Atomic Theory Chapter 4

History sdfgfdsdfgsdfg

Democritus 400BC Greek philosopher Thought all matter was made up of atoms (atomos), which are the basic, indivisible particles of matter

Aristotle Believed all matter was continuous-did not believe in atoms His opinion was accepted for nearly 2000 years

Late 1700’s Scientists agreed that Most natural materials are mixtures of pure substances. Pure substances are elements or compounds. Law of constant composition-a compound always has the same composition.

John Dalton English schoolteacher who tied all three laws together in his atomic theory Dalton turned Democritus’s idea into a scientific theory that could be tested by experimentation.

Dalton’s Atomic Theory Elements are made of tiny particles called atoms. All atoms of a given element are identical. The atoms of an element are different from those of any other element. Atoms of one element combine with others to form compounds. Atoms are indivisible in chemical processes. A chemical reaction changes the way the atoms are grouped together.

Joseph John Thomson Cathode-ray experiment: showed that the atoms of any element can be made to emit tiny negative particles Determined the charge ratio of electrons

William Thomson Plum pudding model-a bunch of positive stuff with the electrons scattered throughout.

Rutherford, Geiger, Marsden-nucleus Gold foil experiment, which led to the discovery of the nucleus. Like bullets through a tissue

Florescent Screen Lead block Uranium Gold Foil

What he expected

Because, he thought the mass was evenly distributed in the atom.

What he got

+

Atomic Structure

Protons The number of protons in an atom determines the element’s identity Nuclear forces hold the nuclear particles together The atomic number equals the number of protons

Electrons Are very small. If the nucleus is a grape, the electrons would be about one mile away. Have a negative charge The arrangements of electrons determines the element’s chemical properties.

Neutrons Mass number= protons+neutrons Neutrons=mass number-atomic number Isotope-atoms that have the same number of protons and electrons but different numbers of neutrons (disproves point 2 of Dalton’s theory) Nuclide-any isotope of any element

Table 2.1 The Mass and Charge of the Electron, Proton, and Neutron

Practice- Give the protons, neutrons, and electrons for each Mercury Sodium Carbon 13C 6

Answers Mercury 80p, 80e, 121n Sodium 11p, 11e, 12n Carbon 6p, 6e, 6n 13C 6p, 6e, 7n 6

Ions An ion is formed when we remove or add an electron to a neutral atom. Cation-a positive ion Anion-a negative ion

Ion Example Regular sodium has 11 electrons, 11 protons, and 12 neutrons. If we take away 1 electron, it would have 10 electrons (-), and 11 protons (+) so the charge would be +1. (Neutrons would stay the same).

Periodic Table

Organization Groups/families-vertical columns Periods-horizontal rows

Metals Conductors Lose electrons (cations +) Malleable and ductile

Nonmetals Brittle Gain electrons (anions -) Covalent bonds

Semi-metals or Metalloids

Alkali Metals

Alkaline Earth Metals

Halogens

Transition metals

Rare Earth Metals (Inner transition metals)

Noble Gases

Periodic Table Label the following on your periodic table Group Charge 1 and 11 +1 2 and 12 +2 13 +3 14 +-4 Group Charge 15 -3 16 -2 17 -1 18 none

+1 -4 -3 -2 -1 +2 +3 +1 +2

Trends Notice that metals tend to give up electrons while nonmetals tend to gain electrons.

Ionic Compounds Contains a metal and a nonmetal (causes ions that is why it is called ionic) The net charge of an ionic compound has to be zero.

Ionic Compound Examples Sodium chloride Na +1 Cl -1 NaCl Magnesium chloride Mg +2 Cl -1 MgCl2