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The Electron. 6.0 Chemistry. Development of the Periodic Table. History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H 2 , N 2 , O 2 , C). By the mid 1800’s, about 60 elements had been identified.

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the electron

The Electron

6.0 Chemistry

development of the periodic table
Development of the Periodic Table
  • History of the Periodic Table – By the end of the 1700’s, scientists had identified only 30 elements (ex. Cu, Ag, Au, H2, N2, O2, C).
  • By the mid 1800’s, about 60 elements had been identified.
  • Sept 1860 – chemists assembled at the First International Congress of Chemists in Germany to settle the controversial issues such as atomic mass. Standard values set for atomic mass and improved communication for research.
johann dobereiner 1817
Johann Dobereiner: 1817

Organized the elements into sets of three with similar properties.

He called these groups triads. The middle element is often the averageof the other two.

Ex) Cl – 35.5

Br – 79.9

I – 126.9


Avg Sr


b john newlands 1866
B. John Newlands: 1866
  • Arranged elements in order of increasing atomic mass.
  • Noticed repeating patterns in the elements’ properties every 8th element.
  • Law of Octaves - properties of elements repeated every 8th element.
  • There were 62 known elements at the time.
c dmitri mendeleev 1869
C. Dmitri Mendeleev: 1869
  • Arranged elements in order of increasing atomic mass.
  • Similar properties occurred after periods (horizontal rows) of varying lengths.
  • Organized the 1st periodic table according to increasing atomic mass and put elements with similar properties in the same column.
  • Periodic – repeating properties or patterns
  • Noticed inconsistencies in the arrangement.
He arranged some elements out of atomic mass order to keep them together with other elements with similar properties. (Notice Te and I)
  • He also left several blanks in his table.
  • In 1871, he correctly predicted the existence and properties of 3 unidentified elements – Sc, Ga and Ge
  • These elements were later identified and matched his predictions.
1 st periodic law
1st Periodic Law

Properties of the elements repeat periodically when the elements are arranged in increasing order by atomic mass

Mendeleev is known as the

Father of Chemistry

#101 honors Mendeleev

d henry moseley 1911
D. Henry Moseley: 1911
  • Studied X-ray spectral lines of 38 metals. Each element had a certain amount of positive charge in the nucleus which are called protons.
  • Analyzed data and found that the elements in the PT fit into patterns better when arranged in increasing nuclear charge, which is the Atomic Number.
  • The Modern Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.
glenn seaborg seaborgium sg 106
Glenn Seaborg “Seaborgium” Sg #106
  • Born in 1912 in Michigan, Seaborg proposed reorganizing the Periodic Table one last time as a young chemist working on the Manhattan Atomic Bomb Project during WWII.
  • He suggested pulling the “f-block” elements out to the bottom of the table.
  • He was the principle or co-discoverer of 10 transuranium elements.
  • He was awarded the Noble prize in 1951 and

died in 1999.

seaborgium is the exception
Seaborgium is the exception…
  • After some argument between the USA and the rest of the world, element 106 was named Seaborgium shortly before he died. This was a matter of some controversy because the International Union of Pure and Applied Chemistry, IUPAC, the body that deals with naming in chemistry, had previously ruled that elements should not be named after living people.
parts of the periodic table
Parts of the Periodic Table
  • Horizontal Rows – PERIODS
    • There are 7 periods in the periodic table
    • Elements in a period do NOT have similar properties.
  • Vertical Columns – GROUPS or FAMILIES
    • Labeled 1-18
    • IA-VIIIA are the Main-group or representative elements.
    • Elements in a group have similar properties.
    • Why?
family names write these on your p t
Hydrogen (1)

Alkali metals (1) – most reactive metals; reactivity increases down the group

Alkaline earth metals (2)

Boron family (13)

Carbon family (14)

Nitrogen family (15)

Oxygen or Chalcogen family (16)

Halogens (17)

Noble gases (18) - inert

Transition elements or metals (3-12): d-block

Inner transition elements or metals (f-block)

Lanthanides or lanthanide series

Actinides or actinide series

Transuranium elements

Family Nameswrite these on your P.T.
e metal nonmetals and metalloids semimetals
E. Metal, Nonmetals and Metalloids (Semimetals):

- Good conductors of heat & electricity

- High melting points most solids at room temperature

  • Metals

Found on the

LEFT side of

the PT

2. Nonmetals

Located on the

RIGHT side of

the PT

3. Metalloids - Properties of both metals & nonmetals

- High luster (shiny)

- Ductile (can be drawn into thin wire)

- Malleable (bends without breaking)

- High densities

- Reacts with acids

- Brittle (easy to break)

- No luster (dull)

- Insulators nonconductors

- Neither ductile nor malleable

- Nonreactive with acids


review of early atomic theories
Review of Early Atomic Theories
  • Dalton 
  • Thomson 

(plum pudding)

Bohr – electrons in a particular path have a fixed energy called energy levels
    • Rungs of a ladder
  • Quantum Mechanical (Schrödinger) Model
    • Electrons better understood as WAVES
    • Does not tell where the electrons are located
    • Electrons have a certain amount of energy - QUANTIZED
light as a wave
Light as a Wave

Characteristics of a Wave

  • Amplitude:Height of the wave from the baseline. The higher the wave the greater the intensity.
  • Wavelength: (λ , “lambda”) in nanometers (1 x 10-9 m). Distance between similar points on 2 consecutive waves.
  • Frequency: (ν , “nu”) The number of waves that pass a fixed point per unit of time. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz)

ex) Radio FM 93.3 megahertz (MHz) is 93.3 x 106Hz (cycles/sec)











D. Electromagnetic Radiation- a form of energy that exhibits wavelike behavior as it travels through space- all forms of EM radiation move at the speed of light

speed of light c
Speed of Light (c)

E. 3.00 x 108 m/s or 186,000 miles/sec.

The relationship between wavelength and

frequency can be shown with the following


This is an indirect relationship.

If λ then ν .

c = λ ν


Visible Light




Gamma Rays


















quantum theory
Quantum Theory
  • Planck’s Hypothesis: (Max Planck 1900)
  • Studied emission of light from hot objects
  • Observed color of light varied with temperature
  • Suggested the objects do not continuously emit E, but emit E in small specific amounts
    • Light is absorbed or emitted in a little packet or bundle called a quantum (quanta –plural).
    • Quantum = minimum amount of E that can be lost or gained by an atom
    • Energies are quantized.

(Think steps not a ramp)






Incandescent light bulbs give off most of their energy in the form of heat-carrying infrared light photons -- only about 10 percent of the light produced is in the visible spectrum. This wastes a lot of electricity. Cool light sources, such as fluorescent lamps and LEDs, don\'t waste a lot of energy generating heat -- they give off mostly visible light. For this reason, they are slowly edging out the old reliable light bulb.

max planck s energy equation
Max Planck’s Energy Equation
  • Proposed that energy is directly proportional to frequency.

E = h

Planck’s equation for each quantum

h = Plank’s constant = 6.626 x 10-34 J.s

This is a direct relationship.

As energy increases, frequency increases.

albert einstein
Albert Einstein

While well-known for the equation E=mc2 , Einstein’s work on the photoelectric effect resulted in being awarded the 1921 Nobel Prize in Physics.

(1879 – 1955)

German Physicist


Albert Einstein and the Photoelectric Effect

Refers to the emission of electrons from a metal when light shines on the metal


  • Electrons are ejected by light of sufficient energy. Energy minimum is different for different metals.
  • The current (# of electrons emitted/s) increases with brightness of the light.

albert einstein and the photoelectric effect
Albert Einstein and the Photoelectric Effect


  • Proposed that light consists of quanta of energy that behaves like particles.
  • Quantum of light = photon = massless particle that carries a quantum of energy.
  • Proposed the Dual Nature of Light: its wave and particle nature.
    • Light travels through space as waves
    • Light acts as a stream of particles when it interacts with matter.


Definition: a method of studying substances that are exposed to some sort of continuous exciting energy.

A. Emission Line Spectra: contains only certain colors or wavelengths (  ) of light.

1. Every element has its own linespectrum (fingerprint).


Continuous Spectrum – White Light

Line Spectrum – Excited Elements

gas discharge tubes
Gas Discharge Tubes
  • Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.
flame tests
Flame Tests
  • used to test qualitatively for the presence of certain metals in chemical compounds. 
  • the heat of the Bunsen flame excites electrons that emit visible light.

Copper(II) sulfate

Lithium chloride

Potassium chloride

Barium nitrate

  • Uses a diffraction grating to diffract the light into particular wavelengths of light.

A Line Spectra result from excited elements - as electrons of an element gain energy and rise to an excited state they then fall back to their ground state in the same pattern producing the same energy drop each time which we see as individual wavelengths of light.

iii atomic spectra and the bohr model of hydrogen 1913
III. Atomic Spectra and the Bohr Model of Hydrogen (1913)

Neils Bohr - Danish Scientist

Explained the bright-line spectrum of hydrogen


  • Added E as electricity to H gas at low pressure in a tube.
  • Emitted E as visible light, was observed through a prism

Result: Hydrogen emitted 4 distinct bright lines of color, aka bright line spectrum


Electrons absorb energy to rise to a higher or excited state and emit energy in the form of a photon of light as they fall back to their ground states.

path of an excited electron as it falls back to the ground state
Path of an excited electron as it “falls” back to the Ground State
  • When electrons gain energy, they jump to a higher energy level (excited state).
  • Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state.
  • As they fall, they emit electromagnetic radiation.
  • Depending on how far they fall determines the type of radiation (light) released.
bohr model of hydrogen
Bohr Model of Hydrogen


  • *Unique line spectrum is due to quantized electron energies.
  • *Electrons are in specific orbits related to certain amounts of energy known as stationary states.
  • *Orbits are related to energy levels.
  • *Energy levels are identified as E1, E2, E3, … (n = 1, 2, 3, …)
  • *Lowest energy level = ground state
  • *Electrons absorb certain amounts of energy to move to a higher energy level farther away from the nucleus = excited state
  • *Electrons return to the more stable ground state and release a photon that has energy equal to the difference in energy between the energy levels.
    • from E2 to E1: Ephoton = E2 – E1 (difference in energy)
the bohr atom for hydrogen a model
The Bohr Atom for Hydrogen -a Model
  • Successful in calculating the wavelength, frequency, & energy of hydrogen’s line spectrum.
  • Successful in calculating the energy needed to remove hydrogen’s electron

H(g) + energy  H+1(g) + 1e-

Calculated ionization E = observed ionization E

= 1312.1 kJ/mol

lyman balmer and paschen series of the hydrogen atom
Lyman, Balmer and Paschen series of the Hydrogen Atom
  • Lyman series: electrons fall to n = 1 and give off UV light.
  • Balmer series: electrons fall to n = 2 and give off visible light.
  • Paschen series: electrons fall to n = 3 and give off infrared light.
When electrons absorb energy they jump to a higher (excited) state.

n=2 n=3 n=4 n=5 n=6 n=7

Electrons are not stable. Radiation (light) is emitted when an electron falls back from a higher level to a lower level.


Infrared Light

Visible Light

Ultraviolet Light

atomic spectra
Atomic Spectra


Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.




limitations of the bohr model
Limitations of the Bohr Model
  • Model could not calculate the wavelengths of observed spectra of multi-electron atoms.
  • Model could not explain the chemical behavior of atoms.
  • Bohr used classical mechanics to understand the behaviors of small particles.
  • The Bohr model is also known as the planetary, solar system, or satellite model.
quantum mechanical model of the atom
Quantum Mechanical Model of the Atom
  • Louie De Broglie (1924-5)
    • Took Einstein’s idea that light can exhibit both wave and particle properties
    • Very small particles (like electrons) display properties of waves.
    • Behavior of electrons in Bohr’s quantized orbits was similar to behavior of waves

French scientist

  • Known: any wave confined to a space can only have specific frequencies
  • De Broglie suggested electrons are waves confined to the space around the atomic nucleus.
  • Electrons could exist only at specific frequencies which correspond to specific energies (E = h quantized E of Bohr)
quantum mechanical model of the atom1
Quantum Mechanical Model of the Atom
    • Experimentally proven in 1927 by diffraction of electrons by Davisson & Germer (showed diffraction of electrons by a crystal of Ni)
  • Wave-Particle Duality of Nature
    • Light and electrons (very small particles like electrons, atoms, molecules) have properties of waves and particles  QUANTUM MECHANICS (based on WAVE properties)

**Large objects obey the laws of classical mechanics**

c werner heisenberg 1927
C. Werner Heisenberg: (1927)
  • Heisenberg’s Uncertainty Principle: states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.
  • You cannot predict future locations of particles.
  • He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.
d erwin schr dinger wave equation 1926
Wave nature of an electron is described by a mathematical equation.

Four quantum numbers in the equation are used to describe an electron’s behavior – location and energy.

Electron is treated as a wave with quantized energy.

Describes the probability of the electrons found in certain locations around the nucleus.

D. Erwin Schrödinger Wave Equation (1926)

(1887 – 1961) Austrian Physicist

electron density
Electron Density

An orbital is a region in

which an electron with a

particular energy is likely

to be found.

Where the density of an

electron cloud is high there

is a highprobability that is

where the electron is

located. If the electron

density is low then there is

a lowprobability.


E. Atomic Orbitals - region around the nucleus where an electron with a particular energy is likely to be found (not the same as Bohr’s orbits!)

  • Orbitals have characteristic shapes, sizes, &


2. Orbitals do not describe how the electronmoves.

3. The drawing of an orbital represents the 3-dimentionalsurface within which the electron is found 90% of the time.

s p and d orbitals
s,p and d orbitals

s orbital

p orbitals

d orbitals

For a more complete representation and presentation of atomic orbitals go to

quantum numbers
Quantum Numbers
  • Each quantum number provides more specific information on the probable location of an electron.
  • Each electron within an atom can be described by a unique set of 4 quantum numbers.
quantum numbers finding an address for each electron
Quantum Numbers - Finding an address for each electron:
  • “state” Principle Quantum Number (n) or the energy level;
    • Describes the relative size of the electron cloud.
    • Positive integer values (n = 1 to n = 7)
  • “city” Sublevel (l)
    • Describes theshape of the electron cloud.
    • The maximum number of sublevels within a level = n
    • Shapes are s, p, d,or f.
    • Lowest energy = s Highest energy = f
quantum numbers cont
Quantum numbers cont.
  • “street” Orbital (ml) odd # of orbitals
    • Describes the orientation or direction in space
      • s – 1orbital
      • p – 3 orbitals (x, y, z)
      • d – 5 orbitals (xy, yz, xz, x2 – y2, z2)
      • f – 7 orbitals (y3 – 3yx2, 5yz2-yr2, x3-3xy2,

zx2-zy2, xyz, 5xz2-3xr2, 5z3-3zr2)

    • Orbitals within the same sublevel have the same energy are called degenerate orbitals
    • An orbital can hold a maximum of 2 electrons
quantum numbers cont1
Quantum numbers cont.
  • “house” Spin (ms)
    • Describes the direction of electron spin in an orbital.
    • The clockwise or counterclockwise motion of electrons.
    • Only electrons with opposite spins can occupy the same orbital.
    • The opposite spin is written as

+1/2 or -1/2 or  or

e electron configurations
E. Electron Configurations:
  • Shorthand notation for indicating the number of electrons in each level, sublevel, and orbital.


  • Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess.
electron configuration rules
Electron Configuration Rules
  • The Aufbau Principle: electrons are added one at a time to the lowest energy orbital available.
pauli exclusion principle
Pauli Exclusion Principle:

1. Each orbital can only hold 2 electrons.

2. The electrons must have opposite spins.

s-sublevel = max 2electrons

p-sublevel = max 6 electrons

d-sublevel = max 10 electrons

f-sublevel = max 14 electrons


hund s rule
Hund’s Rule:
  • Electrons will remain unpaired in degenerate orbitals before they pair up.

incorrect ↑↓↑ __

correct ↑ ↑ ↑

















Increasing energy





Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin.

Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel)

Aufbau Principle: Electrons fill lowest energy levels first (n=1)

electron configuration examples
Electron Configuration Examples:

Ex) electron configuration for Na:

1s2 2s2 2p6 3s1

Ex) orbital filling box diagram for Na:

electron dot diagrams
Electron Dot Diagrams:

Write the symbol for the element.

Place dots around the symbol to represent the

valences & p electrons only.

Do NOT include d & f orbitals in diagram.

p orbital electronss orbital electrons

electron configuration orbital box diagram electron dot diagram
Electron Configuration Orbital Box Diagram Electron-dot Diagram




What does the Tellurium electron-dot resemble???

unpaired vs paired electrons filled and half filled orbitals
Unpaired vs. Paired ElectronsFilled and Half-filled orbitals
  • Atoms with unpaired electrons are said to be paramagnetic. These are weakly attracted to a magnetic field.
  • Atoms with all paired electrons are said to be diamagnetic. These are weakly repelled from a magnetic field.
  • ½ filled and filled orbitals have special stability
noble gas or shorthand electron configurations
Noble Gas or Shorthand Electron Configurations
  • Rb
  • Se
  • At

Draw the Dot Diagrams for these elements

exceptions to the rules
Exceptions to the Rules
  • Max stability - ½ filled and filled orbitals
    • Cr
    • Mo
    • Cu
    • Ag
    • Au
exceptions to the rules1
Exceptions to the Rules
  • Max stability - ½ filled and filled orbitals
    • Cr
    • Mo
    • Cu
    • Ag
    • Au
electron configuration for ions
Electron Configuration for Ions
  • K
  • P
  • Al
  • Se
  • K+1
  • P-3
  • Al+3
  • Se-2
excited vs ground state
Excited vs. Ground State
  • If an electron absorbs energy, it is in an EXCITED state

Ne: 1s22s22p53s1

  • How is this different from the ground state configuration?

Ne: 1s22s22p6

c electron configuration families
C. Electron Configuration & Families

Carbon has 4 valence electrons

  • Valence electrons – outermost electrons (s and p); responsible for bonding and chemical behavior.
  • Elements in the same group have the same number of valence electrons.
electron configurations
Electron Configurations
  • Stable Octet: 8 electrons in the outer level is very stable (includes He)
    • Ions – gain/lose electrons to achieve a stable octet
    • Isoelectronic – same electron configuration
      • Examples: N, O, F, Na, Mg, Al are isoelectronic with Ne – this is called an isoelectronic series
  • Pseudoisoelectronic – same electron configuration but includes the d orbitals
    • Fe+2 is pseudoisoelectronic with Ar