ATOMS AND COMPOUNDSLavoisier (1743-1794) 1. Laws of Chemical Combination

Law of Conservation of Matter Matter is neither created nor destroyed in chemical reactions

Law of Constant Composition or Definite Proportions

Given compound always has same elements in samemass ratio or % composition (both are intensiveproperties)

e.g in sodium chloride, there is 1.54 times mass of chlorine in the compound as there is sodium

Example1

Exampl 2

Law of Multiple Proportions

If two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small whole numbers.

For example, considering two of the carbon oxides: CO and CO2, 100 grams of carbon may react with 133 grams of oxygen to produce carbon monoxide, or with 266 grams of oxygen to produce carbon dioxide. The ratio of the masses of oxygen that can react with 100 grams of carbon is 266:133 2:1, a ratio of small whole numbers.Applies to elements that combine to form more than one compound.

Example 3

2. Daltons Atomic Theory Matter is made up of very tiny, indivisible particles (atoms).

Atoms of each element have the same mass, different from

other elements.

Atoms combine to form molecules in small whole number ratios.

Atoms of some pairs of elements can combine in different

ratios to from different compounds .

5. When #4 occurs the 1:1 ratio will be most stable.

3.Subatomic Particles

Atom is not indivisible, but has smaller particles within- protons, neutrons and electrons. Once the atom is divided into its subatomic particles, it no longer retains the identity of the original element.All protons, neutrons and electrons are identical. We study their mass, charge and location in the atom.

Table Proton electron neutron.

Mass number(A) is an integer, not same as atomic mass.( A = Z + n = p + n )Elements identity depends on number of protons. ( Z = p )Atoms are neutral, protons equal electron. ( p = e )

Isotopes: Elements with same number of protons and differentnumbers of neutrons, they are atoms of the sameelement. e.g: 11 Hand12 H

Parts of the atomThe nucleus: the central portion of the atom containing most of the mass but least of its volume. It is composed ofProtons: positively charged subatomic particles. Each element has a unique number of protons.

Neutrons: neutral subatomic particles are called nuetrons. The atoms of an element may have different numbers of neutrons. If so, they are said to be isotopes of one another. For example, carbon has three different naturally occurring isotopes. All of the isotopes of carbon have six protons in their nuclei. One of the isotopes has six neutrons, one has seven, and one has eight.

Outside the nucleus (where the electrons live):the outer portion of the atom, containing least of the mass yet having most of the volume.

Electrons: very small subatomic particles with negative charge, move through this volume of space outside the nucleus. This volume of space is organized into shells, subshells, and orbitals. That involves quantum mechanics, which we will postpone as long as possible.

Atomic number ( Z) : It is the number of proton in an atom of an element. For example, all of the atoms of the element iron have26 protons, so the mass number of iron is 26. The atomic number is unique for each element. During nuclear reactions, an element may change into another. For clarity and as an aid to balancing nuclear reaction equations, the atomic number may be written as part of theatomic symbol. If so, it appears at the bottom left, as 26Fe.

Mass number ( A) :

the number of protons and neutrons in an atom of a specific isotope of an element. For example, oxygen has an atomic number of eight. The isotope of oxygen having ten neutrons has a mass number of 18. The atomic symbol may include the mass number, if it is relevant. The mass number is written at the top left, as 18O.

Atomic Masses

The actual masses of individual atoms cannot be measured with existing technology.

Relative atomic mass scale (or atomic weight scale) is relative to Carbon-12 (exactly 12 amu).

Some elements have atoms (same Z) with different numbers of neutrons- isotopes, thus different (A).

Hydrogen isotopes & Carbon isotopes & Berkley isotopesAtomic mass is weighted average of As

Example 4

5.Development of the Periodic Table

Mendeleyev (1834-1907)

Based on atomic masses (protons & atomic number, Z, unknown at that time).Periods (rows) show cycle through values for various physical or chemical properties, start new row when cycle repeats with a high or low value.Missing elements lead to discovery of new elements : 31, gallium; 32, germaniumCobalt and nickel out of order by atomic mass- Mendeleyev explained as experimental error in measurement of mass

Example 5

Brown's Periodic TableExplore Chemical History http://www.chemheritage.org/explore/explore.html