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MODERN ATOMIC THEORY
A.K.A.
THE ELECTRON CHAPTER W O R L D O F C H E M I S T RY C H A P T E R 1 1
WHERE DID WE LEAVE OFF?
Ernest Rutherford (1906)
Gold Leaf Experiment Results:
• Atom is mostly empty space.
• Center of atom is positively-charged and
concentrated.
Suggested that electrons traveled around the nucleus.
• Couldn’t explain why the negative electrons
weren’t attracted to positive nucleus, causing
the atom to collapse
MODERN ATOMIC THEORY
Understanding the nature of light and how it
transmits energy allowed us to understand the
structure of the atom.
In increasing energy, ROYGBV
ELECTROMAGNETIC SPECTRUM
ELECTROMAGNETIC
RADIATION
Light exhibits properties of waves as well as
particles. This is known as wave-particle duality
of light.
Most subatomic particles behave as PARTICLES and
obey the physics of waves.
ELECTROMAGNETIC
RADIATION
Wavelength
represented by
“lambda”
Waves have a frequency = the number of wave cycles per second (Hz or s-1)
ELECTROMAGNETIC
RADIATION
Long wavelength low frequency
Short wavelength high frequency
increasing
frequency
increasing
wavelength
ELECTROMAGNETIC
SPECTRUM
Quick Check: X-rays vs Radio waves
Which has the lower frequency?
Which has the shorter wavelength?
Which has more energy?
ELECTROMAGNETIC
SPECTRUM
Radio waves
X rays
X rays
EXCITED GASES
& ATOMIC STRUCTURE
Electricity on.
Excited electrons.
Gas glows.
EMISSIONS OF ENERGY
BY ATOMS
When all e– are in lowest possible energy state, an atom is
in the ground state.
If “right” amount of energy is absorbed by an e–, it can
“jump” to a higher energy level. This is an unstable,
momentary condition called the excited state.
When e– falls back to a lower-energy, more stable orbital (it
might be the orbital it started out in, but it might not), atom
releases the “right” amount of energy as light.
Any-old-value of energy to be absorbed or released is NOT
OK. This explains the lines of color in an emission spectrum.
EMISSIONS OF ENERGY
BY ATOMS
EVERY ELEMENT GIVES A UNIQUE
EMISSION SPECTRUM
ENERGY LEVELS
OF HYDROGEN
• When we add a lot of energy to a sample of H atoms
certain types of photons are produced.
• Photons are “bundles” of energy that make up light.
• We see only selected colors; this means that the
hydrogen atom must have certain discrete energy
levels.
ENERGY LEVELS
OF HYDROGEN
Energy levels are quantized, i.e. only certain values are
allowed. Excited hydrogen atoms always emit photons with
the same discrete colors (wavelengths).
Bohr’s greatest contribution to science
was in building a simple model of the
atom. It was based on an understanding
of the LINE EMISSION SPECTRA
of excited atoms.
• Problem is that the model only
works for Hydrogen.
Niels Bohr
(1885-1962)
ATOMIC LINE EMISSION SPECTRA
AND NIELS BOHR
ATOMIC SPECTRA AND BOHR
Bohr said classical view is wrong!
• e- can only exist in certain discrete orbits
• e- only possess certain amount of energy
called quanta
Bohr received Nobel prize for atomic structure in 1922.
One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.
Planetary Model
EARLY MODEL OF ATOM
PROVEN TO BE
WRONG!!!!
B O Z E M A N E M I S S I O N A N D A B S O R P T I O N S P E C T R A ( 5 : 1 7 M I N )
CRASH COURSE CHEMISTRY:
THE ELECTRON
Schrodinger applied idea of e- behaving as a
wave to the problem of electrons in atoms.
He developed the WAVE EQUATION.
Solution gives set of math expressions
called WAVE FUNCTIONS.
Each describes an allowed energy state of
an e-
E. Schrodinger
1887-1961
QUANTUM OR WAVE MECHANICS
(CURRENT MODEL)
ELECTRONS OCCUPY
ORBITALS
What is an orbital?
• NOTHING like an orbit
• probability map (based on mathematical
calculations)
• the distance the electron will most likely be from
the nucleus
• does not tell us when the electron occupies a
certain point in space or how it moves
Electrons in atoms are
arranged as LEVELS
SUBLEVELS
ORBITALS
ARRANGEMENT OF
ELECTRONS IN ATOMS
• Each energy level has a number called the
PRINCIPAL ENERGY LEVEL, n
• Currently n can be 1 thru 7, because there are
7 periods on the periodic table.
ENERGY LEVELS
ENERGY LEVELS
n=2
n=3
n=4
n=1
n=6 n=7
n=5
TYPES OF SUBLEVELS
• Sublevel(s) make up energy levels.
• The most probable area to find these electrons
takes on a shape. They are named s, p, d, and f.
• Orbitals make up sublevels. They can hold up to
2 e- with opposite spins.
• An electron can spin in one of two directions; we represent as or
• Two electrons must have opposite spins to occupy the same orbital.
PAULI EXCLUSION PRINCIPLE
2
s 1 2 8
s 1 2
s 1 2
p 3 6 18
p 3 6 d 5 10
32
REL ATIVE S IZES OF THE SP HERICAL
1 S , 2 S , AND 3 S ORBITAL S OF
HYDROGEN.
The three p orbitals lie 90o apart in space.
P ORBITALS
THE SHAPES AND LABELS
OF THE FIVE 3D ORBITALS.
F ORBITALS
• f sublevel has 7
orbitals
LABEL THE SUBLEVELS
n=2
n=4
1s
4s
3s
2s
4p
4d
4f
3d
3p
2p
DIAGONAL RULE
• Must be able to write it for the test! This will be question #1 ! Without it, you will not get correct answers!
• The diagonal rule is a memory device that helps you remember the filling order of the orbitals from lowest energy to highest energy. This is called the Aufbau principle.
• Aufbau means building up or construction in German!
s
s 3p 3d
s 2p
s 4p 4d 4f
s 5p 5d 5f 5g?
s 6p 6d 6f 6g? 6h?
s 7p 7d 7f 7g? 7h? 7i?
1
2
3
4
5
6
7
By this point, we are past the current periodic table
so we can stop.
DIAGONAL RULE Steps:
1. Write the energy levels top to
bottom.
2. Write the orbitals in s, p, d, f
order. Write the same number of
orbitals as the energy level.
3. Draw diagonal lines from the top
right to the bottom left.
4. To get the correct order, follow
the arrows!
W H Y A R E D A N D F O R B I T A L S A LWAY S I N
L OW E R E N E RG Y L E V E L S ?
• d and f orbitals require LARGE amounts of energy
• It’s better (lower in energy) to skip a sublevel that
requires a large amount of energy (d and f orbitals)
for one in a higher level but lower energy.
This is the reason for the diagonal rule!
BE SURE TO FOLLOW THE ARROWS IN ORDER!
ELECTRON CONFIGURATIONS
… it is a list of all the electrons in an atom (or ion)
1. Locate element on periodic table to determine
# of electrons (atomic #)
2. Fill orbitals in proper order using diagonal rule.
3. Check total number of electrons.
ELECTRON CONFIGURATIONS
2p4 Energy Level
Sublevel
Number of
electrons in the
sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.
LET’S TRY IT!
Write the electron configuration for the following elements:
1H
3Li
7N
10Ne
1s1
1s2 2s1
1s2 2s2 2p3
1s2 2s2 2p6
LET’S TRY IT!
19K
30Zn
82Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
…5s2 4d10 5p6 6s2 4f14 5d10 6p2
1s2 2s2 2p6 3s2 3p6 4s2 3d10
1s2 2s2 2p6 3s2 3p6 4s1
SHORTHAND NOTATION
• A way of abbreviating long electron
configurations
• Since we are only concerned about the
outermost electrons, we can skip to places we
know are completely full (noble gases), and
then finish the configuration
• Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].
• Step 2: Find where to resume by finding the next energy level.
• Step 3: Resume the configuration until it’s finished.
SHORTHAND NOTATION
• Longhand is 1s2 2s2 2p6 3s2 3p5
• You can abbreviate the first 10 electrons with a noble gas, neon.
[Ne] replaces 1s2 2s2 2p6
• The next energy level after Neon is 3
• So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17
[Ne] 3s2 3p5
EX) CHLORINE
PRACTICE SHORTHAND
NOTATION
Write the shorthand notation for each of the following atoms:
N
K
Ca
I
Bi
[He] 2s2 2p3
[Ar] 4s1
[Ar] 4s2
[Kr] 5s2 4d10 5p5
[Xe] 6s2 4f14 5d10 6p3
TRY THESE!
Write the shorthand notation for:
Cu
W
Au
[Ar] 4s2 3d9
[Xe] 6s2 4f14 5d4
[Xe] 6s2 4f14 5d9
ELECTRON CONFIGURATIONS
We need electron configurations so that we can
determine the number of electrons in the outermost
energy level (valence shell). These are called valence
electrons. The number of valence electrons determines
bonding in molecules and compounds.
VALENCE ELECTRONS
Electrons are divided between core and valence electrons (involved in bonding).
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5
OCTET RULE
The maximum number of valence electrons is
eight.
Atoms like to either empty or fill their
outermost level to achieve a full valence shell.
RULES OF THE GAME
Number of valence electrons for main group
atoms = Group number (for 1A – 8A groups)
• Ex) Mg is in group 2A and has 2 valence e-
• How many valence electrons do krypton
atom have?
IONS AND ELECTRONS
• Recall… negative ions have gained electrons, positive
ions have lost electrons.
• The electrons that are lost or gained should be
added/removed from the highest energy level
(not the highest orbital in energy!)
Skip this section….
Ex) Sn Atom: [Kr] 5s2 4d10 5p2
Sn+4 ion: [Kr] 4d10
Sn+2 ion: [Kr] 5s2 4d10
Note that the electrons came out of the highest energy level, not the
highest energy orbital!
IONS AND ELECTRONS
Atom lost 4e-
from 5s and 5p
Atom lost 2e-
from 5p
Skip this section….
• Ex) Bromine
Atom: [Ar] 4s2 3d10 4p5
Br- ion: [Ar] 4s2 3d10 4p6
Note that the electrons went into the highest energy level, not the
highest energy orbital!
IONS AND ELECTRONS
Atom gained 1e-
into 4p
Skip this section….
TRY SOME IONS!
• Write the shorthand
notation for these:
Br-
Ba+2
Al+3
• Write the longhand
notation for these:
F-
Li+
Mg+2
1s2 2s2 2p6
1s2
1s2 2s2 2p6
[Kr]
[Xe]
[Ne]
Skip this section….
ORBITALS AND THE
PERIODIC TABLE
s block p block d block
f block
Ex) Consider EC of the halogens.
9 F: 1s2 2s2 2p5
17 Cl: 1s2 2s2 2p6 3s2 3p5
35 Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
53 I: … 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
85 At: …………………………...…. ?
The last orbital to fill is the same for
elements in the same group but
different principal energy levels.
EXPLAIN THIS COMIC:
“WHO” IS THIS NEW KID?
ORBITAL DIAGRAMS
• Graphical representation of an electron
configuration
• One arrow represents one electron
• Shows spin and which orbital within a sublevel
• Same rules as before (Diagonal Rule)
ORBITAL DIAGRAMS
One additional rule: Hund’s Rule
• In orbitals of EQUAL ENERGY,
place one electron in each orbital
before making any pairs
• All single electrons must spin the same
way
LITHIUM
Atomic number = 3
1s22s1 3 total e-
CARBON
Atomic number = 6
1s2 2s2 2p2 6 total e-
Here we see for the first time
HUND’S RULE. When placing
electrons in a set of orbitals having
the same energy, we place them singly
as long as possible.
DRAW THESE ORBITAL
DIAGRAMS!
• Oxygen (O)
• Iron (Fe)
• Chromium (Cr)
• Mercury (Hg)
GENERAL PERIODIC TRENDS
As you go down a group:
Larger orbitals
Electrons held less tightly
As you go across a period:
Electrons held more tightly
Higher effective nuclear charge
TRENDS IN ATOMIC SIZE measured in a tomic radi i
As you go down a group… atoms get bigger. Why?
• electrons are added further from the nucleus and there is less attraction
• additional energy levels • shielding effect- each additional energy
level “shields” the electrons from being pulled in toward the nucleus.
As you go across a period…
atoms get smaller. Why?
• increased nuclear charge (positive charge from
protons) in same principal energy level
• each added electron feels a greater and greater +
charge because the protons are pulling in the
same direction, where the electrons are
scattered.
Small
TRENDS IN ATOMIC SIZE
TREND IN ATOMIC SIZE
WHICH IS BIGGER?
• Na or K ?
• Na or Mg ?
• Al or I ?
• CATIONS are SMALLER than the atoms
from which they come.
• The electron/proton attraction has gone UP
and so size DECREASES.
Li,152 pm 3e an 3p
+ Forming a
cation.
ION SIZES
Li+,78 pm 2e and 3p
Skip this section….
ION SIZES
• ANIONS are LARGER than the atoms from which
they come.
• The electron/proton attraction has gone DOWN and
so size INCREASES.
• Trends in ion sizes are the same as atom sizes.
Forming
an anion.
Skip this section….
TRENDS IN ION SIZES
Figure 8.13
Skip this section….
WHICH IS BIGGER?
• Cl or Cl- ?
• K+ or K ?
• Ca or Ca+2 ?
• I- or Br- ?
Skip this section….
This is called the FIRST ionization
energy because we removed only the
OUTERMOST electron
This is the SECOND IE.
IE = energy required to remove an electron from an
atom (in the gas phase).
IONIZATION ENERGY
Mg
• IE increases across a period because
the positive charge increases.
• IE decreases down a group because
size increases (Shielding Effect)
TRENDS IN
IONIZATION ENERGY
WHICH HAS A HIGHER 1 ST
IONIZATION ENERGY?
• Mg or Ca ?
• Al or S ?
• Cs or Ba ?
ELECTRONEGATIVITY
… is a measure of the ability of an atom in a
molecule to attract electrons to itself.
Concept proposed by
Linus Pauling
1901-1994
PERIODIC TRENDS:
ELECTRONEGATIVITY
• Electronegativity decreases down a group of elements
because atoms with more energy levels attract electrons
less (more shielding).
• Electronegativity increases across a period of elements
because more protons, while the energy levels are the
same, means atoms can better attract electrons.
WHICH IS MORE
ELECTRONEGATIVE?
• F or Cl ?
• Na or K ?
• Sn or I ?