Post on 04-Jan-2016
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Electrons in Atoms
• Why do ions have the charges they have? Like Al3+ or Fe2+ or Fe3+ or O2-
• Why does an atom become an ion in the first place?• Why are the BrINClHOF elements the only ones that
make molecules with themselves? Why don’t any other elements do that?
• How do fireworks make colors when they explode?• How do fluorescent lights work?• How do we know how hot the sun is?• How is there life on this planet?• The electrons in atoms can answer all of these
questions and so many others.• It all happens because of electrons……..
Dalton’s Model of the Atom1803
• Atoms are tiny, indestructible spheres
• No internal structure
Thomson’s Model1897
• Referred to as the “plum-pudding” model.
• The whole atom is a sphere of positive charge, with little negative electrons embedded in it.
Rutherford’s Model1911
• Small, dense core of positive charge.
• Electrons circle the nucleus in fixed orbits.
Rutherford’s Model
• Electrons revolve around the nucleus like planets around the sun (fixed orbits).
• This model failed to explain some properties of atoms.
Niels Bohr’s Model1913
• Electrons orbit the nucleus in specific orbits a fixed distance away.
Neils Bohr’s Model (1913)
• They orbit at a particular energy level. They can move to a higher level, but they need energy to do so.
• A quantum of energy is the required amount to move an e- to a higher level. Exactly this amount, no in-between.
Neils Bohr’s Model (1913)• This model of the atom had shortcomings.
It failed to explain some phenomenon in nature.
• So, a better version was still out there waiting to be discovered…..
Waves• To understand the electronic structure of
atoms, one must understand the nature of electromagnetic radiation.
• The distance between corresponding points on adjacent waves is the wavelength ().
Waves
• The number of waves passing a given point per unit of time is the frequency ().
• For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.
Light
Electromagnetic Radiation
• All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.00 108 m/s.
• Therefore,c =
• This all suggests light is a wave….
The Nature of Energy• The wave nature of light
does not explain how an object can glow when its temperature increases.
• Max Planck explained it by assuming that energy comes in packets called quanta.
• Equantum = hν
• Einstein used this assumption to explain the photoelectric effect.
Photoelectric Effect
• The emission of electrons from a metal when light is shined upon the metal.
• Depending on the metal used, only light of a certain wavelength (color) would cause an electron to be emitted.
Photoelectric Effect• Einstein concluded
that energy is proportional to frequency:
Ephoton = hwhere h is Planck’s constant, 6.626 10−34 J-s.
• This suggests light as a particle….
Nature of Energy
• Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or particle, of that light:
c = E = h
The Nature of Energy
• Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.
The Nature of Energy
• For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source.
• Only a line spectrum of discrete wavelengths is observed.
The Nature of Energy
• Niels Bohr adopted Planck’s idea of quanta and explained these phenomena in this way:1. Electrons in an atom can only occupy certain orbits
(corresponding to certain energies).
2. Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by
E = h
The Nature of Energy• The energy absorbed or
emitted from the process of electron promotion or demotion can be calculated by the equation:
• ∆E = -Rh(1/nf2 - 1/ni
2)
• where RH is the Rydberg constant, 2.18 10−18 J, and ni and nf are the initial and final energy levels of the electron.
The Wave Nature of Matter
• Louis de Broglie posited that if light can behave with material properties (photons), matter should exhibit wave properties.
• He demonstrated that the relationship between mass and wavelength was
λ = h/mv
• In other words, if light waves can act like particles, then things can move like waves.
Heisenberg’s Uncertainty Principle
• Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known:
• (x) (mv) h/4π
• For regular-sized objects, the uncertainty is practically zero, but in many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
Quantum Mechanics
• Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
• It is known as quantum mechanics.
• Amazingly accurate in describing electrons and microscopic behaviors, but also exceedingly strange.
Richard Feynman on Quantum Mechanics (1965)
• “There was a time when the newspapers said that only twelve men understood the theory of relativity. But after people read the paper a lot of people understood the theory of relativity. On the other hand I think I can safely say that nobody understands quantum mechanics.”
Schrodinger’s Wave Equations
• The wave equation is designated with a lower case Greek psi ().
• The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.
Quantum Numbers
• Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.
• Each orbital describes a spatial distribution of electron density.
• An orbital is described by a set of three quantum numbers.
Principal Quantum Number, n
• The principal quantum number, n, describes the energy level on which the orbital resides.
• The values of n are integers ≥ 1.
• n also describes the relative size of the orbital, 2 larger than 1, and so on.
Angular Momentum Quantum Number (l)
• This quantum number defines the shape of the orbital.
• Allowed values of l are integers ranging from 0 to n − 1.
• We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals. This is where s, p, d & f come into play
Value of l 0 1 2 3
Type of orbital s p d f
Magnetic Quantum Number (ml)
• The magnetic quantum number describes the three-dimensional orientation of the orbital.
• Allowed values of ml are integers ranging from -l to l:
−l ≤ ml ≤ l.• Therefore, on any given energy level,
there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.
Magnetic Quantum Number (ml)
• Orbitals with the same value of n form a shell.
• Different orbital types within a shell are subshells.
s Orbitals
• The value of l for s orbitals is 0.
• They are spherical in shape.
• The radius of the sphere increases with the value of n.
s Orbitals
• Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.
p Orbitals
• The value of l for p orbitals is 1.
• They have two lobes with a node between them.
d Orbitals
• The value of l for a d orbital is 2.
• Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
Energies of Orbitals
• For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
• That is, they are degenerate.
Energies of Orbitals
• As the number of electrons increases, though, so does the repulsion between them.
• Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.
Spin Quantum Number, ms
• In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.
• The “spin” of an electron describes its magnetic field, which affects its energy.
Spin Quantum Number, ms
• This led to a fourth quantum number, the spin quantum number, ms.
• The spin quantum number has only 2 allowed values: +1/2 and −1/2.
Pauli Exclusion Principal
• No two electrons in the same atom can have exactly the same energy.
• Therefore, no two electrons in the same atom can have identical sets of quantum numbers.
Electron Configurations
• This shows the distribution of all electrons in an atom.
• Each component consists of – A number denoting the energy
level,– A letter denoting the type of
orbital,– A superscript denoting the
number of electrons in those orbitals.
Orbital Diagrams
• Each box in the diagram represents one orbital.
• Half-arrows represent the electrons.
• The direction of the arrow represents the relative spin of the electron.
Hund’s Rule
• “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”
Periodic Table
• We fill orbitals in increasing order of energy.
• Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.
Atomic Orbitals• Region of space where there is a high
probability of finding an electron.
• Principal Quantum #(n) --denotes the energy level of electrons (1,2,3,4,etc.)
• Also denotes the # of sublevels at that energy level (s,p,d,f)
• Sublevels describe the shapes and sizes of orbitals where e- may be found.
Shapes of Orbitals• s-spherical, with nucleus at the center
• p-dumbbell, or figure-8, with nucleus at the center
• d—as shown
• f—as shown
• As you increase energy levels, the shape of each remains the same, but size gets larger.
Electron Configurations• Orbitals of an atom will fill so that the atom
is in its most stable state. There are 3 rules that govern this:
• Aufbau Principle- e- occupy lowest-energy orbitals first
• Pauli Exclusion Principle- 2 e- in same orbital must have opposite spin
• Hund’s Rule- e- occupy orbitals of the same energy so that there’s a max # of same spin e-
Exceptions to Aufbau
• If you did the configuration for Cu according to the three rules, it would look like this:
• 1s22s22p63s23p64s23d9
• In actuality, it is this:
• 1s22s22p63s23p64s13d10
Another…• Chromium, Cr, also is an
exception to the Aufbau Principle
• According to Aufbau, Cr should have this configuration:
• 1s22s22p63s23p64s23d4
• But it actually has this:
• 1s22s22p63s23p64s13d5
Why Would an Atom Do This?
• Because a filled shell is the most stable arrangement, and a half-filled shell is the next best arrangement.
Valence Electrons• The electrons that exist in the
outermost energy level of an atom are valence electrons.
• A full shell or a half-filled shell is the most stable arrangement.
• Noble gases always have a full valence, or a full outer shell, which is what every other element is trying to
achieve. (Max. of 8 valence electrons)
• What does the term orbital describe?
• A region around the nucleus where an electron is most likely to be found.
• What does an element’s electron configuration describe?
• All of the orbitals that the element’s electrons occupy, and how those electrons are distributed.
• We do not need to focus on all the electrons that an atom has, we really only need to focus on the valence electrons. Why?
• Because they are the outermost electrons, and they are the only electrons that can possibly interact with other atoms.
• How many valence electrons does Oxygen have?
• 6
• Why are the alkali metals so reactive?
• They all have an s1 electron (1 valence electron) that they are trying to lose.
• Noble gases are also called inert gases. Why are the noble gases so unreactive?
• Because they have a full outer shell (eight valence electrons) and do not need any more or less electrons.
First Periodic Table• In 1869, the first table
having elements organized by their properties was published by a Russian chemist and professor named Dmitri Mendeleev.
• He listed them in order of atomic mass.
Gallium and Germanium:
Discovered in 1875 & 1886
• Mendeleev arranged the elements in order of increasing mass.
• In the 1860’s, the proton was not yet discovered.
• In 1913, British physicist Henry Moseley arranged the elements in order of increasing atomic number (# of protons).
Some Vocabulary…
Vertical columns are called groups or families.
Horizontal rows are called periods.
• How many elements are in period 2? 8
• How many elements are in period 6? 32
• How many elements are in group 2? 6
The Periodic Law (Cont.)
• Elements within a column of a group have similar properties.
• Properties in a period change as you move across a period from left to right.
• The pattern of properties within a period repeats as you move form one period to the next.
• Periodic Law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Electron Configurations in Groups
Helium (He) 1s2
Neon (Ne) 1s22s22p6
Argon (Ar) 1s22s22p63s23p6
Krypton (Kr) 1s22s22p63s23p63d104s24p6
Lithium (Li) 1s22s1
Sodium (Na) 1s22s22p63s1
Potassium (K) 1s22s22p63s23p64s1
Alkali Metals
Noble Gases
Blocks of Elements
Metals• 80% of elements
Metals
Nonmetals
Metalloids
Metals (Cont.)
• Conductors of heat
• Conductors of electric current
• High luster
• Ductile
• Malleable
• Solids @ room temp. (except Hg)
Nonmetals
• Most are gases @ room temp
• Poor conductors of heat
• Poor conductors of electric current
• Solid nonmetals are brittle
Metalloids
Properties similar to those of metals and nonmetals
Behaviors can be controlled by changing the conditions.
– Example: Silicon
Classifying the Elements
• Group 1A Elements – Alkali Metals
• Group 2A Elements – Alkaline Earth Metals
• Group 7A Elements – Halogens
• Group 8A Elements – Noble Gases
• Group B Elements – Transition Metals
• Below the Main Body- Inner Transition Metals
Periodic Trends
• Atomic Size– Atomic radius – one half the distance between the
nuclei of two atoms of the same element
• Increases from top to bottom within a group• Decreases from left to right across a period
Atomic Radius vs. Atomic Number
Increase within
Group 1
Shielding Effect
Periodic Trends in Atomic Size
Size decreases
Size
Incre
ases
Trends in Ionization Energy
• Ionization Energy– The energy required to remove an electron from an
atom.
• First Ionization Energy– The energy required to remove the first electron
from at atom.
First ionization energy tends to decrease from top to bottom within a group and increase from
left to right across a period.
First Ionization Energy vs. Atomic Number
Why is the 1st ionization energy for
the noble gases higher?
Periodic Trends in Ionization Energy
Energy Increases
En
erg
y D
ecre
ases
• Electronegativity- The ability of an atom of an element to attract electrons when the atom is BONDED to another atom in a compound.
Periodic Trends
• Metallic properties—as shown. As you approach the nonmetals, metallic properties decrease.