1. Chapter 2—Atoms, Molecules, and Ions Chemistry: Principles and Reactions Fall 2007

2. The building blocks of matter: Atoms: once thought smallest; we know now that they’re made of electrons, protons, and neutrons Molecules: groups of atoms bonded together in a specific way (formula) Ions: charged atoms or molecules (+ or -)

3. Atomic theory Dalton developed the atomic model in 1808. –An element is composed of tiny particles called atoms –In a reaction, atoms move from one substance to another, but no atom disappears or is changed into another element –Compounds are formed when atoms of 2 or more elements combine

4. Were atoms the smallest? Was the question for the next ≈ 100 years Thomson and Rutherford confirmed the existence of subatomic particles. In 1897 Thomson carried out an experiment involving a cathode tube—this proved the existence of the electron (1/2000 th of an atom, charge of -1)

5. The proton: In 1911, Rutherford proved the existence of the proton with his now- famous gold-foil experiment He shot α-particles at the foil…most passed through, but some were reflected back at the fluorescent screen

6. The gold foil experiment “It was as though you had fired a 15-inch shell at a piece of tissue paper, and it bounced back and hit you.” --Ernst Rutherford

7. What did it prove? Since α-particles were + charged, there must be some + charged portion of the atom Rutherford reasoned that a nucleus with a + charge existed at the center of the atom…the particles that passed through the foil missed it…the particles that bounced back had struck it.

8. The weird conclusion Most of the atom (and most of the universe) is empty space) Since Rutherford, we have learned that the nucleus contains two types of particles: –+ charged protons almost as big as a hydrogen atom –Neutral particles called neutrons, bigger than a proton

9. Most of the mass of an atom, then, is in the neutron, although the volume is much larger than the neutron

10. Atomic number The thing that makes an element an element is its atomic number (the number of protons) The numbers on the periodic table correspond to the number of protons in the atoms of that element In a neutral atom, the protons and electrons are equal in number

11. Mass number The mass number of an element is determined by adding up the protons and neutrons (electrons are so small they’re negligible) Atoms of the same element (same atomic number) can differ in mass, because the number of neutrons can vary. We call these isotopes

12. How to write it out In the example to the left, the mass number is the top number. The atomic number is the bottom number. How do you determine the number of neutrons from this information? What about the number of electrons? See example 2.1 for clarification

13. What is radioactivity? Some isotopes are stable, and do not decay. Some are unstable, and break down into nuclei of other elements. Stable isotopes have a neutron / proton ratio falling between 1:1 and 1.5:1

14. What happens if they fall apart? Three things are emitted: –Beta particles similar to electrons –Alpha particles, which are heavy helium nuclei with a +2 charge –Gamma rays, which are the things that made Spider-man, the Hulk, and many other superheroes

15. Horizontal rows are known as periods. There are six of them Vertical columns are known as groups. There are 18 of them. The Periodic Table

16. Names to know: Group 1: Alkali Metals Group 2: Alkaline Earth Metals Group17: Halogens Group 18: Noble Gases In general, metals on the left and middle, nonmetals on the right, metalloids in between them: look at the stairway in your book and know it.

17. The table in one sentence The periodic table is an arrangement of elements in horizontal order of increasing atomic number and vertical order of chemical similarity The modern periodic table got its start with Dmitri Mendeleev, who first began grouping elements in order of similar chemical properties—and predicted the discovery of elements that were later found

18. What makes a molecule 2 or more atoms sticking together Most often through covalent bonding, which means they share electrons. We represent molecules through their formula (NH 3 ) or their structure: HNH H

19. What is an ion? An ion is an atom with too few or too many electrons Too few makes it a cation (+ charge) Too many makes it an anion (- charge) Usually, metals lose electrons (+) Usually, nonmetals gain electrons (-) The important point---protons never ever go away.

20. How they join When a metal and nonmetal bond, the metal gives up its electrons to the nonmetal, but the neutral charge is maintained: Na and Cl come together this way: Na gives away an electron (Na +) Cl picks that electron up (Cl-) So what is the total charge on NaCl?

21. Ionic compounds Ionic compounds are formed by an ionic bond between a metal and nonmetal. The metal gives away electrons, and the nonmetal picks them up (thus making them both ions There are very simple rules for how this happens— column 1 gives up 1 electron; column 2 gives up 2 Column 17 accepts 1 electron; column 16 accepts 2; column 15 generally accepts 3 In general, ions are always trying to get to 8 electrons…noble gas structure

22. Polyatomic ions Polyatomic ions are charged units that act together in reactions. The most common are listed on page 41 in table 2.2. You are responsible to know their formulae, names, and charges by one week from today.

23. Naming Compounds—42-44 The simplest way is by formula Cationstake the name of the metal from which they come (sodium= Na+ Metals in the middle can oxidize more than one way—they get a roman numeral listing the charge. Fe 2+ = iron (II) Polyatomic ions are memorized Monatomic anions get the suffix –ide attached

24. Ionic compounds are named as follows: first metal + second metal- suffix Sodium chloride (NaCl) Chromium (III) Nitrate –Cr(NO 3 ) 3

25. Molecular compounds Nonmetals and nonmetals don’t form ionic bonds…they have different rules We use the Greek prefix system (page 43), and attach the prefix to each element N 2 O 5 = dinitrogen pentoxide

26. Some just have names Water = H 2 O Ammonia = NH 4 Acetylene = C 2 H 2 Nitric / Nitrous Oxide = NO / N 2 O

27. The last 2 systems Hydrocarbons—handled in class Acids—handled in class