History of the Atom; Modern Atomic Theory, Subatomic Particles Atomic Theory CP History of the Atom; Modern Atomic Theory, Subatomic Particles

Reasons to look at history Shows that present theory is the work of many scientists over many years To become aware of the logic used in obtaining evidence to support the theory To gain a better understanding of the properties of atoms

Greek Philosophers Democritus (460 – 371 B.C.) Stated that the world was made of tiny particles that are mostly empty space and that they are the building blocks of all matter Tiny particles could not be divided Termed them atomos, or atoms Means “uncut” or “indivisible” Aristotle (384 – 322 B.C.) Proposed that matter was continuous (i.e. came from only four elements – earth, air, fire, and water)

Dalton’s Atomic Theory John Dalton (1766 – 1844) Work marks the beginning of the development of the modern atomic theory Revived Democritus’ ideas and backed them up with scientific experimentation and data! Dalton’s Atomic Theory: Matter is composed of extremely small particles called atoms Atoms are indivisible and indestructible Atoms of a given element are identical in size, mass, and chemical properties Atoms of a specific element are different from those of another element Different atoms combine in simple whole-number ratios to form compounds In a chemical reaction, atoms are separated, combined, and rearranged

Dalton’s Atomic Theory Dalton’s theory successfully explained: The law of conservation of mass (proposed by Lavoisier) States that mass is conserved in any process, such as a chemical reaction. Conservation is the result of the separation, combination, or rearrangement of atoms The law of multiple proportions (proposed by Proust) Certain elements can combine to form two or more different chemical compounds Experimental evidence led to general acceptance of his atomic theory! Not all of the theory was accurate. It has had to be revised with new information

Modern Atomic Theory All matter is made up of a very small particles called atoms Atoms of the same element have the same chemical properties, atoms of different elements have different properties Any natural sample of an element has an average mass characteristic of that element Compounds are formed when atoms of two or more elements unite Atoms are not subdivided in physical or chemical reactions

Protons, Neutrons, and Electrons Defining the atom

The Atom Atom: Smallest particle of an element that retains the properties of the element Size: Consider this: In 2006 the population of Earth was about 6.5 x 109 people A solid copper penny contains 2.9 x 1022 atoms This is 5 trillion times larger than the worlds population!!!!!

Electrons J.J. Thomson 1879 Cathode Ray Tube Provided the first evidence that atoms are made of even smaller particles Cathode-Ray Tube: Vacuum tube, all gases removed Metal electrodes at each end Electricity runs from the negative cathode to the positive anode Using a magnet, the negative charged cathode ray bent towards a positively charged plate

Electrons Observations of the cathode-ray behavior led to the hypothesis that cathode rays consist of identical negatively charged particles Thomson proposed the “Plum Pudding” model of the atom

Thomson’s Model Found the electron 1 unit of negative charge Mass 1/2000 of hydrogen atom Later refined by Millikan to 1/1840 Concluded that there must be a positive charge since atom was neutral Atom was like plum pudding A bunch of positive stuff, with electrons able to be removed

Nucleus Ernest Rutherford – 1908 Discovered the nucleus The core of the atom and contains protons and neutrons Rutherford’s nuclear model proposed that an atom is mostly empty space This is based on studies of the bombardment of thin metal foils with fast-moving positively charged particles

Rutherford’s Experiment

Rutherford’s Experiment His experiment demonstrated that: An atom is mostly empty space There is a small dense, positive piece at the center Alpha particles are deflected if they got close enough to the positive center

Protons and Neutrons Protons: Neutrons: Rutherford also concluded that the nucleus also contained positively charged particles call protons. Subatomic particle with a positive charge equal in magnitude to the negative charge of an electron The number of protons determines the identity of an atom!! Neutrons: Discovered by James Chadwick (Rutherford’s coworker) Showed the nucleus also contained a neutral subatomic particle, the neutron Electrically neutral subatomic particle The number of neutrons can vary within an element  Isotopes!

Rutherford Activity

Structure of the Atom There are two regions The nucleus: Protons and Neutrons Positive charge Almost all of the mass Electron Cloud: Most of the volume of an atom Region were electrons can be found

Comparing Subatomic Particles Properties of Subatomic Particles Particle Symbol Relative Charge Relative Mass (Mass of Proton = 1) Actual Mass (g) Electron e- 1- 1/1840 9.11 x 10-28 Proton p+ 1+ 1 1.67 x 10-24 Neutron n0

Atomic Number Atoms of each element contain a unique positive charge in their nuclei Therefore, the number of protons in an atom identifies that particular atom! The number of protons in an atom is referred to as the atomic number Because all atoms are neutral: The number of protons = the number of electrons If you know the atomic number of an element you now know both the number of protons and the number of electrons! Atomic Number

Mass Number The total number of protons and neutrons in the nucleus of an atom Mass # = p + n Usually found at the bottom of the atomic symbol How to find the number of neutrons: Mass number – atomic number = # neutrons Example: Nitrogen: Mass number: 14 Atomic number: 7 Number of Neutrons: 7

Fill in the Gaps Atomic # Mass # P N E 3 7 39 19 6 10

Mass Number Continued Although a given type of atom will usually contain a certain number of neutrons in the nucleus, a small percentage will not Ex: most hydrogen atoms contain no neutrons A small percentage contain one neutron and a smaller percentage two neutrons What do we call atoms with a different number of neutrons?

Isotopes In nature, most elements are found as a mixture of isotopes. Atoms that have the same number of protons but different numbers of neutrons This means that the number of neutrons in a atom can change! Protons MUST stay the same Different isotopes of the same element have different mass numbers (i.e. different neutrons) A given element has at least two isotopes, and they all have different mass numbers Example: Gadolinium (Gd) has 6 stable isotopes

Isotopes Isotopes are chemically alike because they have identical numbers of protons and electrons It is useful to compare the relative masses of atoms to a standard reference isotope Carbon-12 is the standard reference isotope Carbon-12 has a mass of exactly 12 atomic mass units (amu)

Atomic Mass A weighted average mass of the atoms in a naturally occurring sample of the element Mass of an atom in atomic mass units (amu) Equal to 1/12th of the mass of carbon Weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature This is the number which appears on the periodic table

Atomic Mass Number Example: 99% of all carbon atoms are the isotope containing 6 neutrons, the remaining 1% is the heavier isotope containing 7 neutrons, which raises the aver mass of carbon from 12.000 to 12.011

Calculations Carbon has two major, stable isotopes. Carbon-12 (98.93% abundance) Carbon-13 (1.07% abundance) Carbon-14 is in trace amounts and is radioactive Calculate the average atomic mass of carbon Carbon 12: (0.9893 x 12) = 11.8716 Carbon 13: (0.0107 x 13) = 0.1391 AAM: 12.0107

Calculations Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium? Average atomic mass = (0.722 x 85) + (0.278 x 87) Average atomic mass = (61.37) + (24.186) Average atomic mass = 85.56 amu