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Modern Atomic Model. Electron modeling…. To understand electrons, scientists began comparing them to light. Behavior of light. Light is a wave – similar to water waves Visible light belongs to electromagnetic spectrum. High energy. Low energy. Low Frequency. High Frequency. Spectrum.

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Modern Atomic Model

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electron modeling
Electron modeling…
  • To understand electrons, scientists began comparing them to light.
behavior of light
Behavior of light
  • Light is a wave – similar to water waves
  • Visible light belongs to electromagnetic spectrum

High energy

Low energy

Low Frequency

High Frequency







Infrared .

Long Wavelength

Short Wavelength

Visible Light

behavior of light1
Behavior of light
  • All forms of EMR travel at a constant speed of 3.0 x 108 m/s, when in a vacuum. This value also works for the speed of light in air (because air is mostly a vacuum).
behavior of light2
Behavior of light

Waves can be described in terms of the following:

  • Wavelength
  • Frequency
  • Amplitude
  • Speed
behavior of light3
Behavior of light

Waves can be described in terms of the following:

  • Wavelength - , distance between successive crests (or any 2 corresponding points), for visible light  = 400-750 nm

Behavior of light

  • Frequency - the number of waves that pass a given point per second, units are cycles/sec or hertz (Hz), abbreviated n - the Greek letter nu

c = ln

behavior of light4
Behavior of light
  • Amplitude – height from origin to crest
behavior of light5
Behavior of light
  • Speed – measured in m/s, light moves @ constant speed of 3.0 x 108 m/s, abbreviated as c
parts of a wave





Parts of a wave


parts of wave
Parts of Wave
  • Origin - the base line of the energy.
  • Crest - high point on a wave
  • Trough - Low point on a wave
  • Amplitude - distance from origin to crest
  • Wavelength - distance from crest to crest, abbreviated l (Greek letter lambda)
behavior of light6
Behavior of light
  • Because light has a constant speed, we get a relationship between  &  (c = )
frequency and wavelength
Frequency and wavelength
  • Are inversely related
    • As one goes up, the other goes down.
  • Different frequencies of light go with different colors of light.
  • There is a wide variety of frequencies
    • The whole range is called a continuous spectrum
behavior of light7
Behavior of light

Question 1: If light has  = 633 nm, what is ?

Question 2: Red light travels at 3.0 x 108 m/s and has a l of 700 nm. What is n?

Question 3: Violet light has a n of 7.5 x 1014 Hz. What is l?

wave model problems
Wave model problems

The wave model of light worked well until the beginning of the 20th century. This is because some scientists were observing light and found that what they saw did not fit the wave model.

wave model problems1
Wave model problems

Black body radiation

  • In 1900, Max Planck was studying radiation given off when matter was heated. The physics he knew said that matter could absorb or emit any quantity of energy. The results of his experiments did not fit with that idea.
light is a particle
Light is a Particle
  • Energy is quantized.
    • Light is energy
    • Light must be quantized
    • These smallest pieces of light are called photons.
  • Energy and frequency are directly related.
wave model problems2
Wave model problems
  • A quantum of light was later called a photon. Radiation is emitted or absorbed in whole numbers of photons.
wave model problems3
Wave model problems
  • To relate the quantum of energy and the frequency of the radiation, he created the relationship E = h.
energy and frequency
Energy and frequency
  • E = h x n
    • E is the energy of the photon
    • nis the frequency
    • h is Planck’s constant
      • h = 6.626 x 10 -34 Joules × seconds
wave model problems4
Wave model problems
  • What energy is given off when your stove coils turn red? (Remember that red light has a frequency of 4.29 x 1014 Hz.)
  • Which has greater energy – red or violet light?
wave model problems5
Wave model problems
  • Planck’s ideas were not immediately accepted. It was not until some time later that Albert Einstein used Planck’s equation to work on solving the photoelectric effect.
wave model problems6
Wave model problems

 Photoelectric effect

  • Light shining on certain metals can eject electrons.
wave model problems7
Wave model problems

 Photoelectric effect

  • The fact that light was able to knock electrons loose wasn’t a problem. What wave theory couldn’t explain was why only certain frequencies of light (or higher) could knock out electrons.
wave model problems8
Wave model problems

 Photoelectric effect

  • Einstein proposed that light consisted of energy quanta that behaved as particles – not waves. The quanta were called photons.
wave model problems9
Wave model problems

 Photoelectric effect

  • The photoelectric effect problem was then solved by the idea that radiation is emitted or absorbed in whole numbers of photons or radiation particles.
wave model problems10
Wave model problems

 Photoelectric effect

  • It was later proven that light could definitely act as a particle. So, we now have light acting as both a wave and as particles. (This will be the basis for understanding how e- behave.)
atomic spectrum

Atomic Spectrum

What color tells us about atoms

  • White light is made up of all the colors of the visible spectrum.
  • Passing it through a prism separates it.
if the light entering the prism is not white
If the light entering the prism is not white…
  • By heating a gas or using electricity, we can get the gas to give off colors
  • Passing this light through a prism does something different than white light
atomic spectrum1
Atomic Spectrum
  • Each element gives off its own characteristic colors
  • Can be used to identify the element
  • How we know what stars are made of
wave model problems11
Wave model problems

Bright line spectrum

  • Scientists noticed that you could vaporize an element in a flame to produce different flame colors. You can then use a prism to sort the colors to produce a line spectrum (only certain colors are produced).
wave model problems12
Wave model problems

Bright line spectrum

  • Problem: Each element produced a different line spectrum.

These are called line spectra

    • They are unique to each element.
    • These are emission spectra (the light is emitted or given off)
rutherford s model
Rutherford’s Model
  • Discovered the nucleus
  • Small dense and positive
  • Electrons around nucleus in electron cloud
bohr s model
Bohr’s Model
  • Why don’t the electrons fall into the nucleus?
  • Move like planets around the sun.
  • In circular orbits at different levels.
  • Energy separates one level from another.
bohr s model1
Bohr’s Model




Energy Levels

bohr s model2
Bohr’s Model


  • Further away from the nucleus means more energy.
  • There is no “in between” energy
  • Energy is in Levels




Increasing energy




bohr model
Bohr Model
  • Niels Bohr was able to explain the bright line spectrum. To do so, he created a model with the following parts:
      • The e- could orbit the nucleus only in allowed paths or orbits.
bohr model1
Bohr Model
  • The H atom has set energy possibilities that depend on which orbit the e- occupies.
  • The ground state occurs when the e- is in the orbit closest to the nucleus.
bohr model2
Bohr Model
  • The orbit where the e- is determines the outer dimensions of the atom.
  • The energy of the e- increases as it moves into orbits that are farther and farther from the nucleus (excited atom).
where the electron starts
Where the electron starts
  • The energy level an electron starts from is called its ground state.
changing the energy
Changing the energy
  • Let’s look at a hydrogen atom

Changing the energy

  • Heat or electricity or light can move the electron up energy levels

Changing the energy

  • As the electron falls back to ground state it gives the energy back as light

Changing the energy

  • May fall down in steps
    • Each with a different energy, frequency, and wavelength
the bohr ring atom
The Bohr Ring Atom

n = 4

n = 3

n = 2

n = 1





the bohr ring atom1
The Bohr Ring Atom
  • The farther the electrons fall, the more energy released and higher frequency produced
  • All the electrons can move
bohr model3
Bohr Model
  • Bohr said that the energy of an electron is quantized (like light) so there have to be energy levels (orbits) where the e- can be. The e- must be given a certain amount of energy to “jump” orbits.
bohr model4
Bohr Model
  • If the electron is in the lowest energy level possible (closest to nucleus), that is called the ground state.
bohr model5
Bohr Model
  • If energy is put into the e-, it goes to a higher level or an excited state. The spectral lines produced were due to the energy given off when the e- fell.
bohr model problems
Bohr Model Problems
  • Unfortunately, Bohr’s model only worked for H. What it did do was to get other scientists thinking.
de broglie

De Broglie

Determined that particles of matter could act as waves.

Described the wavelength of moving particles.


Matter exhibits both wave and particle properties!

where are the electrons
Where are the electrons?
  • We know they are outside of the nucleus.
  • We say they are in an electron cloud.
  • But where?
quantum mechanical model
Quantum-mechanical model

Takes into account the following:

  • Treats e- as waves within the atom.
quantum mechanical model1
Quantum-mechanical model
  • e- inside of atoms have specific E and occupy 3-D regions about the nucleus called orbitals. [Orbitals are different than orbits.]
quantum mechanical model2
Quantum-mechanical model
  • The size and shape of the orbitals depends on the E of the e- that occupy them.
  • All orbitals in an atom make up the e- cloud around the nucleus. The e- cloud gives the atom a size and shape.
  • The e- can’t be located exactly in the atom. There are areas of probability to find an e-.
the quantum mechanical model
The Quantum Mechanical Model
  • Has energy levels for electrons.
  • Contains orbitals of varying shapes and sizes
  • It can only tell us the probability of finding an electron a certain distance from the nucleus.
the quantum mechanical model1
The Quantum Mechanical Model
  • The electron is found inside a blurry “electron cloud”
    • An area where there is a chance of finding an electron.
atomic orbitals
Atomic Orbitals
  • Principal Quantum Number (n) = the energy level of the electron.
  • Within each energy level complex math describes several shapes.
    • These are called atomic orbitals
summary of atomic orbitals
Summary of atomic orbitals

# of shapes

Max electrons

Starts at energy level

















by energy level
By Energy Level
  • First Energy Level
  • only s orbital
  • only 2 electrons
  • Second Energy Level
  • s and p orbitals are available
  • 2 in s, 6 in p
  • 8 total electrons
by energy level1
By Energy Level
  • Third energy level
  • s, p, and d orbitals
  • 2 in s, 6 in p, and 10 in d
  • 18 total electrons
  • Fourth energy level
  • s,p,d, and f orbitals
  • 2 in s, 6 in p, 10 in d, and 14 in f
  • 32 total electrons
by energy level2
By Energy Level
  • Any thing past the fourth level - not all the orbitals will fill up.
    • You simply run out of electrons
  • The orbitals do not fill up in a neat order.
filling order
Filling order
  • Lowest energy fill first.
  • The energy levels overlap
  • The orbitals do not always fill up order of energy level.

















Increasing energy




electron configurations
Electron Configurations
  • The way electrons are arranged in atoms.
  • Aufbau principle - electrons enter the lowest energy first
    • This causes difficulties because of the overlap of orbitals of different energies.
  • Pauli Exclusion Principle - at most 2 electrons per orbital with different spins
  • Hund’s Rule - When electrons occupy orbitals of equal energy they don’t pair up until they have to
electron configurations1
Electron Configurations

Notations of e- in atoms:

  • Orbital diagram - unpaired e- represented by  or  , paired e- shown as .
    • Write orbital diagrams for elements 1-10.
  • e- configuration notation - no more lines & arrows, uses # of e- in a sublevel as a superscript over the sublevel designation
    • Write electron configuration notation for elements 1-10.
electron configurations2
Electron Configurations

Notations of e- in atoms:

  • Shortcut for orbital diagrams & e- configuration notation
    • Uses noble gas (group 18 elements) “core”
noble gas shortcuts
Noble Gas Shortcuts…
  • 1. Find element on periodic table.
  • 2. Move up 1 row on table and go to Noble Gas at end of that row
  • 3. Put this noble gas symbol inside [ ].
  • 4. Now, write out what is left over after the [ ].


  • Li =1s2 2s1
    • Noble gas in row above is He
    • [He] 2s1 is the same as 1s2 2s1
    • Be = 1s2 2s2
    • Noble gas in row above is He
    • [He] 2s2 is the same as 1s2 2s2
    • Na = 1s2 2s2 2p6 3s1
    • Noble gas in row above is Ne
    • [Ne] 3s1 is the same as 1s2 2s2 2p6 3s1
electron configurations3
Electron Configurations


  • How many sublevels are found in the 3rd energy level?
  • If there were an 8th energy level, how many sublevels would it have?
  • How many orbitals are in the s sublevel? p sublevel? d sublevel? f sublevel?
  • If there were a sublevel past the f sublevel, how many orbitals would it have?
  • How many orbitals are in the 2nd energy level?
  • How many orbitals are in the 4th energy level?
  • How many orbitals would be in the 6th energy level?
  • How many electrons are able to go into an orbital?
  • How many electrons would there be in the s sublevel? p sublevel? d sublevel? f sublevel?
  • How many electrons would you find in the 1st energy level? 4th energy level?
  • If we had a 5th energy level, how many electrons would it have?
history of the periodic table
History of the Periodic Table


  • Triads of elements with shared properties


  • method for measuring atomic masses and interpreting the results of measurements


  • arranged elements by atomic masses, properties repeated after every 8 elements → law of octaves
history of the periodic table1
History of the Periodic Table

Mendeleev & Meyer

  • arranged elements according to the increase in atomic mass
history of the periodic table2
History of the Periodic Table


  • left spaces for undiscovered elements & predicted properties of those elements
  • credited with discovering periodicity
history of the periodic table3
History of the Periodic Table

2 Questions:

  • Why could most elements be arranged by increasing atomic mass, but a few could not?
  • What was the reason for chemical periodicity?
history of the periodic table4
History of the Periodic Table


  • shooting electrons at various metals to produce X-rays
      • frequencies of the X-rays were unique to the metals
      • assigned a whole number to each element → atomic numbers
          • arrange elements by atomic numbers to get families with similar properties
history of the periodic table5
History of the Periodic Table

Periodic Table

  • an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column (group/family)
history of the periodic table6
History of the Periodic Table

Periodic Law

  • the physical and chemical properties of elements are periodic functions of their atomic numbers
periodic table information1
Periodic Table Information

Types of elements:

  • Metal – lustrous, good conductors, most are solids, malleable, ductile
  • Nonmetal - poor conductors, no luster, neither malleable nor ductile, most are gases, wide variety of other physical properties
  • Metalloid - properties of metals & non-metals
periodic table information2
Periodic Table Information


  • horizontal rows of elements

Group (family)

  • vertical columns of elements
periodic table information3
Periodic Table Information

Group (family)

  • alkali metals – group 1 (except hydrogen)
  • alkaline earth metals – group 2
  • transition metals – all of the d-block elements
  • inner transition metals – all of the f-block elements
  • metalloids – elements that touch the “staircase”
  • halogens – group 17
  • noble gases – group 18 (elements with filled outer shells of electrons)
more electron configurations
More Electron Configurations

Practice writing shortcut e- configurations.

  • Element # 15
  • Element # 8
  • Element # 34
more electron configurations1
More Electron Configurations

Group e- configurations – all elements in family have similar “endings”

Write the shortcut e- configuration for:

  • Li, Na, K
      • all end with __, so group configuration is __
  • C, Si, Ge
      • all end with __, so group configuration is __
more electron configurations2
More Electron Configurations

Practice identifying elements using group configurations:

  • Element in period 3 with group configuration p4
  • Element in period 6 with group configuration s2
  • Element in period 2 with group configuration p6
  • Element in period 7 with group configuration s1
periodic trends1
Periodic Trends

Electron configurations are able to cause periodic variations in elemental properties...

periodic trends2
Periodic Trends

Valence electrons - electrons that may be lost/gained/shared when chemical compounds are formed

  • Period  As we go across a period, the number of valence electrons increases.
  • Group  As we go down a group, we find that the number of valence electrons stays constant.
periodic trends3
Periodic Trends

Size of the atomic radius – one-half the distance between the nuclei of identical atoms joined in a molecule

  • Period  As we go across a period, the general trend is for the atomic radii to decrease.
  • Group  As we go down a group, there is a general increase in atomic radii. This happens because we are seeing more and more energy levels being added.
periodic trends4
Periodic Trends

Ionization energy – energy required to overcome nuclear attraction and remove an electron from a gaseous element

  • Period  As we go across a period, the general trend is for the ionization energy to increase.
  • Group  As we go down a group, there is a general decrease in ionization energy.
periodic trends5
Periodic Trends

Electronegativity - measure of the power of an atom in a chemical compound to attract electrons

  • Period  As we go across a period, the general trend is for the electronegativity to increase.
  • Group  As we go down a group, there is a general decrease in electronegativity.