Chapter 3: Atoms: The Building Blocks of Ma;er

Transcription

1 Chapter 3: Atoms: The Building Blocks of Ma;er

2 The Atom: From Philosophical Idea to Scien7fic Theory The Greek philosopher Democritus (460 B.C. 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word atomos ) He believed that atoms were indivisible and indestruc-ble.

3 Democritus vs. Aristotle His ideas did agree with later scien7fic theory, but did not explain chemical behavior, and was not based on the scien-fic method but just philosophy. Aristotle believed that atoms did not exist because Democritus could not provide experimental data that supported his theory.

4 Development of Atomic Theory The Aristotelean Philosophy Advocated by Plato and Aristotle Main focus: There are no ul7mately indivisible par7cles. This philosophy strengthens throughout Western culture un7l the seventeenth century.

5 Development of Atomic Theory 17 th century The idea of the atom is reborn! Sir Isaac Newton ( ) Becomes major backer of the idea Most thinkers are trying to explain the proper7es of gases at the 7me.

6 The Three Basic Laws of Chemistry All things chemistry are based upon the three basic laws of chemistry: Law of conserva7on of mass Law of definite propor7ons Law of mul7ple propor7ons

7 The Basic Laws of Chemistry Law of Conserva7on of Mass 1789: Antoine Lavoisier discovers during an experiment that the mass of red mercury oxide (before hea7ng) was equal to the mass of the newly formed mercury metal and oxygen gas.

8 The Basic Laws of Chemistry Law of Constant Composi7on (or Definite Propor7ons) : Joseph Proust proposes this law from data obtained in experiments using copper carbonate. Proust found that all samples used had the same composi7on.

9 The Basic Laws of Chemistry 1803: John Dalton creates an explana7on for the law of conserva7on of mass and the law of constant composi7on. As a result of his work, Dalton creates another basic law of chemistry. Law of mul7ple propor7ons

10 Basic Laws of Chemistry LCM: Ma_er cannot be created or destroyed. LDP: A chemical compound contains the same elements in the same propor7ons by mass regardless of size or source of the sample. LMP: If 2 or more different compounds are composed of the same elements, then the ra7o of the masses of elements is always a small whole number.

11 John Dalton & Atomic Theory 1808: John Dalton proposed an explana7on for the 3 basic laws of chemistry. John Dalton is responsible for the current defini7on of an atom. Dalton s Atomic Theory has 5 main points. *You will have a quiz over Dalton s Atomic Theory on Friday!!

12 Dalton s Atomic Theory 1. All ma_er is composed of extremely small par7cles called atoms. 2. Atoms of a given element are iden7cal in size, mass, and other proper7es; atoms of different elements differ in size, mass, and other proper7es. 3. Atoms cannot be subdivided, created, or destroyed.

13 Dalton s Atomic Theory (cont d) 4. Atoms of different elements combine in simple whole- number ra7os to form chemical compounds. 5. In chemical reac7ons, atoms are combined, separated, or rearranged.

14 Does Dalton s Theory S7ll Hold? Not all por7ons of Dalton s Atomic Theory are s7ll valid. We know that atoms can be subdivided into smaller subatomic par7cles. We also know that a given element can have atoms with different masses. But points #1 and #2 have remained unchanged.

15 ATOMIC STRUCTURE

16 Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negamvely charged parmcle: the electron

17 Mass of the Electron The oil drop apparatus 1916 Robert Millikan determines the mass of the electron: 1/1837 the mass of a hydrogen atom; has one unit of negamve charge

18 Conclusions from the Study of the Electron: a) Cathode rays have idenmcal propermes regardless of the element used to produce them. All elements must contain idenmcally charged electrons. b) Atoms are neutral, so there must be posimve parmcles in the atom to balance the negamve charge of the electrons. c) Electrons have so li;le mass that atoms must contain other parmcles that account for most of the mass.

19 Conclusions from the Study of the Electron:! Eugen Goldstein in 1886 observed what is now called the proton - parmcles with a posimve charge, and a relamve mass of 1 (or 1837 Mmes that of an electron)! 1932 James Chadwick confirmed the existence of the neutron a parmcle with no charge, but a mass nearly equal to a proton

20 Thomson s Atomic Model J. J. Thomson Thomson believed that the electrons were like plums embedded in a posimvely charged pudding, thus it was called the plum pudding model.

21 Ernest Rutherford s Gold Foil Experiment ! Alpha parmcles are helium nuclei - The alpha parmcles were fired at a thin sheet of gold foil! ParMcles that hit on the detecmng screen (film) are recorded.

22 Rutherford s Findings! Most of the parmcles passed right through! A few parmcles were deflected! VERY FEW were greatly deflected Conclusions: a) The nucleus is small b) The nucleus is dense c) The nucleus is posimvely charged

23 The Rutherford Atomic Model Based on his experimental evidence: The atom is mostly empty space All the posi7ve charge, and almost all the mass is concentrated in a small area in the center. He called this a nucleus The nucleus is composed of protons and neutrons (they make the nucleus!) The electrons distributed around the nucleus, and occupy most of the volume His model was called a nuclear model

24 Atomic Number Atoms are composed of iden%cal protons, neutrons, and electrons How then are atoms of one element different from another element? Elements are different because they contain different numbers of PROTONS The atomic number of an element is the number of protons in the nucleus # protons in an atom = # electrons

25 Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon Phosphorus Gold

26 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass number = p + + n 0 Element # p + # n 0 # e - Mass no. Carbon Nitrogen Sodium Uranium Radon

27 Nuclear Symbols Contain the symbol of the element, the mass number and the atomic number. Superscript Subscript Mass X number Atomic number

28 Nuclear Symbols n Find each of these a) number of protons b) number of neutrons c) number of electrons d) Atomic number e) Mass Number 80 35Br

29 Nuclear Symbols n If an element has an atomic number of 34 and a mass number of 78, what is the: a) number of protons b) number of neutrons c) number of electrons d) complete symbol

30 Nuclear Symbols n If an element has 91 protons and 140 neutrons what is the a) Atomic number b) Mass number c) number of electrons d) complete symbol

31 Isotopes Dalton was wrong about all elements of the same type being iden7cal Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers. These are called isotopes.

32 Isotopes Frederick Soddy ( ) proposed the idea of isotopes in 1912 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioac7ve materials.

33 Naming Isotopes We can also put the mass number a6er the name of the element: carbon- 12 carbon- 14 uranium- 235

34 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Nucleus Hydrogen 1 (pro7um) Hydrogen- 2 (deuterium) Hydrogen (tri7um)

35 Elements occur in nature as mixtures of isotopes. Isotopes

36 Isotopes Nuclide p + n 0 e - Mass no. Oxygen Arsenic Phosphorus- 31 Copper Carbon