uncertainty principle - we cannot know the precise momentum and
position of an electron (a particle) at any given moment - since to find one aspect
we need to impinge on the other. (we cannot assume a specific path that each
electron in each orbital will take )
d. Probability distributions - instead of a given path we can discuss PROBABILITY
of finding an electron in a given space at a given time. An orbital is a region outside
of the nucleus where an electron can be found 90% of the time.
e. in the s orbital the shape is a sphere - ball . The probability of finding an
electron somewhere within that given shape is much higher than finding it
elsewhere.
f. quantum numbers
1. a series of numbers used to describe the location of every electron (where
electrons have the highest probabilities of being found )
2. principal quantum number = n = 1,2,3,,,,, only integers
describes the location in terms of distance from nucleus (energy level)
n=1 is closest to nucleus, n=2 is farther....
3. angular momentum number (l) integer from 0 to (n-1) for each value of n
l =0 means s orbital, l=1 means p orbital, l =2 means d orbital, l=3 means f
This quantum number define the region within the level (n) where an electron has a
high probability of being found - aka orbital
4. magnetic quantum number ml ranges from -l to l including zero. This relates the
orientation of the orbital in space - relative to other orbitals
5. the final quantum number is spin and this is + 1/ 2 or -1/2, this relates to the
direction the electron within the orbital moves. Each orbital holds 2 electrons
6. key to this is that no two electrons can have identical quantum numbers
(Pauli exclusion principle)
g. orbital shapes and energies
1. the second quantum number describes the orbital where e- are likely found
2. each orbital has a specific shape
3. s orbitals are spherical.
4. s orbital are 1s, 2s, 3s, 4s, 5s...
5. p orbitals begin with the second energy level
6. 3 types of p orbital exist px,py, and pz These correspond to the dimensions in
space. P orbitals are figure-8 or Barbell shapes along the x, y and z axis
7. d orbitals (begin with 3rd energy level) exist in 5 types -see figure 7.21 on p.
290
8. f orbitals exist in 7 different types (shapes not necessary)
h. writing electron configurations for atoms
1. Aufbau principle - a filling order exists for electrons in the space around the
Nucleus (e- fill orbitals of lowest energy first)
2. orbital diagrams - “picture” of electron arrangement in symbols
3. Hund`s rule - electrons will fill so that the maximum number of unpaired
electrons exists in the degenerate(same energy level) orbital -lowest energy
positions fill first
4. shorthand - kernel and valence electrons
Valence electrons – e- in the outermost energy level
Kernel – the rest of the atom
5. orbital filling across a period – know “diagonal rule” or be able to use Periodic
Table.
6. energy levels represented by large # (n)
7. orbitals represented by letters s,p,d,f
8. # of electrons per orbital represented by superscript number
Example:
Sulfur, 16S – 16 e- in total
e- configuration: 1s2 2s2 2p6 3s2 3p4
9. orbital notation – uses ‘boxes’ to represent each individual orbital (sub-orbital or
orientation) and arrows to represent e-. Each box may contain up to 2 e-, provided
they have opposite spin (arrow up and arrow down).