Ch. 4: Atoms and Elements

1. Ch. 4: Atoms and Elements Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry

2. I. Chapter Outline • Introduction • Atomic Theory • The Nuclear Atom • Elements • The Periodic Table • Ions • Isotopes • Atomic Mass

3. I. Introduction • Atoms are the building blocks of everything we experience. • What we smell, what we feel, what we see. • In this chapter, we trace the history of the atom and learn about its makeup.

4. II. The Greeks • From out perspective, matter can be infinitely divided. • However, Leucippus and Democritus (5th century B.C.) believed there was a limit. • Eventually, you will reach something that was “atomos” or “indivisible.” • Unfortunately, their idea was not accepted.

5. II. Revival of the Atom • The idea of the atom lay dormant for over 2000 years. • John Dalton revived the idea in order to explain 3 natural laws that puzzled everyone at the time. • Dalton’s Atomic Theory worked so well that it was quickly accepted.

6. II. Postulates of Dalton’s Theory • Each element is composed of tiny, indestructible particles called atoms. • All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements. • Atoms combine in simple, whole-number ratios to form compounds.

7. II. Atoms • Today, overwhelming evidence points towards the existence of atoms. • Atoms can be imaged and arranged!

8. III. Not “Atomos” • Dalton’s theory treated atoms as permanent, indestructible building blocks that composed everything. • However, J.J. Thomson discovered electrons, which were much smaller than an atom and negatively charged! • Since atoms are neutral, where’s the positive charge?

9. III. Plum Pudding • J.J. Thomson proposed the plum pudding model of the atom. • Electron “raisins” • “Pudding” of positive charge

10. III. Rutherford likes Plum Pudding • Ernest Rutherford was a student of J. J. Thomson. • He tried to prove the plum pudding model by shooting a-particles at gold foil. • Note that a-particles are 7000x more massive than an electron and have a positive charge.

11. III. Rutherford’s Expectation

12. III. Rutherford’s a-Particle Experiment

13. III. Conclusions from Rutherford’s Experiment • Most of an atom’s mass and all of its positive charge exist in a nucleus. • Most of an atom is empty space, throughout which tiny electrons are dispersed. • By having equal numbers of positively-charged particles (protons) and electrons, an atom remains electrically neutral.

14. III. Rutherford’s Interpretation

15. III. The Nuclear Atom • Surprisingly, an atom is mostly empty space! • The nucleus holds 99.9% of the atom’s mass.

16. III. Components of an Atom • Protons. Positively-charged particles in the nucleus. Mass of 1.67262 x 10-27 kg or 1.0073 amu. • Neutrons. Neutral particles in the nucleus. Mass of 1.67493 x 10-27 kg or 1.0087 amu. • Electrons. Negatively-charged particles. Mass of 9.1 x 10-31 kg or 0.00055 amu.

17. III. Charge • Charge is a fundamental property. • To designate charge, the sign GOES AFTER the magnitude, e.g. 2+. • Matter is charge neutral.

18. IV. An Atom’s Identity • The number of protons in an atom determines its elemental identity.

19. IV. Referring to Elements • Since each element has a unique # of protons, we could refer to elements using Z, the atomic number, which equals the # of protons in an atom. • e.g. The Z = 2 element. • More commonly, we use an element’s name or chemical symbol. • e.g. The element helium, or He.

20. IV. Chemical Symbols • Chemical symbols are a one or two letter abbreviation of an element’s name. • First letter always capitalized; second letter is LOWERCASE. • Some symbols are based on historical names: e.g. Au from aurum.

21. IV. The Periodic Table

22. IV. Sample Problem • Find the name and atomic number of the following elements. • V • N • Hg • Rh • Mo

23. V. Organizing Chemical Info • Dmitri Mendeleev was the first to organize information of elements according to periodic law, i.e. when arranged properly, elements show repeating properties.

24. V. Mendeleev’s Breakthrough • Mendeleev placed elements with similar properties in vertical columns. • He left blank spaces where he thought elements should exist.

25. V. Three Types of Elements

26. V. Sample Problem • Categorize the elements below as either a metal, nonmetal, or metalloid. • Ru (ruthenium) • Se (selenium) • I (Iodine) • Ba (barium) • Es (einsteinium) • Kr (krypton)

27. V. Main Group vs. Transition

28. V. Families of Elements

29. V. Sample Problem • To which group (new numbering system) does each of the following elements belong? If the group has a name, indicate that as well. • Br (bromine) • N (nitrogen) • Cs (cesium) • Mn (manganese)

30. VI. Atoms Can Lose/Gain e-’s • In chemical reactions, it’s common for atoms to lose or gain electrons and become ions. • ion: a particle that has a charge • Examples: • Na  Na+ + e- • I + e-  I-

31. VI. Origin of the Charge • The charge arises from the different number of protons and electrons in the atom. • Ion Charge = # protons - # electrons • A neutral Na atom has 11 protons and 11 electrons. If it loses and electron… • Ion Charge = 11 – 10 = 1+

32. VI. Cations and Anions • An ion is fundamentally different than a neutral atom, so it needs a different name. • cation: a positively-charged ion • anion: a negatively-charged ion • Note that cations and ions have different properties than their parent atoms, e.g. Na vs. Na+.

33. VI. Sample Problem • Determine the charges of the ions described below. • A chromium atom that has lost 3 electrons. • A sulfur atom that has gained 2 electrons. • An iron atom (Fe) that has 24 electrons. • A phosphorus atom (P) that has 18 electrons.

34. VI. Ions and the Periodic Table • The charge of an ion can be predicted by the position of its parent element on the periodic table IF it’s a main group element. • Simply count the number of spaces to the nearest noble gas (forward or backward). • If you go forward, it’s an anion; if you go backward, it’s a cation.

35. VI. Predicting Ion Charge

36. VI. Sample Problem • What are the ions that form from atoms of the following elements? • aluminum (Al) • tellurium (Te) • rubidium (Rb) • oxygen (O)

37. VII. Isotopes • Protons are the only thing that determines the identity of an atom. • Therefore, it’s possible for atoms of the same element to have different masses due to differing number of neutrons. • isotopes: atoms with the same number of protons, but different numbers of neutrons

38. VII. Percent Natural Abundance • The different types and amounts of each isotope is determined by nature. • Note that in an isotope, the # of neutrons varies which makes the mass number (A) vary as well.

39. VII. Referring to Isotopes • Isotopes can be represented using the A, Z, X symbol.

40. VII. Referring to Isotopes • Alternatively, the X, A notation can be used.

41. VII. Sample Problem • How many protons and neutrons are in a potassium isotope with a mass number of 39? What are the three ways to represent this isotope?

42. VIII. What’s the Mass of an Atom? • It depends! • Are we talking about the mass of a specific atom, i.e. a given isotope? • If so, it’s just approximately the mass number. • Are we talking about in general? • Then it’s more complicated…

43. VIII. Atomic Mass • Not all atoms of the same element have the same mass, but we can calculate an average. • The atomic mass is the weighted average mass of an element which accounts for all isotopes and their percent natural abundances.

44. VIII. Calculating Atomic Mass • The equation below enables calculation of atomic mass.

45. VIII. Sample Problem • Calculate the atomic mass of magnesium using the information in the table below.