Atomic Theory of Matter: Atoms, Molecules, and Ions

1. Chapter 2Atoms, Molecules, and Ions Krishna Trehan AP Chemistry 2008-2009 Mrs.Molchany

2. 2.1 - The Atomic Theory of Matter • The idea of atoms was formulated when scientists tried to explain properties of gasses. • Gas composed of invisible objects in constant motion. • Isaac Newtown proposed the idea that atoms were the chemical building blocks of nature.

3. Daltons Atomic Theory • 1. Each element is composed of small particles called atoms • 2. All atoms of a given element are identical • 3. Chemical reactions do not alter atoms, they can be neither created nor destroyed. • 4. Compounds are formed when atoms of more than one element combine.

4. Daltons Theory • Atoms are the most basic building blocks of matter that retain chemical identities of elements. • Law of constant composition – in a compound the relative numbers and kinds of atoms are constant. • Law of conservation of mass – the total mass of materials present is the same before and after a reaction occurs.

5. Daltons Theory • Law of multiple proportions – When elements combine, they do so in the ratio of small whole numbers. i.e. When Carbon and Oxygen react, the produce CO or CO₂ and not something like CO1.6

6. 2.2 The Discovery of Atomic Structure • The atom is composed of subatomic particles. • There are only two kinds of charges : positive, and negative • Law of Electrostatic Attraction like charges repel one another, unlike charges attract.

7. Cathode Rays and Electrons • An tube is pumped almost empty of air. High voltage produces radiation, which is called a cathode ray, because it comes from the negative electrode, or cathode. • The rays could not be seen, but detected. Depending on which gas is in the tube, the ray will give off a certain colored light.

8. J.J.Thompson on Cathodes and Rays • J.J. Thompson concluded that cathode rays are not waves, but particle masses. • Discovered the ratio of the electrons charge to mass as 1.76 x 10⁸ coulombs per gram. He did this by bending the path of a cathode using magnets. By observing the magnitude of the path change, he was able to derive this number.

9. Robert Millikan • Robert Millikan tried to find the mass of an electron by using the Oil Drop Experiment • Millikan allowed small drops of oil (which had obtained extra electrons) to fall on top of electrically charged plates. • From these observations he found that the charge on an electron was

10. Millikan • By using J.J. Thomson’s value for the electrical charge to mass ratio, Millikan found the presently accepted mass of the electron.

11. Ernest Rutherford • Rutherford discovered three types of radiation: alpha, beta, and gamma. • Alpha and beta are bent by electric fields, while gamma rays aren't. • Alpha: are larger than beta rays, and have a positive charge (attracted to negatively charged plates). • Beta: have a negative charge (attracted to positively charged plates. • Considered the radioactive equivalent of cathode rays.

12. The Nuclear Atom • Thompson proposed the idea that since electrons were relatively small, they held a small fraction in the total area of the atom. • He then made the “Plum Pudding” model where individual negatively charged electrons were spread throughout a positive sphere.

13. Gold Foil Experiment • Thompson sent a beam of alpha particles through a thing piece of gold foil. He discovered that all of the alpha particles passes through. But after reviewing the experiment, he saw that some alpha particles were deflected in other directions, some even bounced back. This contradicted the Plum Pudding model.

14. Gold Foil Experiment • He postulated from this experiment that all of the positive charge in the atom was concentrated in a small dense region, called the nucleus. • This meant that most alpha particles simply passed the nucleus, but the ones that bounced back, hit the nucleus and were repelled by its positive charge.

15. 2.3 The Modern View of Atomic Structure • Charge of an electron is -1.602 x 10 -19 C • Charge of a proton is 1.602 x 10 -19 C • 1.602 x 10 -19 is called the electronic charge. • C = coulomb • Atoms have the same number of electrons and protons, so they have no net electrical charge

16. Atomic Structure • Protons and Neutrons reside in the nucleus • A large majority of an atoms area is outside the nucleus, where the electrons are. • Atoms have very little mass, so we use the atomic mass unit (amu) to measure mass. • One amu is equivalent to: • A proton is 1.0073 amu and a neutron is 1.0087 amu • Angstroms (Å), which are 10-10 m, are used to measure diameter

17. Isotopes, Atomic Numbers, and Mass Numbers • Protons are what make elements unique • All atoms of an element have the same number of protons in the nucleus. • Isotopes areatoms of a given element that differ in the number of neutrons. • Atomic number – the number of protons. • Mass number – total number of protons and neutrons in an element

18. 2.4 The Periodic Table • The arrangement of elements in order of increasing atomic number, with elements having similar properties placed in vertical columns is known as the periodic table. • Each column is a group. • Elements in the same group and have similar physical and chemical properties.

19. Metallic Elements • All the elements on the left side and in the middle of the periodic table are metals. • Metals are generally lustrous and are good conductors of heat and electricity. • All metals (besides mercury Hg) are solid at room temperature.

20. Nonmetallic Elements • The metals are separated for nonmetals by a diagonal staircase line that runs from boron (B) to astatine (At). • Hydrogen, though a nonmetal, is on the left side of the periodic table. • The state of matter at which nonmetals are at room temperature, vary from element to element.

21. Metalloids • Elements that have the characteristics of both metals and nonmetals are called metalloids. • They can also be called semi-metals.

22. 2.5 Molecules and Molecular Compounds • Molecule – an assembly of two or more atoms tightly bound together • Diatomic – molecule composed of two atoms, both of which are the same element. • Molecular compounds arecompounds composed of molecules. • For example, in water there are two hydrogens, and one oxygen. The two hydrogens are a diatomic molecule.

23. Molecular and Empirical Formulas • Molecular Formulas – chemical formulas that indicate the actual numbers and types of atoms in a molecule. • Empirical Formulas – chemical formulas that give only the relative number of atoms of each type in a molecule. • Always show the smallest whole number ratio • i.e. C₂H₄ = CH₂ • Molecular formulas provide more/accurate information about a molecule

24. Picturing Molecules • Structural formula - shows which elements are attached to which in a chemical formula. • Generally, structural formulas do not represent actual images, but simply provide a sketch of the structure. H₂O H₂O₂

25. 2.6 Ions and Ionic Compounds • If an electron is added or removed from an atom, then the atom now has a charge, or has become an ion. • Cation = positively charged = loss of electron • Anion = negatively charged = gain of electron • Polyatomic ion – ions that consist of atoms joined as in a molecule, but have a net positive/negative charge : NO₃⁻1 or SO₄⁻2

26. Predicting Ionic Changes • Atoms gain or lose electrons so as to end up with the same number of electrons as the noble gas closest to them. • The noble gases are non-reactive and form few compounds.

27. Ionic Compounds • Ionic compound - compounds that contains positively charged ions and negatively charged ions • When sodium and chlorine react, an electron from sodium goes to chlorine which leaves us with Na⁺+ Cl⁻ After that, the opposite charges attract one another and Na and Cl are held together. • Ionic compounds are combinations of metals and nonmetals. • Molecular compounds are generally composed of nonmetals only.

28. Chemical Compounds • Chemical compounds are always electrically neutral, so we need to make sure the ions in an ionic compounds occur in a ratio where the total positive charge is EQUAL to the total negative charge. Mg2+ + N3- Mg₃N₂ 3 (number of Mg) x 2 (charge of Mg) = 6 2 (number of N ) x -3 (charge of N) = -6 -6 + 6 = 0 which means we have an overall neutral charge.

29. 2.7 Naming Inorganic Compounds Chemical Nomenclature On page 56 of the old textbook you can find a list of many different cations. On page 57 of the old textbook you can find a list of many different anions. Most of these are found on the pink sheets we got at the beginning of the year.

30. Positive Ions (Cations) • Cations formed from metal atoms have the same name as the metal. Example : Na⁺ = sodium ion • IF a metal can form cations of differing charges, the positive charge is given by a Roman numeral in parentheses following the name of the metal: Example : Fe2+ = iron(II) ion Fe3+ = iron(III) ion

31. Most elements that have multiple charges are transition metals • An older form of naming elements was using the latin name (these are found on the pink sheets Mrs.Molchany gave us) • Fe2+ = ferrous ion Fe3+ = ferric ion • Cations formed from nonmetal atoms have names that end in –ium. NH₄⁺ = ammonium ion H₃0⁺ = hydronium ion

32. Negative Ions (Anions) • Monatomic anions have names formed by dropping the ending of the name of the element and adding the ending –ide. H⁻ = hydride ion CN⁻ = cyanide ion. • Polyatomic anions containing oxygen have names ending in –ate or –ite. NO₃⁻1 = nitrate NO₂⁻1 = nitrite

33. Negative Ions (Anions) (cont.) • Anions derived by adding H⁺ to an oxyanion are named by adding as a prefix the word hydrogen or dihydrogen. CO₃2- (carbonate ion)  HCO₃⁻ (hydrogen carbonate ion)

34. Ionic Compounds • Naming ionic compounds follow the format: • cation name followed by the anion name BaBr₂ = barium bromide Al(NO₃)₃ = aluminum nitrate Cu(ClO₄)₂ = copper (II) perchlorate

35. Names and Formulas of Acids • Acids based on anions whose names end in –ide. Anions whose names end in –ide have associated acids that have the –hydro prefix and an –ic. • Cl⁻(chloride)  HCL (hydrochloric acid) • Acids based on anions whose names end in ite or ate. Anions whose names end in ate have acids ending in ic, while those ending in ite, have acids ending in ous. • ClO⁻ (hypochlorite) = HClO (hypochlorous acid)

36. Names and Formulas of Binary Molecular Compounds 1. The name of the element farthest to the left in the periodic table is written first 2. If both elements are in the same group, the lower one is named first 3. The name of the second element is given an –ide ending. 4. Greek prefixes are used to indicate the number of each element (ie. Mono = 1) Cl₂O = dichlorine monoxide