atomic structure & the periodic table - mount st. mary's...atomic structure & the...
TRANSCRIPT
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Atomic Structure &
The Periodic Table
The Greek Philosophers
• Democritus – believed that all matter is made up of tiny particles that could not be divided
• Aristotle -- thought that matter was made of only four elements
• Fire
• Earth
• Air
• Water
Dalton’s Atomic Theory (1808)
• All elements are composed of atoms
• All atoms of the same element have the same mass, and atoms of different elements have different masses*
• Compounds contain atoms of more than one element
• In a particular compound, atoms of different elements always combine in the same way
J. J. Thomson
• Provided the first evidence that atoms are made up of even smaller particles.
• “Plum Pudding” model – electrons were
evenly spread out among a positively
charged mass (chocolate chip ice
cream)
Thomson’s Experiment
• Hypothesis – The beam was a stream of charged particles that caused the air to glow.
• Experiment – Positive and negative plates were put on either side of the tube.
• Results – The beam was bent towards the positive plate.
• Conclusion – There are negative particles being given off of the atoms that are attracted to the positive plate. (Atoms are made of smaller pieces!)
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Ernest Rutherford’s Gold Foil Experiment
Hypothesis: that alpha particles would pass straight through a thin sheet of gold
Results
The alpha particles did not pass straight through, but instead were deflected in
various directions.
Conclusion
The positive charge of an atom IS NOT evenly spread out, but concentrated in
a very small area (nucleus).
Subatomic Particles• Proton (p+)
• Positive charge
• Found in nucleus
• Neutron (n)• Neutral (no charge)
• Found in nucleus
• Approximately the same mass as a p+
• Electron (e-)• Negative charge
• Found orbiting around the nucleus
• MUCH smaller than p+ or n (1/1836)
Atomic Number (Z)
• The number of protons found as a part of the nucleus in an atom
• Unique for each element
• Cannot change without changing the
type of element
Mass Number
• The total number of particles that make up the nucleus
• Mass number = p+ + n
• n = mass number – p+
• Atoms of the same element may have
varying mass numbers because the
number of neutrons can vary
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Isotopes
• Atoms of the same element with
different mass numbers (numbers of neutrons)
Ways to write isotopes
• Element – mass number(Carbon – 12)
• Symbol – mass number
(C – 12)
• Mass number Symbol (12 C)
•Ways to write isotopes, cont.
C)( SymbolNumber MassNumberAtomic
126
Wave Nature of Energy
Electromagnetic (EM) Radiation
• any form of energy that radiates in all
directions from a single source
• consists of oscillating electric and
magnetic fields that carry energy
EM Spectrum• discovered by James Maxwell
• c = speed of light (true for ALL EM radiation)
• c = 3.0 x 108 m/s
• = frequency
• the number of waves passing a given point during a unit of time
• units are “per seconds” = s-1 = Hertz (Hz)
• = wavelength (m)
wave formula c =
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(2.901 m)
What is the wavelength of a radio
wave broadcasted at 103.4 MHz?
• heated solid objects emit visible light where intensity and color depend on temperature
• “black” because there is no light emitted before heating
blackbody radiation
• as the temperature of a black body increases
• decreases
• increases
• E increases
• cannot be explained by classical physics
Wien-Planck law
• a fixed amount of energy
• the smallest amount of energy that can be emitted or absorbed as EM radiation
• E = h
• h = Planck’s constant = 6.626 x 10-34 Js
• E = energy (Joules, J)
• one quanta = h , two quanta = 2 h
quantum (Max Planck 1858-1947)
• energy packet that behaves like a particle
photon
• clean metal surfaces exposed to light will emit electrons
• emitted electron = photoelectron
photoelectric effect
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3.37 x 10-19 J/photon
What is the energy of one photon of
yellow light ( = 589 nm)?
• extended Planck’s quantum theory to say that
energy has mass!
• E = mc2
• E = energy of a photon (J)
• m = mass of photon (kg)
• c = speed of light
Einstein
• ties together Planck’s and Einstein’s work
E = mc2 = h
m = hv/c2
(c = )
m = h/ c
Compton’s Theory
4.79 x 10-19 kg
What is the mass of a photon of
violet light ( = 415 nm)?
• m = h/ v
• m = mass (kg)
• h = Planck’s constant
• = wavelength (m)
• v = velocity (m/s)
• If mass is large, wavelength is small
• Newtonian mechanics applies
• If mass is small, wavelength is large
• Quantum mechanics applies
De Broglie (Louis) equation
8.18 x 10-34 m
Calculate the wavelength of a 2.53 g
bullet traveling 320. m/s.
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• Planck and Maxwell work toward wave theory
• Einstein work toward particle theory
• EM has characteristics of both a wave and a particle at all times
Wave-Particle duality Niels Bohr
• 1st postulate – atom has only
certain allowable (quantized) energy states
• depend upon energy level occupied by e-
• Energy levels identified with integers (n=1,
2 …)
E = -kZ2/n2
k = 2.179 x 10-18 J (Rydberg constant)
Z = atomic number
n = principal quantum number (energy level)
energy of an electron in energy
level n
• 2nd postulate – atom doesn’t radiate energy in one of it’s energy states
• 3rd postulate – atom changes energy states by absorbing or emitting photons of specific frequencies
• if a photon is absorbed, an electron is promoted to a higher energy level (called an excited state) E > 0
• When an electron “relaxes” back to its ground state, a photon is emitted E < 0
2f
2i
18
n
1
n
110x179.2E
2.09 x 10-18 J
Calculate the energy absorbed by
an electron jumping from the 1st
energy level to the 5th.
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-5.448 x 10-19J
Calculate the energy of an electron
moving around a H atom in the 2nd
energy level.
Energy Levels
• The possible energy an electron
can have
• Similar to steps
• Higher steps have more energy
• Going down a step means energy was released
Energy Levels, cont.
• Seven possible energy levels that
correspond to the rows on the periodic table
• Also called “shells”
Bohr’s Model of the Atom
• Each circle represents an energy
level
• 1st level: 2 electrons
• 2nd level: 8 electrons
• 3rd level: 18 electrons
• 4th level: 32 electrons
Erwin Schrödinger
• Developed mathematical model
to describe the motion of electrons
• Work leads to electron cloud
model
Orbital
• Region of space where an electron
is likely to be found
• Shapes are created by taking
many pictures of the electron’s
location over a period of time
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Pauli Exclusion Principle
• Only two electrons will occupy the
same orbital
• A series of the same type of orbital
is also called a sub-level or sub-
shell.
s-orbital
• Shaped like a sphere
• Only one type (b/c it can only point one direction!)
Electron Probablity Map – p orbital p - orbital
• Has one shape that can point
three ways
• Total of 3 types of p-orbitals (px,
py, pz)
Electron Probability Maps
d orbitalsd - orbital
• Five total types
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f - orbitals
• Seven possible types
•How Orbitals Interact
http://micro.magnet.fsu.edu/electromag/java/atomicorbitals/index.html
Electron Configurations(More sophisticated than Bohr’s!)
• Know the total number of electrons
in the atom!
• Use the diagonal rule
Aufbau Principle
• Electrons fill orbitals that have the
lowest energy first
• Follow the DIAGONAL RULE!
•The Diagonal Rule
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
Exceptions to the Diagonal Rule
Expected Found
Cr [Ar]4s23d4 [Ar]4s13d5
Mo [Kr]5s24d4 [Ar]5s14d5
Cu [Ar]4s23d9 [Ar]4s13d10
Ag [Ar]5s24d9 [Ar]5s14d10
Au [Ar]6s25d9 [Ar]6s15d10
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Orbital Diagrams
Shows the placement of electrons in the orbitals…
Ex. Sodium
1s22s22p63s1
1s 2s 2p 3s
Hund’s Rule
Electrons will fill each orbital on a sublevel before pairing with each other.
Ex. Carbon
• Diamagnetic
• all electrons are paired
• s2, p6, d10, f14
• Paramagnetic
• One or more electrons is unpaired
Quantum Numbers
Principal Quantum Number (n)
n = energy level (1, 2, 3, …)
Azimuthal Quantum Number (ℓ)
ℓ = type of orbital (s=0, p=1, d=2, f=3)
Magnetic Quantum Number(which orbital)
m = from - ℓ to + ℓ
Electron Spin (first or second electron)
s = ½
What are the quantum numbers for the 3rd
and 6th electrons in neon?
N 1s22s22p6
3rd electron n=2, ℓ=0, m=0, s=+½
6th electron n=2, ℓ=1, m=-1, s=-½
1s 2s 2p
Electron Configurations on the Periodic Table
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Shorthand electron configurations
Use Noble Gas shortcuts ONLY!
Kr (36 e-)
1s22s22p63s23p64s23d104p6
Mo (42 e-)
1s22s22p63s23p64s23d104p65s24d4
[Kr] 5s24d4
Isoelectronic series
• Species that have the same electron configuration
• E.g. O2-, F-, Ne, Na+, Mg2+
John Newlands
• Arranged the 16 known
elements in order of increasing
atomic mass
Dmitri Mendeleev
• Arranged the periodic table by
increasing atomic mass and stacking elements with similar
properties in the same column
• Was able to predict the properties of elements not yet discovered by
looking at the blank spaces in his table
Henry Moseley
• Fine-tuned the periodic table
by placing the elements in
increasing atomic number
(protons were not known during
Mendeleev’s time)
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Organizing the Periodic Table
• Metals are left of the stair-step
line
• Non-metals are to the right of
the stair-step line
• Metalloids (have properties of
both) are found adjacent to
stair-step line
Other Periodic Arrangements •Atomic Informationfrom the Periodic Table
1
1.00794
H
Atomic Number
Atomic Mass
Element Symbol
Atomic Mass
• The weighted average of all
possible isotopes for an element
• Atomic Mass =
(%A)(Mass of A) + (%B)(Mass of B)
Period
• A row in the periodic table
corresponding to the energy
level on which the electrons
exist
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Group
• A column on the periodic table
in which the elements have
similar properties
• Similar properties are created
by similar electron
configurations
1A
2A 3A 4A 5A 6A 7A
8A• A Groups = Representative Elements
B groups = transition elements
These are called the inner transition elements, and they belong here
Group 1A are the alkali metals (but NOT H)
Group 2A are the alkaline earth metals
H
• Group 8A are the noble gases
• Group 7A is called the halogensValence Electrons
• Electrons that are in the outermost
shell
• Involved in bonding with other
atoms
• Same amount as “A” group #
• Examples: H has 1, N has 5, P has 5
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Octet Rule
• Atoms are most stable when they have full outer shells
• 8 electrons (s2, p6) for most
representative elements (s2 for H, He)
Ion
• an atom (or group of atoms) that has gained or lost electrons
• anion – negative charge – gained electron(s) – tend to be larger than parent atom
• cation – positive charge – lost electron(s) – tend to be smaller than parent atom
Monoatomic Ions
• Ions from single atoms
• Charge can be predicted by location on the periodic table
• Cation – positively charged ion – lost e-
• Anion - negatively charged ion –gained e-
+1 +2
+3 -3 -2 -1
Atomic Size
• Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.
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Radius
Atomic Size - Group trends• As the atomic number increases
each atom has another energy level,
so the atoms get
bigger.
H
Li
Na
K
Rb
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Atomic Size - Period Trends
• Going from left to right across a period,
the size gets smaller.
• Electrons are in the same energy level.
• But, there is more nuclear charge.
• Outermost electrons are pulled closer.
Na Mg Al Si P S Cl Ar
Ionization Energy• Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).
• Removing one electron makes a 1+ ion.
• 2nd IE > 1st IE
• IE for non-valence electrons >>> 1st IE
IE - Group trends
• As you go down a group, the first IE decreases (less energy required, easier to remove electron)
• because • the electron is further away from the attraction of the nucleus
• Increased sheilding
Shielding
• The electron on the outermost energy
level has to look
through all the other
energy levels to see
the nucleus.
• Second electron has
same shielding, if it is
in the same period
IE - Period trends
• All the atoms in the same period have the same energy level.
• But, increasing nuclear charge
• So IE generally increases from left to right (higher energy, more
difficulty to remove electron)
Trends in Electronegativity
• Electronegativity is the tendency for an
atom to attract electrons to itself when
it is chemically combined with another
element.
• They share the electron, but how
equally do they share it?
• The higher the EN the “stronger” an atom is at attracting electrons.
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Electronegativity Group Trend
•The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has.
•Thus, more willing to share.
•Low electronegativity.
Electronegativity Period Trend
• Metals let their electrons go easily
• Thus, low electronegativity
• Nonmetals want more electrons.
• Try to take them away from others
• High electronegativity.
Electron Affinity
• The amount of energy RELEASED when an electron is added to an atom
• the more negative, the more energy
that is released, the more likely an atom
will add an electron
Trends in three atomic properties
Trends in metallic behavior