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Modern Chemistry Chapter 3

Modern Chemistry Chapter 3. Atoms: The Building Blocks of Matter. Chapter 3 Sec 1. The Atom: From Philosophical Idea to Scientific Theory In Ancient Greece (400B.C.) Democritus was the first on record ever to say matter was composed of atoms Atom is based on the Greek for indivisible.

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Modern Chemistry Chapter 3

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  1. Modern ChemistryChapter 3 Atoms: The Building Blocks of Matter

  2. Chapter 3 Sec 1 The Atom: From Philosophical Idea to Scientific Theory • In Ancient Greece (400B.C.) Democritus was the first on record ever to say matter was composed of atoms • Atom is based on the Greek for indivisible

  3. More Ch 3-1 • Aristotle who was next generation form Greece did not believe in atoms……. • It was not until 18th Century that experimentation began to gather evidence to prove the existence of atoms

  4. Foundations of Atomic Theory • By the late 1700’s the definition of an element became widely accepted • An Element is something that cannot be broken down by ordinary chemical means. • It was also commonly believed that elements combined to form compounds.

  5. More Fundamentals of Atomic Theory • Chemical reactions were defined as the transformation of a substance or substances into one or more new substances. • By the 1790’s the analysis of matter became more quantitative because reliable balances were becoming available.

  6. Law of Conservation of Mass • This law states that mass is neither created or destroyed during ordinary chemical reactions or physical changed. • Back in the day people had funny ideas…..Spontaneous generation, magic, potions, alchemy

  7. Law of Definite Proportion • This law states that a chemical compound contains the same elements in exactly the same proportion by mass regardless of the size of the sample or the source of the compound. • Example: water one drop vs gallons • Example: water lake, pond, ice

  8. Law of Multiple Proportions • This law states that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers • Example: hydrogen peroxide and water • H2O2H2O

  9. Dalton’s Atomic Theory • 1. All matter is composed of extremely small particles called atoms. • 2. Atoms of a given element are identical in size, mass, and other properties. • 3. Atoms cannot be subdivided, created or destroyed. • 4. Atoms of different elements combine in simple whole-number ratios to form compounds • 5. In chemical reactions, atoms are combined, separated, or rearranged.

  10. Modern Atomic Theory • Much of John Dalton’s Theory is still true today….. • All matter is composed of atoms • Atoms of any one element differ in properties from atoms of another atom. • Some of Dalton’s theory has been discarded: isotopes, nuclear reactions, etc…..

  11. Chapter 3 Sec 2 Structure of the Atom • Atom is defined as the smallest particle of an element that retains the properties of that element. • Today it is known that the atom has two basic parts: the nucleus and the electron cloud. The three basic subatomic particles are electrons, protons and neutrons

  12. Discovery of the Electron • Scientists of the late 1800’s were trying to discover the relationship between matter and electricity by passing electric current through various gasses at low pressures. • These devices were called cathode ray tubes. • It was noticed that when current passed through the tube the cathode glowed. They hypothesized that this was due to a stream of particles.

  13. The Electron Two main observations: • 1. Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have an negative charge. • 2. The rays were deflected away from a negatively charged object.

  14. The Electron • J.J. Thomson 1897 said the ratio of charge to mass was always the same regardless of the type of metal used or type of gas used. • He named this particle the “ELECTRON” • In 1909 American physicist Robert Millikan measured the charge of an electron. • Today we know the mass 9.109*10-31 kg

  15. Discovery of the Atomic Nucleus • J.J. Thomson believed electrons were imbedded throughout the atom ”plum pudding” model of the atom. • In 1911, Ernest Rutherford, Hans Geiger, Ernest Marsden conducted the Gold Foil Experiment. Bombard thin gold foil with alpha particles • They are given credit for discovering the nucleus of the atom. They concluded there was densely packed matter with a positive charge present in the gold foil = the nucleus • See table on top of page 76.

  16. The Nucleus • Analogy if the nucleus is a marble the atom would be the size of a football field • Electron negative charged equal but opposite of the proton positive charge • Neutrons are not charged but have a mass slightly larger than the proton.

  17. Composition of the Atomic Nucleus • Nuclear Forces are short range proton-neutron, proton-proton, neutron-neutron forces holding nuclear particles together. • The size of the atom is extremely small! Atomic radii range from 40 to 270 picometers ( reminder that is 10-12 of a meter) • Nuclei are extremely dense about 2*108 metric tons/cm3

  18. Chapter 3 Sec 3 Counting Atoms • Atomic number is equal to the number of protons in an atoms nucleus • Isotopes different forms of the same element with differing masses due to differing neutrons • Mass Number is the total number of protons and neutrons added together • The Mole Avogadro’s number = 6.02x10 23 • Molar Mass

  19. More Ch 3 sec 3 • The MOLE Avogadro’s number of molecules. • Molar Mass (always read Periodic Table amu’s to a hundredth of a gram) • Atoms and molecules are soooooooo small that the scale we use them on is huge!

  20. Molar Mass • With a Periodic Table and calculator by your side!  • Molar mass of an element is just read off the Periodic table and rounded correctly! • Molar mass for a compound is added using masses from the periodic Table and the subscripts in the formula.

  21. Example: Molar Mass • H2O water • Molar mass • Hydrogen (1.01 grams) 2 • Oxygen (16.00 grams) 1 • 2.02g + 16.00g= 18.02grams ANSWER!

  22. Example: Molar Mass • C6H12O6 glucose • Carbon (12.01g)6= 72.06g • Hydrogen(1.01g)12= 12.12g • Oxygen (16.00g)6= 96.00g • And the total is!!!!!!! 180.18 grams is the molar mass of glucose

  23. Molar Mass • Aluminum nitrate Al(NO3)3 Molar mass • Aluminum (26.98g) 1 = 26.98 • Nitrogen (14.01g) 3 = 42.03 • Oxygen (16.00g) 9 =144.00 • And the total is !!!!! 213.01 grams

  24. MOLE Dimensional Analysis • Amadeo Avogadro said one mole is equal to 6.02x1023 particles……. • 6.02 EE 23 or 6.02 exp 23 don’t type x10……….. If you have any questions about your calculator please see me! • 2 mole is 12.04 x10 23 particles • .5 mole is 3.01 x10 23 particles • Mole calculations are FUN!

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