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Plan for Wed, 24 Sept 08

Plan for Wed, 24 Sept 08. Lecture Classifying and separating mixtures (1.9) How do we know that there are atoms? (2.2) Ok, there are atoms. What do they look like? (2.5) Re-introducing the periodic table (2.6-7) Chemical nomenclature (2.8) Make sure to read: 2.3 (Dalton’s atomic theory)

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Plan for Wed, 24 Sept 08

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  1. Plan for Wed, 24 Sept 08 • Lecture • Classifying and separating mixtures (1.9) • How do we know that there are atoms? (2.2) • Ok, there are atoms. What do they look like? (2.5) • Re-introducing the periodic table (2.6-7) • Chemical nomenclature (2.8) • Make sure to read: • 2.3 (Dalton’s atomic theory) • Sections you can skip: • 2.1 (early history of chemistry), 2.4 (early experiments to characterize the atom)

  2. Pure Substances vs Mixtures Pure substances cannot be separated into different components by physical means. Mixtures can be separated into different components. The unique properties of the components of a mixture can be used to separate the components.

  3. Homogeneous vs Heterogeneous Mixtures A Homogeneous Mixture is a mixture that has a uniform appearance and composition throughout A Heterogeneous Mixture is a mixture in which you can identify the component parts just by looking at it Homogeneous mixtures are often indistinguishable from pure substances on the macroscopic level. You have to consider an unknown sample on the molecular level to determine if it is a mixture or pure substance.

  4. Separation Techniques • Since components in a mixture retain their identities, we can exploit the unique properties of the components to separate them. • The more similar the properties are, the more difficult it is to separate the components.

  5. Separation of Heterogeneous Mixtures...by particle size Cacao beans grow in the jungles of equatorial countries, many of which are rather dangerous. Before chocolate manufacturers can roast the cacao beans from their suppliers, they must remove sticks, stones, and sometimes bullets!! freshly harvested cacao beans plus some other junk chocolate fountain! roasted cacao beans

  6. Separation of Heterogeneous Mixtures…by magnetism In this example, the magnetic property of the iron filings is used to separate it from the non-magnetic sulfur powder.

  7. Physical Properties of Homogeneous Mixtures • Pure substances have a unique set of physical properties that are different from any other pure substance • The physical properties of mixtures are not unique...they differ with the relative amounts of the components in the mixture Change in Boiling Point of a Solution vs. a Pure Liquid

  8. Separating Liquid-Liquid Mixtures Change in Boiling Point of a Solution vs. a Pure Liquid

  9. Ch 2 – Atoms Molecules and Ions What did we know in the 1800’s? • Most natural materials are mixtures of pure substances. • Pure substances are either elements or combinations of elements called compounds. • A given compound always contains the same proportions (by mass) of the elements...this is the law of definite composition

  10. Law of Definite Composition A given compound always has the same composition, regardless of where it comes from. ...or, a chemical compound always contains exactly the same proportion of elements by mass. Water: 8 g oxygen (O) to 1 g hydrogen (H) Carbon dioxide: 2.7 g oxygen (O) to 1 g carbon (C)

  11. Law of Multiple Proportions When two elements combine to form more than one compound, the different weights of one element that combine with the same weight of the other element are in a simple ratio of whole numbers. What this means at the particulate level is that when elements combine, they do so in the ratio of small whole numbers. For example: carbon and oxygen react to form CO or CO2, but not CO1.8. Fig. 5-2, p. 121

  12. Negatively-charged electrons Uniform positive charge Ok, there are atoms. What do they look like? • It was known that atoms were neutrally-charged, but that they contained negatively-charged particles called electrons. • Lord Kelvin (yes, that Lord Kelvin) proposed the “Plum-Pudding” Model (we will call it the “Chocolate-Chip Cookie” Model) of atomic structure. • In this model, the atom is composed of discrete, negatively-charged electrons embedded in a cloud of uniform positive charge.

  13. The Nuclear Atom 1911 – Ernest Rutherford demonstrated the nuclear nature of the atom in which the empty space is 10,000 to 100,000 times larger than the size of the nucleus.

  14. Atomic Structure • Every atom contains small, dense nucleus. • All of the positive charge and most of the mass are concentrated in the nucleus. • The nucleus is surrounded by a large volume of nearly empty space that makes up the rest of the atom. • The rest of the atom is thinly populated by electrons, the total charge of which exactly balances the positive charge of the nucleus. If an atom was the size of a baseball stadium, the nucleus would be the size of a fly on home plate.

  15. What’s in an Atom? • Electron • mass = 9.11 x 10-28 g 0.000549 amu (call this 0 amu) • charge = -4.8080 x 10-10 esu -1 • Proton • mass = 1.67 x 10-24 g 1.00728 amu (call this 1 amu) • charge = +4.8080 x 10-10 esu +1 • Neutron • mass = 1.68 x 10-24 g 1.00867 amu (call this 1 amu) • charge = 0 esu 0 (where esu = electrostatic unit; amu = atomic mass unit) 1 amu = 1.66 x 10-24 g 1 amu = 1/12 the mass of one carbon atom

  16. What do the numbers of different particles mean? • # Protons = chemical identity of the atom (which element is it?) • In an electrically-neutral atom, the number of protons in the nucleus is exactly balanced by the number of electrons. • # Electrons = ionic character of the atom. An ion has either more or fewer electrons than the electrically-neutral atom. • anion = more electrons, so ion is negatively-charged • cation = fewer electrons, so ion is positively-charged • # Neutrons = isotopic character of the atom • an atom of an element usually comes in at least 2 or 3 different isotopes (sometimes more) • usually there will be one isotope that is far more abundant than the others • If the number of protons is changed, the chemical identity of the atom is changed.

  17. Atomic number (Z) Atomic symbol (X) 16 16 O O 8 O 8 Atomic Weight (related to A) A X Z 16.00 A few definitions. . . • Atomic number (Z): the number of protons in the nucleus of an atom • Mass number (A): the sum of the numbers of protons and neutrons in the nucleus of an atom • Atomic Mass: the mass of one atom, expressed in amu • Atomic Weight: an average of the atomic masses of the most common isotopes Atomic Symbol Element symbol (X) For example: or In the periodic table...

  18. 60 Co 27 37 Cl 17 238 U 92 Let’s count some particles # protons # electrons # neutrons 27 27 60 – 27 = 33 Cobalt-60 Chlorine-37 17 17 37 – 17 = 20 Uranium-238 92 92 238 – 92 = 146

  19. QUESTION Of the following three choices, which would have the greatest number of neutrons? 1. 137Ba2+ 2. 128Te2– 3. 133Cs A = 137 Z = 56 # neutrons = 137 – 56 = 81 A = 128 Z = 52 # neutrons = 128 – 52 = 76 A = 133 Z = 55 # neutrons = 133 – 55 = 78

  20. QUESTION Of the following, which would NOT qualify as an isotope of 35Cl? 1. 36Cl 2. 35Cl– 3. 37Cl– Not an isotope because # neutrons is the same.

  21. QUESTION Calcium plays several critical roles in the functioning of human cells. However, this form of calcium is the ion made with 20 protons and 18 electrons. Therefore the ion would be… 1. positive and called an anion. 2. positive and called a cation. 3. negative and called an anion. 4. negative and called a cation.

  22. 1A 8A Group 3A 6A 7A 4A 5A 2A Meet the Periodic Table Period Hello!! • Alkali Metals ... soft, shiny metals; react vigorously with water; rarely found in elemental form • Alkaline Earth Metals ... soft, shiny metals; react less vigorously with water than alkali metals; rarely found in elemental form • Halogens ... gases: F, Cl; liquid: Br; solid: I; highly reactive; F is the most reactive element; all quite toxic; not found in elemental form • Noble Gases ... all gases; largely unreactive, although Kr and Xe can form compounds; found in minute quantities in the atmosphere

  23. 1A 8A B 3A 6A 7A 4A 5A 2A Si As Ge Te Sb At Po Non-metals • Metals … good conductors of heat, electricity; malleable solids. Tend to lose electrons in reactions to form cations. • Non-metals … poor conductors; not malleable. Tend to gain electrons in reactions to form anions. • Metalloids … both metallic and nonmetallic properties Metals

  24. QUESTION From the following list select the element that is most likely to become an anion during a chemical reaction. 1. Hydrogen 2. Tungsten 3. Germanium 4. Bromine Nonmetals tend to gain electrons in reactions.

  25. QUESTION Of the following, which is most likely to become a cation as a result of a chemical reaction? What would be the charge on that cation? 1. N; –3 2. Ne; +1 3. Na: +1 4. Not enough information given to predict. Metals tend to lose electrons in reactions.

  26. Chemical Nomenclature Outline • Chemical formulas of elements • Naming binarycompounds (two elements) • compounds containing a metal and a nonmetal, aka ionic compounds • Type I: metal forms one kind of cation (we have seen these before) • Type II: metal can form more than one kind of cation • compounds containing two nonmetals • Naming compounds that contain polyatomic ions • Naming Acids and their anions • acids that do not contain oxygen typically give monatomic ions (exceptions include CN-) • acids that contain oxygen (“oxyacids”) always give polyatomic ions

  27. Formulas of Elements • The chemical formula of most elements is the atomic symbol, e.g. Li, Os, Xe, Pu • Some elements form diatomic molecules. Memorize these: H2, N2, O2, F2, Cl2, Br2, I2 Horses Need Oats For Clear Brown I’s • Other exceptions: P4, S8, C60

  28. Binary Ionic Cmpds (Type I) Metal can form only one kind of cation Note: These cations are not isoelectronic with any noble gas!! Ni2+ Ga3+ Zn2+ Ag+ Cd2+

  29. Binary Ionic Cmpds (Type I) Cations (Mn+): name of atom + cation • Magnesium: Mg2+ ... magnesium cation • Cesium: Cs+ ... cesium cation Anions (Xm-): root of atom name + -ide • Fluorine: F- ... fluoride anion • Sulfur: S2- ... sulfide anion • Selenium: Se2- ... selenide anion Compound Name: <cation name> <anion name> Chemical Formula: MmXn

  30. Example • Name the following compound: KCl potassium chloride • Name the following compound: CaBr2 calcium bromide • Write the formula for: barium hydride BaH2 • Write the formula for: aluminum sulfide Al2S3

  31. Binary Ionic Cmpds (Type II) Metal can form more than one kind of cation (see table 6.7) These are a few common examples of metals that can form different cations. This is not an exhaustive list.

  32. Binary Ionic Cmpds (Type II) Cations (Mn+): name of atom + (oxidation state) + cation • Oxidation state = charge. Use Roman numerals. • Iron: • Fe2+ ... iron(II)cation • Fe3+ ... iron(III)cation • Chromium: • Cr2+ ... chromium(II)cation • Cr3+ ... chromium(III)cation Anions (Xm-): root of atom name + -ide (same as before) Compound Name: <cation name> <anion name> Chemical Formula: MmXn

  33. Example • Name the following compound: CuCl copper(I) chloride • Name the following compound: MnO2 manganese(IV) oxide • Write the formula for: vanadium(V) fluoride VF5 • Write the formula for: tin(IV) bromide SnBr4

  34. The first element in the formula is named first, and the full element name is used. The second element is named as though it were an anion. Prefixes are used to indicate the number of atoms present. The prefix mono- is never used for naming the first element. i.e., CO is carbon monoxide, not monocarbon monoxide. Binary Covalent Cmpds (Type III) (two nonmetals, or a nonmetal and a metalloid)

  35. Example • Name the following compound: BF3 boron trifluoride • Name the following compound: I2O7 diiodine heptoxide • Write the formula for: phosphorus trichloride PCl3 • Write the formula for: dinitrogen trioxide N2O3

  36. Table 2.5 Common Polyatomic Ions

  37. Ionic Compounds with Polyatomic Ions • Writing names and formulas is pretty much the same as for binary compounds… • cation is named first, anion is named second • multiples of the ions are taken to ensure charge neutrality • Example: • Na2SO4 … 2Na+ + SO42- sodium sulfate • manganese(II) hydroxide … Mn2+ + 2OH- Mn(OH)2 • The tricky part is learning where one ion ends and the next begins. • Example: • KHSO4 … K+ + HSO4- potassium hydrogen sulfate • NH4C2H3O2 … NH4++ C2H3O2- ammonium acetate • NaH2PO3 ... Na+ + H2PO3- sodium dihydrogen phosphite • Lucky for you guys, there are only a few common polyatomic cations: NH4+ and Hg22+

  38. Polyatomic ion(s) present? No Yes No Naming procedure is similar to naming binary ionic compounds. Take organic chemistry Flowchart for Naming Compounds

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