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History and Progression of Atomic Theory. Structure of the Atom. An atom is the smallest particle of an element that retains the chemical properties of that element. The nucleus is a very small region located at the center of an atom.

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structure of the atom
Structure of the Atom
  • An atom is the smallest particle of an element that retains the chemical properties of that element.
  • The nucleusis a very small region located at the center of an atom.
  • The nucleus is made up of at least one positively charged particle called a protonand usually one or more neutral particles called neutrons.

Surrounding the nucleus is a region occupied by negatively charged particles called electrons.

  • Protons, neutrons, and electrons are often referred to as subatomic particles.

400 BC

  • This is the Greek philosopher Democritus who began the search for a description of matter more than 2400 years ago.
    • He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?
democritus atomic theory 400bc

Democritus’ Atomic Theory 400BC

Democritus asserted that space contained an infinite number of particles

Named atomos, “not cutting” or "indivisible”

Atoms are eternal and invisible; absolutely small, so small that their size cannot be diminished; totally full and incompressible.

Atoms are homogeneous, differing only in shape, arrangement, position, and number

  • To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes.
  • Atoms were infinite in number, always moving and capable of joining together.
dalton s theory early 1800 s
Dalton’s Theory (early 1800’s)
  • He deduced that all elements are composed of atoms. Atoms are indivisible and indestructible particles.
  • Atoms of the same element are exactly alike.
  • Atoms of different elements are different.
  • Compoundsare formed by the joining of atoms of two or more elements in specific ratios.
  • In chemical reactions, atoms are combined, separated, and rearranged
  • Explained the Law of the Conservation of Mass

Matter is neither nor destroyed during ordinary chemical reactions or physical changes

modern atomic theory
Modern Atomic Theory
  • Atomic Theory has been modified
  • We now know that atoms are divisible into smaller particles
  • And that elements can have atoms with different masses (isotopes)
j j thomson english physicist1897
J.J. Thomson (English physicist1897)
  • Used Cathode Ray tube to determine the presence of

– (electrons) and + (protons) particles.




Current passed through a glass tube filled with various gases at low pressure from the cathode to the anode (terminals of the voltage source)

  • Stream of particles glowed (cathode ray)
  • Cathode rays deflected away from negatively charged objects
  • Led to the hypothesis that cathode ray particles are negatively charged
  • These negatively charged particles were named electrons
j j thomson
J.J. Thomson
  • Plum Pudding Model -- the structure of an atom is something like pudding. He assumed that the basic body of an atom is a spherical object containing electrons & protons randomly confined in homogeneous jellylike material. Positive charges cancel the negative charges.
  • This model was soon disproved


ernest rutherford eng 1911
Ernest Rutherford (Eng.1911)
  • Atoms have a central positive nucleus surrounded by negative orbiting electrons.
  • This idea was the result of his famous Gold Foil Experiment(see next slide). This experiment involved the firing of radioactive alpha particles through gold foil.
  • Alpha particles = + charged particles about 4x mass of H
  • This model suggested that most of the mass of the atom was contained in the small nucleus, and that the rest of the atom was mostly empty space.
  • Most of particles passed straight through the foil but approximately 1 in 8000 were deflected.
  • Conclusion- volume of the nucleus very small compared to the volume of the atom
  • To put the size or the nucleus in perspective, if the nucleus of an atom was the size of a marble, the atom would be the size of a football field!
niels bohr danish 1913
Niels Bohr (Danish 1913)
  • The Bohr Model is probably familiar to us as the "planetary model" of the atom is used to symbolize atomic energy.
  • Electrons orbit the nucleus much like planets orbiting the Sun.
  • Electrons travel in certain orbits at specific distances from the nucleus with definite energy levels (energy shells or energy levels)
  • Lower energy state near the nucleus and higher energy state as you move away from the nucleus
  • Electrons gain or lose energy by jumping between orbits, absorbing or emitting electromagnetic radiation

Think of rungs on a ladder, except the distance between the rungs is not constant!

chadwick english 1932
Chadwick (English 1932)
  • James Chadwick discovered a third type of particle, which he named the Neutron.
    • Neutrons help to reduce the repulsion between protons and stabilize the atom's nucleus.
    • Neutronsalways reside in the nucleus of atoms and they are about the same mass and size as protons.
    • Neutrons do not have any electrical charge; they are electrically neutral
quantum model of the atom
Quantum Model of the Atom
  • Louis de Broglie (French scientist 1924)
    • Hypothesized that electrons have wavelike properties
    • Electrons exist at certain frequencies corresponding to specific energies
    • Confirmed experimentally by investigating diffraction (bending of wave as it passes through small opening) and interference (waves overlap)
cloud theory
Cloud Theory
  • Based on the work of many scientists
  • Based on the mathematical approach of Quantum Mechanics
  • Electrons are assigned regions of space (Orbitals) not pathways (Orbits)
  • Electrons are moving around the nucleus rapidly in no predictable path producing a cloud of e-’s over time. Think of a rapidly moving fan blade.
cloud theory of today
Cloud Theory of Today

Electron Cloud

quantum theory
Quantum Theory
  • Describes mathematically the wave properties of electrons and other very small particles
  • Schrodinger Wave Equation (Austrian 1926)- showed mathematically how electrons can exist as waves of certain energies and therefore certain frequencies
heisenberg uncertainty principle heisenberg german 1927
Heisenberg Uncertainty Principle(Heisenberg- German 1927)
  • Electrons act as particles and waves
  • How can you locate electrons in an atom?
    • Observe interference with photons
    • This knocks the electron off course
  • “It is impossible to determine simultaneously both the position and velocity of an electron or any other particle”

Atomic Number

  • Atoms of different elements have different numbers of protons.
  • Atoms of the same element all have the same number of protons.
  • The atomic number(Z) of an element is the number of protons of each atom of that element.
mass number
Mass Number
  • The mass number is the total number of protons and neutrons that make up the nucleus of an isotope.
how to find of protons
How to find # of Protons
  • Number of protons in an atom is ALWAYS equal to the Atomic Number
number of electrons e
Number of Electrons (e-)
  • Atoms – Number of protons and electrons are equal (overall neutral charge)
  • Ions
    • Loss of electron(s) makes positive ions
    • Gain of electron(s) makes negative ions
loss or gain of electrons
Loss or Gain of Electrons
  • Atom Ion
  • Na  Na+1 + 1e- (LOSS)
  • +11 +11
  • -11 -10
  • 0 net +1 net charge
  • Cl2 + 2e-  2Cl- (GAIN)
  • +17 +17
  • -17 -18
  • 0 net -1 net
how to calculate of neutrons
How to calculate # of Neutrons
  • Atomic Mass (rounded to integer)
  • - Atomic Number
  • ----------------------------------------------------
  • Number of Neutrons in the nucleus
  • Atomic Mass – Atomic Number = # Neutrons
atom contents
Atom Contents
  • Protons (p+) – always equal to Atomic #
  • Electrons (e-)
    • In Atoms – Same as the # of Protons
    • In Ions – Net charge after e-’s have been lost or gained in an attempt to become stable
      • Loss of e-’s = Positive charge
      • Gain of e-’s = Negative charge
  • Neutrons (n0) = Atomic Mass – Atomic #
  • Isotopes
  • Isotopes are atoms of the same element that have different masses.
  • The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons.
  • Most of the elements consist of mixtures of isotopes.
designating isotopes
Designating Isotopes
  • Hyphen notation: The mass number is written with a hyphen after the name of the element.
  • uranium-235
  • Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number.

The number of neutrons is found by subtracting the atomic number from the mass number.

  • mass number − atomic number = number of neutrons
  • 235 (protons + neutrons) − 92 protons = 143 neutrons
  • Nuclideis a general term for a specific isotope of an element.

Sample Problem A

  • How many protons, electrons, and neutrons are there in an atom of chlorine-37?
  • Solution:
  • atomic number = number of protons = number of electrons
  • mass number = number of neutrons + number of protons
  • # P+= 17
  • # e-= 17
  • # N0= mass number - #P+= 37 – 17 = 20

Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound


Law of multiple proportions: if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

relative atomic masses
Relative Atomic Masses

The standard used by scientists to compare units of atomic mass is the carbon-12 atom, which has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu.

One atomic mass unit,or 1 amu, is exactly 1/12 the mass of a carbon-12 atom.

The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

average atomic masses of elements
Average Atomic Masses of Elements

Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.

Calculating Average Atomic Mass

The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

average atomic masses of elements continued
Average Atomic Masses of Elements, continued

Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 amu, and 30.85% copper-65, which has an atomic mass of 64.927 amu.

The average atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.

average atomic masses of elements continued1
Average Atomic Masses of Elements, continued

(0.6915 × 62.929 601 amu) + (0.3085 × 64.927 794 amu) = 63.55 amu

The calculated average atomic mass of naturally occurring copper is 63.55 amu.

hotel head to toe activity
Hotel Head to Toe Activity
  • How many floors were there in the hotel?
  • How many double beds were in the s rooms? The p rooms? The d rooms? The f rooms?
  • How many guests would each room hold?
  • What was special about how guests had to sleep in the beds?
  • We will now compare our hotel and its rooms and guests with electrons filling orbitals in atoms

Guests = electrons

  • Floors = Energy levels (1, 2, 3, 4, 5, 6, or 7)
  • Rooms = Energy sub-levels (s, p, d, and f)
  • Double beds = Orbitals
  • The way guests sleep in the bed = Electron spin
quantum numbers
Quantum Numbers
  • Quantum numbers specify the properties of atomic orbitals and electrons in orbitals
    • Energy level- principal quantum number (n)
    • Orbital shape- angular momentum quantum number (l)
    • Orientation around the nucleus- magnetic quantum number (m)
    • Electron spin- spin quantum number
principal quantum number n
Principal Quantum Number (n)
  • Indicated main energy level for the electron
  • Values of 1, 2, 3, …
  • n=1 at the orbital closest to the nucleus, lowest energy level
  • As n gets larger, electron’s energy and distance from the nucleus increase
angular momentum quantum number l
Angular Momentum Quantum Number (l)
  • Orbital shape; sublevels
magnetic quantum number m
Magnetic Quantum Number (m)
  • Orientation around the nucleus
  • Orbitals with same shape but different location in 3-D coordinate system
  • Only one s orbital
  • Three p orbitals; px, py, and pz
  • Five d orbitals
  • Seven f orbitals
orbitals shapes
Orbitals shapes
  • S or spherical orbital
orbital shapes
Orbital Shapes
  • P orbitals or perpendicular orbitals
  • y x z
orbital shapes1
Orbital Shapes
  • D orbitals or Diagonal orbitals
spin quantum number 1 2 or 1 2
Spin Quantum Number (+1/2 or -1/2)
  • Each orbital can contain up to two electrons
  • The two electrons must have different spin states
aufbau rule
Aufbau Rule
  • An electron occupies the lowest-energy orbital that can receive it.
  • Example: Take the least number of steps to get to your hotel room. Start filling rooms on the first floor and then the second floor. Sometimes it takes fewer steps to go up a floor to an s room, than to walk to the end of the hall on the current floor (d or f room)
hund s rule
Hund’s Rule
  • "Orbitals of equal energy are each occupied by one electron before any is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin"
  • Example: One person fills each bed in the room before doubling up!
  • Example: each p orbital at a particular energy level must receive one electron before they can have 2 electrons
pauli exclusion principle
Pauli Exclusion Principle
  • No two electrons in the same atom can have the same four quantum numbers
    • Think of your home and address
    • Street number, city, zip code, state
  • Two electrons is the same orbital must have different spin quantum numbers (opposite spin)
  • Example: sleep head to toe
  • Electron Configuration packet