Chapter 2: Atoms, Molecules and Ions

1. Chapter 2: Atoms, Molecules and Ions

2. Early Models of Atoms • Democritus (460-400B.C.) first suggested the existence of these particles, which he called “atoms” for the Greek word for “uncuttable”. They lacked experimental support due to the lack of scientific testing at the time. • Plato and Aristotle formulated the notion that there can be no ultimately indivisible particles, so the “atomic” view faded for a number of years. • John Dalton (1766-1844) performed experiments to study the ratios in which elements combine in chemical reactions. He formulated hypotheses and theories to explain his observations, which became Dalton’s Atomic Theory. • All elements are composed of tiny indivisible particles called atoms. • Atoms of the same element are identical. The atoms of any one element are different from those of any other element. • Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction. • Atoms of different elements can physically mix together or combine in simple, whole number ratios to form compounds.

3. Dalton’s Atomic Theory • According to Dalton’s atomic theory atoms are the smallest particles of an element that retain the chemical identity of the element. His theory explains several simpler laws of chemical combination from his time. • Law of constant composition: In a given compound, the relative numbers and kinds of atoms are constant. (4) • Law of conservation of mass: The total mass of materials present after a chemical reaction tis the same as the total mass present before the reaction (3) • Law of multiple proportions: If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. • Example: CO2 and CO H2O2 and H2O

4. Discovery of Atomic Structure • As scientists began to develop methods for more detailed probing of the nature of matter, we discovered more. Atoms are now known to be divisible as they can be broken down to even smaller particles by atom smashers. • J.J. Thomson (1856-1940) discovered electrons using cathode ray tubes. Another CRT • Robert Millikan (1868-1953) carried out experiments to determine the charge of an electron (-). He also determined the ratio of the charge to the mass of an electron. • In 1886, E. Goldstein observed a cathode ray tube and found rays traveling in the opposite direction to that of the cathode rays. He called these rays canal rays and concluded that they must be positive particles, which are now called protons. • In 1932, James Chadwick confirmed the existence of yet another subatomic particle: the neutron. Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton. See simulation

5. After discovering these subatomic particles, scientists wondered how they were put together. • JJ Thompson thought since the electrons contributed such a small fraction of the atoms mass, they were probably an equal fraction of it size so it was like “Plum Pudding”. • In 1911, Ernest Rutherford and his coworkers performed the Gold Foil Experiment to further study the phenomenon. • Concluded that most of the mass of each atom and all of its positive charge reside in a very small, extremely dense region which is called the nucleus. The rest of the atom is mostly empty space.

6. Modern View of Atomic Structure • Since the time of Rutherford, physicists have learned much about the nucleus. Although many other parts have been discovered, chemists tend to only work with three main particles since they determine chemical behavior: Electron, Neutron and Proton • Electron has a charge of -1.602 X 10-19 C and a proton has a charge of 1.602 X 10-19 C so this quantity of Coulombs is known as one electronic charge and atomic and subatomic particles usually have a charge that is multiples of this. Neutrons have no charge and are electrically neutral. • Atoms have extremely small masses so instead of using the real numbers, atomic mass units (amus) are used. Protons and neutrons are very similar in mass but it would take 1836 electrons to equal 1 proton so most of an atoms mass is in the nucleus. • Atoms are also extremely small with diameters between 1 X 10-10 and 5 X 10-10 so they are usually expressed with angstroms, which is 10-10.

7. Illustrating the Size of an Atom • The diameter of a US penny is 19 mm. The diameter of a silver atom, by comparison is only 2.88 A. How many silver atoms could be arranged side by side in a straight line across the diameter of a penny? 19 mm 1X10-3m1A1 Ag atom = 6.6 X107 Ag atoms 1mm 1X10-10m 2.88 A This is over 66 million silver atoms could sit side by side across a penny!

8. Atomic Number • The number of protons in the nucleus of an atom of that element, which is the primary difference that distinguishes each element. • For an atom with no charge, this is also the number of electrons since the positive charge of the protons cancels the negative charge of the electrons.

9. Mass Number • Most of the mass of an atom is found in the nucleus so the total number of protons and neutrons equals the mass number. • If you know the atomic number and mass number you can determine the composition of that atom. • The composition can be represented by the shorthand notation using the element symbol, atomic number and mass number. • For gold, Au is the symbol for the element and the atomic number is subscript and mass number is superscript on the left side. 197 Au 79 • Do Sample Exercises 2.3 and Practice Exercises in that box on pg 46

10. Isotopes • Atoms that have the same number of protons but different number of neutrons. • Affects the shorthand notation of the element. • Do Sample Exercises 2.3 and Practice Exercises in that box on pg 46

11. Atomic Mass • Today we can determine the masses of individual atoms with a relative high degree of accuracy but since they are so small atomic mass units are used with hydrogen being 1 amu. • The average atomic mass for an element due to the different isotopes, the mass of those isotopes and the natural percent abundance. It is also known as atomic weight. • Add up the different atomic mass of each atom and then divide by the number of atoms. • Or, multiply mass by % and then determine average mass. • Sample Exercise 2.4 and Practice Exercise on pg 47

12. Mass Spectrometer • The most direct and accurate means for determining atomic and molecular weights.See pg 48

13. Periodic Table • The arrangement of elements in order of increasing atomic number, with elements having similar properties placed in vertical column. • Atomic number, symbol, name, atomic weight are found in each square for each element. Some tables have additional information as well. Example • Can be arranged according to metals, non-metals and metalloids, solid liquid and gases, and by family. Example

14. Molecules and Molecular Compounds • Even though the atom is the smallest representative sample of an element, only the noble gas elements are normally found in nature as isolated atoms. All others form either molecules or ions. • A molecule is an assembly of two or more atoms tightly bound together by a covalent bond created by two atoms sharing electrons. • Diatomic atoms form diatomic molecules (remember 7… start at 7 form a 7 and hydrogen). • Compounds that are composed of molecules that contain more than one type of element are molecular compounds. • Most molecules are composed of nonmetals. • Chemical formulas that indicate actual number and types of atoms in a molecules are called molecular formulas. Such as H2O, C6H12O6, and C2H4. • Empirical formulas give only the relative number of atoms, they are basically the reduced formula. Such as H2O, CH2O, and CH2. • Do Sample Exercise 2.6 and Practice Exercise on pg 53.

15. Picturing Molecules • The molecular formula of a substance describes the composition but doesn’t show how they come together. • Structural formula: shows which atoms are attached to which. • Atoms are represented by their symbol and the bonds are represented by lines. • Perspective Drawing: shows actual geometry to give some sense of three-dimensional shape. • Ball-and-stick Models: shows atoms a spheres and bond as sticks. Accurately represents the angles at which the atoms are attached to one another within the molecules. • Space-filling Model: shows what the molecule would look like if the atoms were scaled up to size.

16. Ions • Some atoms can gain or lose electrons to try and get the same number of electrons as the nearest noble gas, when an electron is gained or lost from a neutral atom a charged particle occurs called an ion. • An ion with a positive charge (lost an electron) is called a cation, where as an ion with a negative charge (gained an electron) is called an anion. • In general, metals atoms tend to lose electrons to form cations and nonmetals tend to gain electrons to form anions. • In addition to simple single atom ions, there are polyatomic ions, which consist of atoms joined as a molecule but they have a net positive or negative charge. • Ionic charge can be predicted by determining how many electrons an atom has to lose to become like the nearest stably arranged noble gas. • Do Sample Exercises 2.7 and 2.8 and Practice Problems on pg 55.

17. Ionic Compounds • When a positive ion such as Na comes close to a negative ion such as Cl, their opposite charges are attracted and form an ionic compound connected by a ionic bond. • Generally, they are combinations of metals and nonmetals such as Na and Cl. • Ions in ionic compounds are arranged in three-dimensional structures. • The formula for an ionic compound is always an empirical formula (most reduced form) because there is no discrete molecule of NaCl. • Chemical compounds are always electrically neutral, so the empirical formula shows the ratio of the ions for this to be true. • For example, Mg2+ and N3- would have to be Mg3N2. • Sample Exercise 2.10 and Practice Exercises on pg 58

18. Lithium Fluoride Naming Inorganic Compounds • To obtain information about a particular substance you must know its chemical name and formula, the system used for this is chemical nomenclature. Some compounds also have common names in addition to their chemical nomenclature such as water. • The rules for naming a compound is based on divisions of substances into categories. The major division is between inorganic and organic. • Among the inorganic compounds the three basic divisions are ionic compounds, molecular compounds and acids. Sodium Nitrate Potassium Oxide Aluminum chloride

19. Naming Positive Ions (Cations) • Cations formed from metals atoms have the same name as the metal found on the periodic table. These are monatomic ions. • Mg2+Magnesium ion and K+ Potassium ion • If a metal can form different cations, the positive charge is indicated by a roman numeral in parentheses following the name of the metal. These are usually transition metals. • Cu2+Copper (II) ion and Cu+ Copper (I) ion • Cations formed from nonmetal atoms have name that end in –ium. These are polyatomic ions. • NH4-Ammonium ion and H3O+  Hydronium ion

20. Naming Negative Ions (Anions) • The names of monatomic anions are formed by replacing the ending of the name of the element with –ide. • O2-oxide ion and N3- nitride ion • Polyatomic anions containing oxygen have names ending in -ate or -ite. • NO3-  nitrate ion and NO2- nitrite ion • Anions derived by adding H+ to an oxyanion are name by adding hydrogen or dihydrogen as a prefix as appropriate. • HCO3-hydrogen carbonate ion and H2PO4- dihydrogen phosphate ion

21. Naming Ionic Compounds • Names of ionic compounds consist of the cation name followed by the anion name. • Do Sample Exercises 2.12 and 2.13 and Practice Exercises on pg 63

22. Naming Acids • You know a molecule is an acid because its cation is hydrogen. • Acids containing anions whose names end in -ide are named by changing the -ide ending to -ic adding the prefix hydro- to this anion name and then following with the word acid. • HClHydrochloric Acid • Acids containing anions whose name end in -ate or -ite are named by changing the -ate to -ic and -ite to -ous and then adding the word acid. • HNO3Nitric Acid & HClO2Chlorous Acid • Do Sample Exercise 2.14 and Practice Exercise on pg 64-65.

23. Naming Binary Molecular Compounds • The name of the element farther to the left in the periodic table is usually written first. Except Oxygen is written last with all except Flourine. • If both elements are in the same group in the periodic table, the one having the higher atomic number is named first. • The name of the second element is given and -ide ending. • Greek prefixes are used to indicate the number of atoms of each element. Although mono is never used with the first element. • mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca • N2O5 Dinitrogen pentoxide • Sample Exercise 2.15 and Practice Exercise on pg 65