4-1 Notes. Defining the Atom

Transcription

1 4-1 Notes Defining the Atom

2 Early Models of the Atom All matter is composed of atoms Atoms are the smallest particles of an element that retains their identity in a chemical reaction Greek philosopher Democritus (460B.C.-370B.C) was the first to suggest the existence of atoms He believed that atoms were indivisible and indestructible His ideas did not explain chemical behavior

3 Dalton s Atomic Theory 2000 years later John Dalton used experimental methods to transform Democritus s ideas on atoms into a scientific theory Dalton studied the ratios in which elements combine in chemical reactions. He developed the atomic theory which had 4 main points

4 Dalton s Atomic 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 of 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 are never changed into atoms of another element as a result of a chemical reaction

5 Sizing up the Atom Atoms are extremely small The radii of most atoms fall within the range of 5 x m to 2 x m. Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes. By using this microscope, individual atoms can be moved around and arranged in patterns This technology is important in future applications in medicine, communications, solar energy, and space exploration

6 4-2 Notes Structure of the Nuclear Atom

7 Subatomic Particles Dalton s atomic theory is still accepted today except for 1 thing, that atoms are not indivisible. An atom can be broken apart into 3 different particles electrons, protons, and neutrons.

8 Electrons Discovered by English physicist J.J. Thompson His experiment included passing electric current through gases while they were sealed in a tube (fig. 4.4) This produced a cathode ray (glowing beam) that traveled from the cathode to the anode Since the ray was attracted to the + plate and repelled by the plate, he concluded that the particles had a negative charge Named them electrons Robert Millikan later concluded that each electron carried a -1 charge.

9 Protons There were four things that were known to be true about atoms atoms have no net electric charge electric charges are carried by particles of matter electric charges always exist in whole-numbered ratios when a given number of negatively charged particles combine with the same number of positively charged particles, an electrically neutral particle is formed Therefore, there had to be a positive particle in the atom. In 1886 Eugen Goldstein found evidence of a positive particle he named a proton He also determined the mass was 1840 times that of an electron

10 Neutrons In 1932, English physicist James Chadwick ( ) discovered the neutron Neutrons are subatomic particles with no charge but a mass approximately equal to that of a proton

11 The Atomic Nucleus Based on experiments Ernest Rutherford performed in 1911 he concluded the following The majority of an atom is empty space. All the positive charge & most of the mass is concentrated into a small region Called the nucleus Composed of protons and neutrons The electrons are distributed around the nucleus and occupy almost all the volume of the atom. Rutherford s model was an improvement of Thompson s but it turned out to be incomplete

12 4-3 Notes Distinguishing Among Atoms

13 Atomic Number Elements are different because they contain different numbers of protons The atomic number = the number of protons in an element Find the # of protons for H, C, K

14 Mass Number Protons + neutrons = the mass number The number of neutrons in an atom is the difference between the mass number and the atomic number # of neutrons = mass # - # of protons Find the # of neutrons for C, Nb, Ag, The composition of an atom is often written in shorthand notation using atomic number and mass number Ex. Gold - 197

15 Isotopes Isotopes are atoms that have the same number of protons but different numbers of neutrons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers Hydrogen has three isotopes What are they?

16 Atomic Mass Since the actual masses of elements are extremely small, they are difficult to work with. Therefore, we find the relative masses of atoms using a reference isotope as a standard The isotope chosen was Carbon-12, and it was given a mass of exactly 12 atomic mass units (amu) An atomic mass unit is defined as 1/12 th of the mass of carbon 12

17 Atomic Mass Carbon 12 has 6 protons and 6 neutrons, therefore each proton or neutron is 1 amu Due to the existence of isotopes, you can t accurately predict the atomic mass of any element. Most elements occur as a mixture of 2 or more isotopes, each with a specific mass and natural percent abundance The atomic mass of an element is the weighted average mass of the atoms in a naturally occurring sample of the element

18 Calculating Atomic Mass To calculate the atomic mass of an element: 1. multiply the mass of each isotope by its natural abundance 2. Add the products The sum is the weighted average mass of the atoms Sample Problem 4.2

19 The Periodic Table A periodic table is an arrangement of elements based on a set of repeating properties Each element is identified by its symbol in a square Each horizontal row of the periodic table is called a period Each vertical column of the periodic table is called a group or family Element of the same group have the same physical and chemical properties