Chapter 2 Atoms and Elements

1. Chapter 2Atoms and Elements

2. Dalton’s Atomic Theory • Elements are made of tiny particles called atoms. • Each element is characterized by the mass of its atoms. Atoms of the same element have the same mass, but atoms of different elements have different masses. • Chemical reactions only rearrange the way atoms are combined; the atoms themselves are unchanged.

3. Conservation of Mass • Law of Mass Conservation: Mass is neither created nor destroyed in chemical reactions.

4. Dalton’s Atomic Theory • Law of Definite Proportions: Different samples of a pure chemical substance always contain the same proportion of elements by mass. • Any sample of H2O contains 2 hydrogen atoms for every oxygen atom

5. Dalton’s Atomic Theory • Law of Multiple Proportions:When two elements form two different compounds, the mass ratios are related by small whole numbers. • It is possible for the same two elements to combine in different mass ratios to give different substances. • Example NO and NO2

6. The Discovery of Electrons • Cathode-Ray Tube (Thomson, 1856–1940): Cathode rays consist of tiny negatively charged particles, now called electrons.

7. The Discovery of Electrons • Oil Drop Experiment (Millikan, 1868–1953): Applied a voltage to oppose the downward fall of charged drops and suspend them. • Voltage on plates place 1.602176 x 10-19 C of charge on each oil drop. • Millikan calculated the electron’s mass as 9.109382 x 10-28 grams.

8. The Structure of Atoms • Discovery of Nucleus:Rutherford irradiated gold foil with a beam of alpha particles to search for positive charged particles. Most of the particles passed through but some were deflected at large angles, why?

9. The Discovery of Electrons • Thomson’s conclusions: • The cathode rays are made of tiny negatively charge particles • Every material tested contained these same particles • Thomson believed that these particles were the ultimate building blocks of matter • These particle became known as electrons

10. The Structure of the Atom • Thomson: • The structure of the atom is like plum pudding • Negatively charged particle is a sphere of positive charge

11. The Structure of the Atom • The Plum Pudding model: • The mass of the atom is due to the mass of the electrons within • The atom is mostly empty space because negative particles repel

12. Radioactivity • Curie and Becquerel discovered that certain elements would constantly emit small, energetic particles and rays • These energetic particles could penetrate matter

13. Radioactivity • Rutherford discovered there were three type of radioactive emissions • Alpha: • mass of 4x a hydrogen atom and + charge • Beta: • mass of 1/2000th of a H atom and - charge • Gamma • Energy rays not particles

14. Rutherford’s Experiment • Shoot alpha particles at a very thin sheet of matter and show that they all pass through

15. Rutherford’s Experiment • Results: • 98% of the particles went straight through • About 2% of the particles went through but were deflected by large angles • About 0.01% of the particles bounced off the gold foil

16. Rutherford’s Experiment

17. Rutherford’s Experiment • Conclusions: • The atom contains a tiny dense center called a nucleus • The atom is mostly empty space • The electron are dispersed in the empty space around the nucleus • The nucleus is positively charged • The nucleus of the atom must have a particle of the same amount of charge but opposite in sign

18. The Structure of Atoms

19. The Structure of Atoms • Atomic number ( Z)= Number of protons in atom’s nucleus= number of electrons around atom’s nucleus.

20. The Structure of Atoms Mass Number (A)=Number of protons (Z) +Number of neutrons (N)

21. The Periodic Table of the Elements

22. Protons, Electrons and Neutrons • Protons: determines the element • Number of protons = atomic number • Electrons: determines the reactivity • If neutral # of electrons = # of protons • If charged: the sum of the charged particel must equal the charge • Neutrons: determines the nuclear stability • The atomic mass (whole number) - the # of protons = the # of neutrons

23. Protons, Electrons and Neutrons • Ions: different number of electrons than neutral • Cations: have a positive charge, less electrons than neutral • Anions: have a negative charge, more electrons than neutral • Isotopes: different number of neutrons • Can be stable or radioactive • Have different atomic mass • The elemental mass on the periodic table is usually the most common isotope • The mass on the periodic chart is calculated from the sum of all the weighted masses of the naturally occurring isotopes

24. Examples: P, E, and N

25. Example: calculation of atomic mass • If copper is 69.17% Cu-63 with a mass of 62.9396amu and the rest Cu-65 with a mass of 64.9278amu, what is copper’s atomic mass?

26. Patterns

27. Patterns • Metals: shiny, ductile, malleable, conduct heat, conduct electricity • Nonmetals: dull, brittle,poor conductor of heat and electricity • Metalloids: semimetals, semiconductors, show properties of both metals and nonmetals

28. Some Important Families • Noble gases: • nonreactive

29. Some Important Families • Alkali metals: • Very reactive metals

30. Some Important Families • Alkaline Earth metals: • Reactive metals

31. Some Important Families • Halogens: • Reactive nonmetals

32. Patterns and Ions

33. Counting Atoms by the Mole • A mole is defined as a unit having 6.022x1023particles • A mole can be used to determine the number of atoms or molecules in a sample.

34. Counting Atoms by the Mole • Recall: To find the number of moles in a sample divide the number of grams by the molar mass • To find the number of particle in that sample multiply the number of moles by the Avogadro’s number

35. Example: Counting Atom • Determine the number of copper atoms in a penny weighing 3.10g.