CHEMISTRY: Atomic Structure

Early Atomic Theory
-a little bit of history


Democritus (400 BC)


Thought that the world was made of empty space and tiny particles which he called "atoms." He proposed "atoms" as a concept but had no experimental proof.

Aristotle (384-322 B.C.)


Proposed that matter was continuous and not made up of smaller particles. He called this continuous substance "hyle." This belief was held till the 17th century.

Isaac Newton and Robert Boyle (1627-1691)(17th century)


Proponents of the idea of atoms, but could offer no proof

John Dalton (1766-1844)


Studied the observations of others concerning chemical reactions

Antoine Lavoisier (1743-1794)


Found that when a chemical change occurred in a closed system, the mass after the reaction equaled the mass before the reaction. He proposed that in ordinary chemical reactions, matter can be changed in many ways, but it cannot be created or destroyed. This became known as The Law of Conservation of Mass.

Joseph Proust (1754-1826)


Observed that specific substances always contain elements in the same ratio by mass. This became known as The Law of Definite Proportions.

Dalton's Atomic Theory

Formed the basis of our present atomic theory while trying to explain Lavoisier and Proust.


Dalton's Atomic Theory
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass and other properties. Atoms of different elements differ in size, mass and other properties.
3. Atoms cannot be subdivided, created or destroyed.
4. Atoms of different elements can combine in simple, whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.


Dalton's ideas explained the "Law of Conservation of Mass" and the "Law of Definite Proportions."


Dalton's ideas are not entirely correct, but they did form the basis of our current thinking.





Dalton's Law of Multiple Proportions


The ratio of mass of one element that combine with a constant mass of another element can be expressed in small whole numbers. (SO2, SO3; NO, NO2; SnO, SnO2)


J.L. Gay-Lussac (1778-1850) observed that gas reactions (under conditions of constant temperature and constant pressure) the volumes of reacting gases and gaseous products are in the ratio of small whole numbers.





Avogadro's (1776-1856) Hypothesis (see Ch. 11)

Equal volumes of gases, under the same conditions, have the same number of molecules





Structure of the Atom

Early Research on Atomic Particles

In the mid-1800s, experiments showed that the atom is made up of smaller particles


Discovery of the Electron

J.J. Thomson (1897)


Worked with the cathode ray tube and is credited with the discovery of the electron.

Robert Millikan (1909)


Measured the charge of an electron using his "oil drop apparatus"

Discovery of the Nucleus

Lord Rutherford (1912-13) Gold Foil Experiment


In the work he conducted with Hans Geiger and Ernest Marsden, Ernest Rutherford bombarded a very thin sheet of gold foil with a stream of positively-charged subatomic particles (alpha particles). A few of the particles "bounced back."

He concluded that most of the atom is empty space. He proposed that all the mass was concentrated in the center of the atom and this center was positively charged. This center we now call the NUCLEUS.

E. Goldstein


Discovered positively-charged particles (protons) using a modified cathode ray tube.

James Chadwick (1932)


Discovered the neutron, a particle with the same mass as a proton but having no charge.

The discovery of these subatomic particles led to a major revision of Dalton's hypothesis.


Composition of the Nucleus

The nucleus consists of protons and neutrons and has an overall positive charge which is balanced by the negative charge of the electrons. The mass of a proton is 1836 times the mass of an electron. A neutron has slightly more mass that the proton.


The number of protons determines the identity of the atom. The number of protons and neutrons combined determines the mass.


Isotopes


Number of protons in an atom determines its identity.


Number of electrons must equal the number of protons in a neutral atom. The electrons determine an atom's chemical reactivity (how it behaves).


Atoms of the same element can differ in mass (same number of protons but a different number of neutrons). The number of neutrons (plus the number of protons) determines the mass of an atom. These different forms of the atoms are called ISOTOPES.


Moseley


Worked with X-ray tubes and discovered that the wavelength of the generated X-rays was dependent on the number of protons.

(Later, Moseley proposed organizing the Periodic Table according to the atomic number rather than the atomic mass)



Properties of Subatomic Particles Particle Symbol Relative Electric Charge Mass No. Relative Mass (amu)
electron e- -1 0 0.0005486
proton p+ +1 1 1.007276
neutron n0 0 1 1.008665


Dalton's theory changed again. It now states that "all atoms of an element contain the same number of protons but may contain different numbers of neutrons."


Nuclide


a particular kind of atom which contains a definite number of protons and neutrons

Nucleons


a general term pertaining to the particles which make up the atomic nucleus. These are the protons and the neutrons

Atomic Number (Z) - the number of protons


Atomic Mass (A) - The total number of nucleons in an atom:


atomic mass = number of protons + number of neutrons


number of neutrons = atomic mass - atomic number


number of neutrons = A - Z





Atomic Mass

The proton and the neutron have essentially the same mass.


An atomic mass unit (amu) is used to measure the mass of an atom. One amu is the mass of one proton or one neutron. However, we need a reference standard.


Carbon-12 was chosen as the standard reference for atomic mass. The nucleus has 6 protons and 6 neutrons. All other elements are compared to Carbon-12.


Carbon-12 atom is defined as having a mass of 12 amu


\ Atomic Mass Unit (amu) = 1/12 the mass of a Carbon-12 atom = 1.661 x 10-27 kg





Average Atomic Mass

The mass of each type of a specific atom (each nuclide) is different due to the differing numbers of neutrons. However, most elements have more than one isotope and each isotope is present in nature at some percentage. The average atomic mass (listed on the Periodic Table) is the weighted average of the various isotopes for that element. The average atomic mass can be calculated for any element if the masses of the individual isotopes and their percent abundances in nature are known.


Mass of Isotope #1 x % abundance =


Mass of Isotope #2 x % abundance =


Mass of Isotope #3 x % abundance = ____________


Average Atomic Mass S =








THE MOLE

Chemical Symbols and Formulas


Represent elements and compounds


Elemental symbol


one atom of an element

Formula of a compound


one molecule or one formula unit

Groups of Items


We often represent groups of items using such terms as:


Dozen (12)


Gross (144)


Acre (43,560 ft2; 4840 yd2; 160 rod2)


Yard (3 ft)


Yard of concrete (cubic) = 3x3x3 = 27 ft3





Mole

A term chemists use to refer to a specific number of atoms, ions, molecules, or formula units

The SI base unit representing the chemical quantity of a substance

1 mole = 6.02 x 1023 items

1 mole of particles has a mass in grams equal to that of one particle in amu.

If a single atom has a mass of 14 amu, then a mole of the same substance has a mass of 14 g.




amu/atom = grams/mole





1 mole atoms = 6.02 x 1023 atoms


1 mole molecules = 6.02 x 1023 molecules


1 mole ions = 6.02 x 1023 ions


1 mole formula units = 6.02 x 1023 formula units





Molar Mass

The mass of one mole of atoms, ions, molecules, (anything)


Convert grams to # atoms and grams to moles and vice versa


Convert grams to moles, and then convert moles to # molecules (see the worksheet on this procedure under the Resource section)


Problems:

Problem 1. You have 6.02 x 1023 atoms of Ca (calcium). How many moles do you have?


6.02 x 1023 x 1 mole/6.02 x 1023 atoms = 1.0 mole





Problem 2. How many moles do you have in 11.5 g ethanol (C2H5OH)?


1 mole ethanol has a mass of 46 g


(2x12) + (6x1) + (1x16) = 46 g


11.5 g C2H5OH x 1 mole C2H5OH/46 g C2H5OH = 0.250 g





Problem 3. You have 1.2 x 1025 molecules NH3, how many moles do you have and what is its mass?


Process: convert molecules to moles to grams


1.2 x 1025 molecules NH3 x 1 mole NH3/6.02 x 1023 x 17 g/mole = 338 g NH3





Additional Problems:


Convert 0.638 mole Ba(CN)2 to grams


Convert 50.4 g CaBr2 to moles


Convert 1.26 moles NbI5 to grams


Convert 86.2 g C2H4 to moles