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Unformatted text preview: AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own The Atomic Theory I. Democritus: 5th century B.C. Greek philosopher. A. Expressed belief that all matter consists of very small, indivisible particles called atomos (meaning uncuttable or indivisible). II. John Dalton: English scientist and school teacher that formulated a precise definition of the indivisible building blocks of matter called atoms. A. Elements are composed of extremely small particles called atoms. B. All atoms of a given element are identical, having the same size, mass, and chemical properties (Figure 1). C. The atoms of one element are different from the atoms of all other elements (Figure 1). D. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction (Figure 1). 1. Law of Definite Proportions: different samples of the same compound always contain its constituent elements in the same proportion by mass. a. If the ratio of the masses of different elements in a given compound is fixed, the ratio of the atoms of these elements in the compound also must be constant. 2. Law of Multiple Proportions: if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers (Figure 2). a. Different compounds made up of the same elements differ in the number of atoms of each kind that combine. E. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. 1. Law of Conservation of Mass: matter can be neither created nor destroyed. a. Because matter is made of atoms that are unchanged in a chemical reaction, it follows that mass must be conserved as well. Figure 1 Figure 2 ~1~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own The Structure of the Atom I. Atom: basic unit of an element that can enter into chemical combination; extremely small and indivisible; made up of subatomic particles (protons, neutrons, and electrons). II. The Electron A. Radiation: emission and transmission of energy through space in the form of waves. B. Cathode Ray Tube: glass tube from which air has been evacuated (Figure 3). 1. When the two metal plates are connected to a high-voltage source, the negatively charged plate (cathode) emits an invisible ray. 2. Cathode ray is drawn to positively charged plate (anode) where it passes through a hole and continues traveling to the other end of the tube. 3. When the ray strikes the specially coated surface, it produces a strong fluorescence. C. Experiments with Cathode Ray Tube 1. When ± electrical field is on, ray is repelled from the ± charge. 2. When + electrical field is on, ray is attracted to the + charge. 3. Electromagnetic Theory: a moving charged body behaves like a magnet and can interact w/electrical and magnetic fields through which it passes. 4. Therefore, ray must consist of negatively charged particles, electrons. D. J.J. Thomson: British physicist who discovered the electron by using the cathode ray tube and electromagnetic theory. 1. Determined the ratio of electric charge to the mass of an individual electron. 2. Ratio = ±1.76 x 108 C/g, where C stands for coulomb (unit of electric charge). E. R.A. Millikan: American physicist who determined the charge of electron (Figure 4). 1. Oil Drop Experiment a. Sprayed oil droplets from an atomizer. b. Droplets settled into a beam of X-rays, becoming charged w/electrons. c. Applied an electric voltage to the top (+) and bottom (±) of the chamber so that the droplets would stop falling and remain stationary. 2. Calculating Charge of Electron a. Knowing the density of oil, he calculated the mass of each oil droplet (volume x density = mass (m)). b. Force of gravitational attraction: Force = mG, where G = acceleration due to gravity. c. Electrical force: F = Ee, where E = applied voltage and e = charge of electron. d. The two forces are equal when voltage is adjusted, so that the oil droplet remains stationary and mdropletG = Ee. e. Charge of individual electron = ±1.6 x 10±19 C. 3. Mass of Electron = = 9.10 x 10±28 g (extremely small mass). III. Radioactivity: spontaneous emission of particles and/or radiation (coined by Marie Curie). A. Anything that spontaneously emits radiation is said to be radioactive. B. Wilhelm Röntgen: German physicist responsible for discovering the X-ray. 1. Noticed that cathode rays caused glass and metals to emit very unusual rays. 2. Highly energetic radiation penetrated matter, darkened photographic plates, and caused a variety of substances to fluoresce. 3. Rays could not be deflected by magnet so did not contain charged particles. ~2~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own C. Antoine Becquerel: French physicist who discovered radioactivity in uranium. 1. Found that exposing thick wrapped photographic plates to certain uranium compound caused them to darken, even w/o stimulation of cathode rays. 2. Rays from uranium compound were highly energetic and could not be deflected by magnet but differed from magnets b/c they arose spontaneously. D. Three Types of Rays From Decay of Radioactive Substances (Figure 5) 1. Alpha ( ) Rays/Particles: consist of + particles; deflected by + plate. 2. Beta ( ) Rays/Particles: electrons; deflected by ± plate. 3. Gamma Rays: consist of high-energy rays; no charge; not affected by external field. Two features of atoms discovered by IV. The Proton and the Nucleus (Ernest Rutherford, Hans early 1900s: Geiger, Ernest Marsden) 1. Contain electrons. A. Gold Foil Experiment: used thin foils of 2. Are electrically neutral (to gold as targets for particles from a maintain electric neutrality, atom must contain same # of protons and electrons). radioactive source (Figure 6). -------------------------------------------------B. Results of the Experiment J.J. Thomson¶s Plum Pudding Model: 1. Majority of particles passed atom is a uniform, positive sphere of through foil w/little or no deflection. matter in which electrons are embedded. 2. In rare cases, particles would scatter at large angles and may The SI unit for atomic dimensions is in bounce back in the direction from terms of picometer (pm). which it had come. 1 pm = 1 x 10±12 m C. Conclusions 1. Most of atom must be empty space (explains why majority of particles passed through gold foil w/little or no deflection). 2. Atom¶s + charges are all concentrated in nucleus: dense central core within the atom and constitutes most of the atom¶s mass (explains why some particles were deflected at large angles). 3. Proton: positively charge particles in the nucleus (weighs 1.67262 x 10±24 g). 4. Electrons are spread out around the nucleus. V. The Neutron A. Ratio of mass of He atom to that of a H atom is 2:1 (known that H has 1 p+ and He has 2 p+); proven to be wrong because actual ratio of He atom to H atom should be 4:1. B. James Chadwick: British physicist who proved the existence of neutrons. 1. Bombarded a thin sheet of beryllium with particles, and very-high energy radiation similar to ray was emitted by metal. 2. Later experiments showed rays contained 3rd subatomic particle (neutron: electrically neutral particles having a mass slightly greater than that of protons, 1.67493 x ± 10 24 g). ~3~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own Figure 3 Figure 4 Figure 5 Figure 6 ~4~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own Atomic Number, Mass Number, and Isotopes I. Atomic Number (Z): number of protons in the nucleus of each atom of an element. A. In neutral atom, number of p+ and e± are equal; therefore, Z also represents # of e±. B. The atomic number can help to determine the chemical identity of an atom. II. Mass Number (A): total number of neutrons and protons present in the nucleus of an atom of an element. A. A = # of p+ + # of neutrons = Z + # of neutrons B. # of neutrons = A ± Z Ex. Give the number of protons, neutrons, and electrons in each of the following species: (a) (b) (a) # of protons = 8 # of electrons = 8 # of neutrons = 9 (b) # of protons = 80 # of electrons = 80 # of neutrons = 120 III. Isotopes: atoms that have the same atomic number but different mass numbers. A. Notation for Atomic Number and Mass Number of an Atom of Element (X): B. Hydrogen has three isotopes: , (deuterium), (tritium). C. Two common isotopes of uranium: (uranium-235) and (uranium-238). IV. Chemical properties of an element are determined primarily by the protons and electron in its atoms; neutrons do not take part in chemical changes under normal conditions. ~5~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own The Periodic Table I. Periodic Table: chart in which elements having similar chemical and physical properties are grouped together (organize large volumes of available information about the structure and properties of elemental substances and help to make predictions about chemical behavior) (Figure 7). A. Atoms are arranged by atomic number. B. Period: horizontal rows. C. Group: vertical columns. II. Three Categories of Elements A. Metal: good conductor of heat and electricity. B. Nonmetal: poor conductor of heat and electricity (17 elements). C. Metalloid: properties that are intermediate between those of metals and nonmetals (8 elements). III. Names of Certain Groups on Periodic Table A. Alkali Metals: group 1A elements including Li, Na, K, Rb, Cs, and Fr. B. Alkaline Earth Metals: group 2A elements including Be, Mg, Ca, Sr, Ba, and Ra. C. Halogens: group 7A elements including F, Cl, Br, I, and At. D. Noble Gases (Rare Gases): group 8A elements including He, Ne, Ar, Xe, and Rn. Figure 7 ~6~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own Molecules and Ions I. Molecules: aggregate of at least two atoms in a definite arrangement held together by chemical forces, called chemical bonds. A. Contain atoms of same element or atoms of 2 or more elements joined in fixed ratio. B. Like atoms, molecules are electrically neutral. C. Of all the elements, only the 6 noble gases exist in nature as single atoms (monatomic gases). D. Diatomic Molecules: molecules that contain only two atoms from a single element or different elements (i.e. H2, O2, N2, F2, Cl2, Br2, I2, HCl, CO). E. Polyatomic Molecules: molecules containing more than two atoms (i.e. O3, NH3, H2O). I. Ions: an atom or a group of atoms that has a net positive or negative charge (Figure 8). A. During chemical reactions, protons in the nucleus remain the same, but electrons may be lost or gained. 1. Cation: an ion with a net positive charge due to the loss of 1 or more electron from a neutral atom (i.e. Na lose electron Na+; Mg2+; Fe3+). 2. Anion: an ion with a net negative charge due to the gain of 1 or more electron from a neutral atom (i.e. Cl gain electron Cl±; S2±; N3±). 3. With very few exceptions, metals tend to form cations and nonmetals form anions. B. Ionic Compound: compound formed from cations and anions (i.e. NaCl). C. Monatomic Ions: contain only one atom (i.e. Na+ and Cl±). D. Polyatomic Ions: ions that contain two or more atoms with a net + or ± charge (i.e. NH4+, OH±, CN±). Figure 8 ~7~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own Chemical Formulas I. Chemical Formulas: express the composition (describe the element present and ratio in which the atoms are combined) of molecules and ionic compounds in terms of chemical symbols. A. Molecular Formula: shows the exact number of atoms of each element in the smallest unit of a substance (i.e. H2O = 2 hydrogens + 1 oxygen). 1. Subscript indicated the number of atoms of an element present. 2. If there is one atom of an element in the molecule, no subscript is necessary. 3. Allotrope: one of two or more distinct forms of an element (i.e. O2 and O3; carbon diamond and graphite). B. Structural Formula: shows how atoms are bonded to one another in a molecule (i.e. water = H ± O ± H). 1. The lines signify chemical bonds between two atoms. C. Molecular Models (Figure 9) 1. Ball-and-Stick Model a. Atoms are wooden or plastic balls with holes in them; balls are all the same size and each type of atom is represented by a specific color. b. Sticks or springs are used to represent chemical bonds. c. Angles formed btw atoms approximate bond angles in actual molecule. d. Advantages: shows the 3-D arrangement of atoms clearly; fairly easy to construct. e. Disadvantages: balls are not proportional to the size of atoms; sticks greatly exaggerate the space between atoms in a molecule. 2. Space-Filling Model a. Atoms are represented by truncated called held together by snap fasteners, so that the bonds are not visible. b. Balls are proportional in size to atoms. c. Advantages: More accurate b/c they show the variation in atomic size. d. Disadvantages: time-consuming to construct; do not show 3-D positions of atoms. D. Empirical Formula: tells us which elements are present and the simplest wholenumber ratio of their atoms (but not necessarily the actual number of atoms in a given molecule). 1. Hydrogen peroxide (H2O2) has an empirical formula of HO. 2. Hydrazine (N2H4) has an empirical formula of NH2. NOTE: empirical formulas are simplest chemical formulas, written by reducing subscripts in molecular formulas to smallest possible whole #; molecular formulas are true formulas of molecules. Ex. Write the molecular formula of methanol from its ball-and-stick model, shown to the left. CH3OH Ex. Write the molecular formula of chloroform from its balland stick model, shown to the left. CHCl3 E. Formula of Ionic Compounds (Figure 10) 1. Formulas of ionic compounds are usually the same as their empirical formulas b/c ionic compounds do not consist of discrete molecular units. ~8~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own 2. Ionic compounds are formed so that the arrangement of cations and anions is such that the compounds are electrically neutral. 3. Charges of the cation and anion are not shown. 4. To be electrically neutral, the sum of the charges on the cation and anion in each formula unit must be zero. 5. If the charges on the cation and anion are numerically different, the subscript of the cation is numerically equal to the charge on the anion, and the subscript of the anion is numerically equal to the charge on the cation. Ex. Determine the formula of potassium bromide. Potassium (K) is a cation with a 1+ charge; bromine (Br) is an anion with a 1 charge. The sum of the charges is +1 + (1) = 0, so no subscripts are necessary. Formula of potassium bromide is KBr. Ex. Determine the formula of zinc iodide. Zinc (Zn) is a cation with a 2+ charge; iodine (I) is an anion with a 1 charge. To make the sum of the charges equal to zero, you need two iodines to balance the charge of the zinc: +2 + 2(1) = 0. Formula of zinc iodide is ZnI2. Ex. Determine the formula of aluminum oxide. Aluminum (Al) is a cation with a 3+ charge; oxygen is an anion with a 2 charge. To make the sum of the charges equal to zero, you need to have two aluminums and three oxygens: 2(+3) + 3(2) = 0. Formula of aluminum oxide is Al2O3. Figure 9 Figure 10 ~9~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own Naming Compounds I. Chemical Nomenclature: the naming of chemical compounds. A. Organic Compounds: compounds that contain carbon, usually in combination with elements such as hydrogen, oxygen, nitrogen, and sulfur. B. Inorganic Compounds: compounds that do not contain carbon with a few exceptions: CO2, CO, CS2, CN±, CO32í, HCO3í. II. Ionic and Molecular Compounds A. Binary Compounds: compounds formed from just two elements. B. Ternary Compounds: compounds consisting of three elements. C. Transition Metals 1. Can form more than one type of cation (i.e. iron Fe2+ and Fe3+). 2. Assign the ending ±ous to the cation w/fewer positive charges and the ending ±ic to the cation w/more positive charges (i.e. Fe2+ ferrous ion; Fe3+ ferric ion). a. Does not provide info. regarding actual charges of 2 cations involved. b. ±ous and ±ic designations provide names for only two different elemental cations. 3. Stock System: designate different cations with Roman numerals, where I = 1+, II = 2+, and so on (i.e. MnO manganese (II) oxide; Mn2O3 manganese (III) oxide). Metal-Nonmetal General Form = MxNy M = any metal ion or NH4+. N = any nonmetal element and OHí and CNí. x, y = integers. 1. Use the full name of the metal or NH4+. 2. Write the stem of the nonmetal¶s name, hydroxide, or cyanide. 3. Add the suffix ±ide to the stem of the nonmetal. Ex. NaCl sodium chloride Al2O3 aluminum oxide Nonmetal-Nonmetal General Form = AxBy A and B represent different nonmetals. x, y = integers. 1. Write numerical prefix (1 = mono, 2 = di, 3 = tri, 4 = tetra, 5 = penta, 6 = hexa, 7 = hepta, 8 = octa, 9 = nona, 10 = deca) for x unless x = 1. 2. Add the full name of the nonmetal, A. 3. Write the numerical prefix for y unless y = 1. 4. Add the stem of the name of the second nonmetal, B. 5. Add the suffix ±ide to the stem. NOTE: if first element, A, is hydrogen, then no numerical prefixes are needed; prefix ³mono´ may be omitted for the 1st element; for oxides, the ending ³a´ in the prefix is sometimes omitted. Ex. NO2 nitrogen dioxide N2O4 dinitrogen tetroxide ~ 10 ~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own Cation-Polyatomic Anion Compounds General Form = (metal ion or NH4+)x(polyatomic anion)y 1. Use the full name of the metal ion/ammonium ion. 2. Name the polyatomic anion. Ex. Ex. Na2SO4 sodium sulfate (NH4)2CO3 ammonium carbonate Name the compounds: (a) Cu(NO3)2 (b) KH2PO4 (c) NH4ClO3 (d) SiCl4 (e) P4O10. (a) copper (II) nitrate (b) potassium dihydrogen phosphate (d) silicon tetrachloride (e) tetraphosphorous decoxide Ex. (c) ammonium chlorate Write chemical formulas for: (a) carbon disulfide (b) disilicon hexabromide (c) mercury (I) nitrite (d) cesium sulfide (e) calcium phosphate (a) CS2 (b) Si2Br6 (c) Hg2(NO3)2 (d) Cs2S (e) Ca3(PO4)2 Compound Cation: metal or NH4+ Anion: monatomic or polyatomic Binary compounds of nonmetals Ionic Cation has only one charge (Alkali earth metal/alkaline earth metal cations and Ag+, Al3+, Cd2+, Zn2+) Cation has more than one charge (Other metal cations) Use prefixes for both elements present (prefix "mono-" usually omitted for 1st element). Name metal first. Name metal first. If monatomic anion, add "-ide" to the root of the element name. Specify charge of metal cation with Roman numeral in parentheses. If polyatmoic anion, use name of anion. Molecular Add "-ide" to the root of the second element. If monatomic anion, add "ide" to the root of the element name. If polyatmoic anion, use name of anion. III. Acids and Bases A. Acid: a substance that yields hydrogen ions, H+, when dissolved in water. 1. Formulas for acids contain one or more hydrogen atoms as well as an anionic group íí> HxA (aq), where H = hydrogen, x = integer, and A = nonmetal. a. Write the prefix hydroí. b. Add the stem of the name of the nonmetal (exception: if A = sulfur, write sulfur instead of sulf). c. Add the suffix ±ic to the stem. d. Add the word acid. Ex. HF hydrofluoric acid H2S hydrosulfuric acid ~ 11 ~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own 2. Oxoacids: acids that contain hydrogen, oxygen and another element (the central element) (i.e. HNO3 nitric acid; H2CO3 carbonic acid; H2SO4 sulfuric acid). a. Formula usually written w/H 1st, then central element, and then O. b. 2 or > oxoacids often have same central atom but different # of O¶s. c. Starting with oxoacids whose names end with ±ic, we use the following rules to name these compounds. i. Addition of 1 O atom to ±ic acid: The acid is called ³per«-ic´ acid (i.e. HClO3 chloric acid HClO4 perchloric acid). ii. Removal of 1 O atom from ±ic acid: The acid is called ³íous´ acid (i.e. HNO3 nitric acid HNO2 nitrous acid). iii. Removal of 2 O atoms from ±ic acid: The acid is called ³hypo«ous´ acid (i.e. HBrO3 bromic acid HBrO hypobromous acid). 3. Oxoanions: anions of oxoacids. a. When all H ions are removed from the ±ic acid, the anion¶s name ends with ±ate (i.e. H2CO3 carbonic acid CO32í carbonate). b. When all H ions are removed from the ±ous acid, the anion¶s name ends with ±ite (i.e. HClO2 chlorous acid ClO2í chlorite). c. The names of anions in which one or more but not all the H ions have been removed must indicate the # of H ions present (i.e. H3PO4 phosphoric acid; H2PO4í dihydrogen phosphate; HPO42í hydrogen phosphate; PO43í phosphate). Common Parent Acids H2CO3 H3BO3 H4SiO4 HNO3 H2SO4 H2S2O3 H2CrO4 Ex. carbonic acid boric acid silicic acid nitric acid sulfuric acid thiosulfuric acid chromic acid H2Cr2O7 H3PO4 H3AsO4 HClO3 HBrO3 HIO3 HMnO3 dichromic acid phosphoric acid arsenic acid chloric acid bromic acid iodic acid manganic acid HOCN cyanic acid HSCN thiocyanic acid H2C2O4 oxalic acid H2C8H4O4 phthalic acid HC2H3O2 acetic acid H(NH2)SO3 sulfamic acid Name the following oxoacids and oxoanion: (a) H3PO3 and (b) IO4í. (a) phosphorous acid (b) periodate Oxoacid Removal of all H+ ions Oxoanion per-«-ate per-«-ic acid + [O] Representative ³-ic´ acid -ate í [O] ³-ous´ acid -ite í [O] hypo-«-ous acid hypo-«-ite ~ 12 ~ AP Chemistry Chapter 2: Atoms, Molecules, and Ions Do not submit this in as your own B. Base: a substance that yields hydroxide ions (OHí) when dissolved in water. KOH potassium hydroxide Ex. NaOH sodium hydroxide Ba(OH)2 barium hydroxide NOTE: ammonia (NH3) is also considered a common base b/c it yields OHí ions when dissolved in water: NH3 + H2O NH4+ + OHí. IV. Hydrate: compound that has a specific number of water molecules attached to it íí> salt ‡ x H2O, where x = a small whole number. A. Name the salt. B. Give the numerical prefix for the value of x. C. Add the word hydrate. Ex. CuSO4 ‡ 5 H2O copper (II) sulfate pentahydrate ZnF2 ‡ 4 H2O zinc fluoride tetrahydrate NOTE: water molecules can be driven off by heating; anhydrous: compound that has no H2O associated w/it. V. Peroxides: compounds that contain additional oxygen atoms; generally have one more O atom than the regular oxide. A. Writing the Formula of a Peroxide Given the Name 1. Write the formula of the normal oxide. 2. Add one more O atom to the formula. 3. Do not alter or reduce subscripts. lithium peroxide Li2O2 Ex. hydrogen peroxide H2O2 B. Naming Peroxides Given the Formula 1. Write the name of the 1st element. 2. Add the word peroxide. Ex. K2O2 potassium peroxide MgO2 magnesium peroxide ~ 13 ~ ...
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This note was uploaded on 03/12/2012 for the course CHEM 215 taught by Professor Tinich during the Spring '12 term at Neosho CC.

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