The amazing world of the Atom

· Home · Schedule · Tutor · Labs · Demonstrations · Chem Matters · Makeup · Notes

What Is An Atom?

I - Earliest Models Of The Atom

~ 400 B.C.

    1. Democritus, a Greek Philosopher, came up with the idea of the atom by thinking about it.
    2. Democritus' hypothesis was that all things were made up of small, indivisible particles, which he called atoms.

 384-322 B.C.

    1. Aristotle rejected Democritus’ theory.
    2. He believed all things were infinitely divisible and composed of only one substance called hyle. This idea held throughout the Middle Ages of Europe.

1700’s

    1. Charle Coulomb showed the relationship between the amount of electric force that two charged particles exert upon each other and the distance separating them.
        1. Laws of attraction and repulsion (opposites attract, likes repel)
        2. As distance increases the electrical force decreases.

B. Antoine Lavoisier studied closed systems and chemical reactions.

        1. Law of conservation of matter.
        1. He found that in ordinary chemical reactions, matter can be changed in many ways, but it cannot be created or destroyed.

C. Joseph Proust observed that specific substances always contain elements in the same ratio by mass. This came to be known as the law of definite proportions.

 

II - Smaller Than The Smallest

1800’s

    1. John Dalton (1803) combined the work of Lavoisier and Proust to help him develop his atomic theory.

1. He hypothesize that all matter is made up of very small particles called atoms. Dalton’s ideas were similar to those of Democritus.

2. Dalton believed that atoms of different elements were not the same, and that atoms of one element were exactly alike.

3. Dalton also stated that atoms can join with other atoms in small ratios to form compounds.

4. The law of multiple proportions.

a. Because atoms cannot be divided these ratios can be expressed in small whole numbers.

B. Avogadro, in 1811, suggested that equal volumes of different gases, under the same conditions (pressure, temperature) contain same number of atoms.

C. Mendaleev developed the periodic table of elements, which finalized the belief in the existence of atoms.

D. In 1896, Henri Becquerel accidentally discovered that uranium ores emit invisible rays that fog photographic plates.

E. Marie Curie and her husband, Pierre, found that these rays came from the atoms and were smaller than the atom. Marie was the first to use the term radioactivity.

        1. She also discovered two new radioactive elements, polonium and radium.

2. Becquerel and the Curies put an end to Dalton's idea that the atom was indivisible.

F. J.J. Thomson developed his model of the atom when he discovered that atoms are divisible.

1. In 1897, he invented the cathode ray tube. In a vacuum tube a luminous beam flowed between the negative (Anode) and positive (Cathode) terminals of a high voltage supply.

2. Using the cathode ray tube with magnets and electrical plated he discovered that the cathode ray was made up of negatively charged particles.

a. He calculated these particles to have a mass 1/2000 as much as a hydrogen atom, the

lightest known atom.

b. He called these sub-atomic particles electrons.

 

 

 Make-up Questions 

The Story So Far

DEFINITIONS

Element - A type of matter that cannot be chemically broken down.

Atom - The smallest particle of an element that posses the properties of that element.

Molecule - A particle made up of two or more atoms.

Compound - Two or more elements combined by chemical bonds.

 

III – The Sub – Atomics

1900’s

    1. Max Planck, in 1900, introduced the idea of quanta to explain the spectrum of light emitted by certain heated objects.
        1. Energy is not given off continuously, but in small packets called quanta.

B. In 1905, Albert Einstein broadened Plank’s idea to explain photoelectric effect.

1. Einstein mathematically determined that light was made up of particles that had discrete energies.

2. He called these particles photons. Until that time, light was thought to be only energy carried in the form of waves.

 

C. Robert Millikan, in 1909, calculated the charge on the electron in his "oil drop" experiment.

1. Electrons were transferred from the atomizer to the oil drops. The negatively charged particles fell through a vacuum under the influence of gravity.

2. Millikan adjusted the voltage on the plates to offset the pull of gravity. He found that each droplet had a different charge, but all the charges were a multiple of one small charge. He assumed that small charge to be the charge of one electron.

 

1900’s Continued

  D. Ernest Rutherford developed his atomic model from observations made in his gold- foil experiment.

1. Rutherford found that one of the invisible rays discovered by the Curies had a positive charge, was extremely small, and moved very fast. He called this an alpha particle.

2. Rutherford bombarded a thin piece of gold with alpha particles.

a. The majority of particles passed straight through the foil. This indicated to him that the atom was mostly empty space.

b. Some particles did however hit something in the foil. From the fact that the alpha particle is positively charged and the scattered pattern the particles produced, Rutherford stated that in the center of each atom is a dense nucleus that is positively charged.

3. In 1911, Rutherford developed an atomic model that had a positively charged nucleus and was

surrounded by electrons.

4. Rutherford also predicted the existence of uncharged particle in the nucleus.

 Make-up Questions 

 

1900’s continued

 

    1. In 1913, Neils Bohr combined the work of Plank (spectral analysis) and Rutherford (positively charged nucleus) to develop his atomic model.

1. Bohr’s model had a positively charged nucleus surrounded by electrons that were in defined orbits around this nucleus.

2. The distance of each orbit from the nucleus was fixed.

3. An electron could change orbitals if it was sufficiently excited with incoming energy.

4. With his model, he could calculate the spectral lines for hydrogen.

a. This was because he predicted that a photon of light energy was emitted each time an

electron fell from one orbital to another.

b. The color of the photon was determined by how far the electron fell to its original

orbit.

 

 

Properties of Sub-Atomic Particles

Electrons Protons Neutrons

1._______________ 1._______________ 1._______________

_______________ ________________ _______________

2._______________ 2._______________ 2._______________

_______________ _______________ _______________

3._______________ 3._______________ 3._______________

_______________ _______________ _______________

4._______________ 4._______________ 4._______________

_______________ _______________ _______________

5._______________ 5._______________ 5._______________

_______________ _______________ _______________

6._______________ 6._______________ 6._______________

_______________ _______________ _______________

 

Atomic Number : 1.___________________________________

2.___________________________________

3.___________________________________

 

Atomic Mass Number :

A ____________ number closest to the atomic mass. It is

the ________ of the number of ________ and __________

in the nucleus of the atom.

Atomic Number = Number of

___________

Number of Protons = Number of

______________

Number of Neutrons = Atomic Mass Number

Minus the Number

of __________

 

Notice !

You will notice, the atomic mass number is the same as

the Molar mass or Gram formula mass number.

Atomic Mass The ______________ average of the masses of the ___________ of that

element.

Isotope Atoms that have the ___________ number of ____________ but

_________________ numbers of _____________.

Mass Number Mass Number

30 __ __

X X

15 __ __ (isotope)

Atomic Number Atomic Number

Periodic Trends

in Atomic Size

Using the above figure make the following two generalizations.

1. Within a group (column) atomic size generally tends to _______________ when going from top to bottom. A result of increasing the number of principal _______________ levels.

2. Within a period (row) atomic size generally tends to _______________ when moving left to

right. A result of the effective ______________ charge acting on its electrons.

On your Periodic Table "A" mark the trends using arrows, with the arrow pointing in the direction of increasing size. Be sure to include the legend in the "KEY :" area of the chart.

 

IV - Wave Mechanical Model

    1. In 1924, Louis de Broglie proposed that electrons have the properties of waves.
    2. Werner Heisenberg developed the uncertainty principle.
        1. The uncertainty principle when applied to the electron states that it is inherently impossible for us to know simultaneously both the exact momentum of the electron and its exact location in space.
    1. Erwin Schrodinger, in 1926, proposed an equation that incorporates both the wavelike and particlelike behavior of the electron. This is known as a wave function.
    2. From the work of Broglie, Heisenberg, and Schrodinger we have the wave mechanical model of the atom.

1.

 

 

 

 

Why was the wave mechanical model of the atom developed?

 

 

Because electrons behave as waves they move in different patterns. Draw and label the "s" and three "p" orbitals here:

Electron Arrangements

Elements are arranged in the periodic table according to _______________ number. The atomic number corresponds not only to the number of ______________ in the nucleus of the neutral atom, but also to the number of ______________ in that atom.

 

The arrangement of electrons in an atom is called the _________________ configuration

All things in nature seek to be at the ______________ possible energy level.

Think of objects falling from the sky seeking the lowest potential energy of the

Earth’s surface. The same is true for ________________. The most ____________ , or ground, electron configuration of an atom is that in which the _____________ are in the ____________ possible energy states.

The orbitals of an atom are ____________ in order of ________________ energy.

 

In the Quantum Mechanics model of the atom, the Principal _______________ levels (n=1, 2, 3, etc.) of electron orbitals are divided into ____________ based on the energy an electron would release if it should fall to the ____________ state. Each energy _____________ corresponds to a different electron cloud ____________, denoted by the letters ______ (spherical shaped clouds), ______ (dumbbell shaped clouds), ________ and ________ (complex shaped clouds). The cloud ______________ is where there is a high ______________ of finding an electron.

Electron Arrangement Rules

1. Electrons enter the ______________ energy level first. (Aufbau Principle)

2. ____________ orbital (sublevel) has more than _________ electrons. (Pauli

Exclusion Principle)

3. __________ electron ___________ takes place until all orbitals

(sublevels) in an energy level are half filled. (Hund Rule)

4. ___________ filled or ____________ filled levels have extra stability.

 

Sample of electron filling (electron configuration)

 

 

 

 

 

 

 

 

 

 

 

 

 

 

If you understand how the periodic table is organized, it is not necessary to memorize the order in which __________ fill. You can write the electron configuration of an element based on its __________ in the periodic table.

 

 

 

 

 

 

 

 

 

 

 

Electron-dot symbols are a ____________ way of showing the ________________ electron shell of an atom. The electron-dot symbol for an element consists of the chemical __________________ for the element plus a __________ for each outer shell (valence) electron. Dots are placed in __________ "regions" around the atomic symbol: the top, the bottom, and the left and right sides. Each of these regions can accommodate _________ electrons.

 

 

Marking your periodic table on the "Eight Families of Elements"

1. Mark the generalized electron configuration for each family.

2. Mark the generalized electron-dot symbol for each family.

3. Complete the atomic information on the "Adopt An Element" paper.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Avogadro Meets Mr. Mole

By doing some quick mathematics with the _________ of protons, neutrons and

electrons, chemists were able to calculate that one __________ of atoms contains

approximately ____________ atoms in it. This number, 6.02 X 1023, is called

_______________ Number.

The same is true for molecules, there are _______________ molecules in a mole.

(Be careful not to confuse moles (fuzzy creatures) with molecules!)

These _______________ can be written as: ____________________ or __________________

 

 

 

 

Avogadro and You

Solve the following challenges showing ALL of your WORK using the Picket Fence.

1. How many atoms are in 3 moles of calcium?

 

 

 

2. How many atoms are in 2.5 moles of Na?

 

 

 

3. How many molecules are in 0.3 moles of strontium nitrate?

 

 

 

4. How many molecules are in 23 grams of sodium?

 

 

 

5. How many atoms are in 86 grams of beryllium?

 

 

 

 

6. How many atoms are in 54 kilograms of fluorine?

 

 

 

 

7. How many molecules are in 24 g of CuSO4?

 

 

 

8. How many molecules are in 17 mg of O3, ozone?

 

 

 

9. Calculate the number of molecules of CaCl2 in 750 mL of a .55 M solution.

 

 

 

 

10. Suppose an atom of carbon had a diameter of 1 cm. The earth has a

circumference of 45,000 Km. How many circumferences (time around) of the

world would 6.02 X 1023 atoms cover?