Atoms The smallest particle of an element that retains its identity in a chemical reaction.

Atomic Theory Democritus believed that atoms were indivisible and indestructible. He did not have experimental support nor did he explain chemical behavior. It took 2000 years after Democritus for the real nature of atoms and events at the atomic level to be established

Dalton’s Atomic Theory Using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory 1. All elements are composed of tiny indivisible particles called atoms 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 3. Atoms from different elements can physically mix together or can chemically combine in simple whole- number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

Structure of the Nuclear Atom Three kinds of subatomic particles: electrons, protons, and neutrons Electrons are negatively charged subatomic particles, mass=9.11 X 10-28 Protons are positively charged subatomic particles, mass = 1.67 X 10-24 Neutrons are subatomic particles with no charge, mass 1.67 X 10-24

Experiments JJ Thomson discovered the electron by passing an electric current through gases at low pressure. The gases were sealed in a glass tube fitted at both ends with metal disks called electrodes. One electrode became positively charged (anode) and the other negatively charged (cathode) This created what was called a cathode ray which was attracted to a positive plate and repelled by a negative plate

Experiments Eugen Goldstein observed a cathode-ray tube and found rays traveling in the opposite direction of those in the cathode rays. He referred to these as canal rays and concluded they were positively charged Later they became known as protons

Original Atomic Model

Rutherford’s Gold-Foil Experiment Used alpha particles (He atoms that have lost their two electrons and have a double positive charge) as a narrow beam directed at a very thin sheet of gold foil. Hypothesis-beam would easily pass straight through with small deflections Results-most passed through with no deflection, some bounced back Resulted in the atomic model containing a nucleus or central core of an atom that is composed of protons and neutrons.

Rutherford’s Atomic Model

Atomic Number Elements are different because they contain different numbers of protons Atomic Number: The number of protons in the nucleus of an atom of that element. This is used to identify an element. Example: Carbon’s atomic number is 6 because there are 6 neutrons in each Carbon atom’s nucleus For each element the number of protons equals the number of electrons. Atoms are electronically neutral, so the negative charge must equal the positive charge.

Mass Number The total number of protons and neutrons in an atom If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition. The number of neutrons in an atom is the difference between the mass number and atomic number. Number of neutrons= mass # - atomic # The composition of any atom can be represented in shorthand notation using atomic number and mass number

Isotopes Atoms that have the same number of protons but different numbers of neutrons

Isotopes Because isotopes of an element have different number of neutrons, they also have different mass numbers. Despite differences, isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior. To the right is an example of carbon isotopes

Atomic Mass Since the 1920’s the atomic mass has been able to be determined by using a mass spectrometer Because the actual masses of individual atoms are so small the atomic mass unit was developed Atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom. It is more useful to compare the relative masses of atoms using a reference isotope (carbon-12) as a standard. This isotope was assigned a mass of exactly 12 amu Example: He-4 with a mass of 4.0026amu has about one-third the mass of a carbon-12 atom whereas nickel-60 has about 5 times the mass of a carbon-12 atom.

Atomic Mass Continued… A carbon-12 atom has 6 protons and 6 neutrons in its nucleus, and a mass set at 12 amu Since the protons and neutrons account for nearly all of this mass a single proton or neutron is about 1 amu In nature most elements occur in two or more isotopes The atomic mass of an element is a weighted average of the atoms in a naturally occurring sample of the element. A weighted average reflects both the mass and the relative abundance of the isotopes as they occur in nature.

Calculating Atomic Mass for an Element Multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. Example: Element X has two natural isotopes. The isotope with a mass of 10.012amu has relative abundance 19.91%. The isotope with a mass of 11.009amu has a relative abundance of 80.09% Calculate the atomic mass of this element.

Calculating Atomic Mass For An Element Solution: Knowns: Isotope 10X has mass=10.012 amu abundance: 19.91%=0.1991 Isotope 11X has mass=11.009 abundance: 80.09%=0.8009 For 10X 10.012amu X 0.1991 = 1.993 For 11X 11.009 X 0.8009 = 8.817 For element X 1.993+8.817 = 10.810 amu The calculated mass value is closer to the mass of the more abundant isotope, which is what you would expect

The Periodic Table A periodic table allows you to easily compare the properties of one element (or a group of elements) to another element (or group of elements) Period- Each horizontal row of the periodic table. The properties of the elements vary as you move across it from element to element. Group (or family)- Each vertical column of the periodic table. Elements within a group have similar chemical and physical properties.

The Periodic Table