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Atomic Theory & Chemical Nomenclature Chapter 2 HW:16,17, 24, 30,34,35,43,57,59 10th edition: 61,66,68,71 11th edition: 65,70,72,77 Early Evidence for Atomic Theory • Democritus – matter is made of separate, discrete particles (~ 450 BC) • Total mass of a reaction stays the same – Law of Conservation of Mass (1789 by Antoine Lavoisier) • When elements combined to make compounds, they always occurred in specific ratios – Law of definite proportions (1799 by Joseph Proust) – Water always decomposes in a ratio of 8 g O2 to 1 g H2. • Pollen particles found to exhibit apparent random motion (1827 by Robert Brown) – Einstein explained the phenomenon in 1905 as being the result of microscopic collisions of atoms (water molecules). John Dalton’s Atomic Theory (A New System of Chemical Philosophy: 1808) 1. Elements are composed of small particles called atoms. Atomos - uncuttable 2. All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements (size, mass, properties). 3 Dalton’s Atomic Theory (1808) 3. Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms is fixed. H2O, NH3, MgCl2, CO2 4. A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. 2Fe + 4H2O → 2H2 + 2Fe(OH)2 4 Dalton’s Atomic Theory # 3 Supported Proust’s Law of Definite Proportions Dalton’s Atomic Theory # 4 16 X + 8Y 8 X2Y Supported Lavoisier’s Law of Conservation of Mass Cathode Ray Tube: Thomson; e- Charge/Mass Magnetic field (N/S) result Neither or Both fields + - Electric field (+/-) result Cathode Ray Tube Video J.J. Thomson (1897), measured charge/mass of e7 e- = -1.76 x 108 C/g (1906 Nobel Prize in Physics) Cathode Ray in a discharge tube The ray is invisible, but the fluorescence of a zinc sulfide coating on the glass causes it to appear green. Cathode rays are attracted to positive charges. They consist of negatively charged electrons 8 Millikan’s Oil Drop Experiment Measured mass of e(1923 Nobel Prize in Physics) e- charge = -1.60 x 10-19 C Thomson’s charge/mass of e- = -1.76 x 108 C/g Millikan’s Experiment e- mass = 9.10 x 10-28 g 9 Thomson’s Model (1904): Plum-pudding model - Knew atoms were neutral - But they had electrons (e-) - Theorized an evenly distributed (+) charge “cloud” 10 Radioactivity • Recent discovery of X-rays (1895) • Many substances were discovered to also give off high energy rays in late 1890’s • Spontaneous emission of particles and/or radiation by elements via decay breakdown Alpha particle: +2 • Term coined by Marie Curie Gamma radiation: 0 Beta particle: -1 (uranium compound) Ernest Rutherford’s Gold Foil Experiment (with Hans Geiger and student Ernest Marsden) (Rutherford had already won 1908 Nobel Prize in Chemistry for work in radioactivity) particle velocity ~ 1.4 x 107 m/s (~5% speed of light) 1 out of every 20K deflected at > 90 1. atoms positive charge is concentrated in the nucleus 2. proton (p) has opposite (+) charge of electron (-) 3. mass of p is 1840 x mass of e- (1.67 x 10-24 g) 12 atomic radius ~ 100 pm = 1 x 10-10 m Rutherford’s Model of the Atom nucleus radius ~ 5 x 10-3 pm = 5 x 10-15 m 20,000X smaller “If the atom is the Houston Astrodome, then the nucleus is a marble on the 50yard line.” 13 James Chadwick’s Experiment (Neutron) (1932) (1935 Nobel Prize in Physics) Rutherford’s proposal accounted for only the mass of protons H atoms: 1 p+; He atoms: 2 p+ mass He/mass H should then be = 2/1 Actual measured ratio = 4/1 Chadwick subjected 9Be sheet to alpha radiation + 9Be 1n + 12C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10-24 g mass p ≈ mass n ≈ 1840 x mass e- 15 Atomic Number, Mass Number, and Isotopes Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Mass Number A ZX Atomic Number 1 1H 235 92 2 1H U (D) 238 92 *Position may change Element Symbol 3 1H U (T) Atomic # is always a whole number Ions and Isotopes • Ion – Charged particle (atoms), various e– Element with unequal numbers of electrons and protons – Atoms can gain/lose e- to form ions (never P+’s) Mg → Mg+2 + 2eN + 3e- → N-3 • Isotopes – chemically similar atoms, various N0 – Same # of P+, unequal numbers of neutrons – Atoms can vary neutrons to form isotopes (not P+’s) The Isotopes of Hydrogen 99.985% 0.015% Trace; radioactive Natural Abundance 18 Practice: How many P+, N0, e-? Ne P+ 10 N0 20-10 = 10 e10 Ca 20 20 40-20 = 20 20 * 30 65-30 = 35 30 – 2e- = 28 20 10 40 65 14 +2 Zn -3 N *Carbon-14 7 6 14 - 7 = 7 14 - 6 = 8 *the atom is neutral, # e- = # P+ 7 + 3e- = 10 6 the ion is positive = it has lost ethe ion is negative = it has gained e- *If we know the element we can find atomic # on Periodic Table *if only one number is listed, it is always referring to the atomic mass Practice: How many P+, N0, e-? P+ N0 e- P 15 15 14 15 59 Ni+3 28 31 25 I- 53 70 54 8 9 8 34 46 36 29 123 Oxygen-17 Se-2 80 Group Period Noble Gas Halogen Transition Metals Alkali Metal Alkaline Earth Metal The Modern Periodic Table 21 Chemistry In Action Natural abundance of elements in Earth’s crust: Natural abundance of elements in human body: 22 A molecule is an aggregate of 2 or more atoms in a definite arrangement held together by chemical forces. Composed of only non-metals H2 H2O NH3 CH4 NaCl ? *Note: If it contains a metal it is not a molecule Ionic compound A molecule is an aggregate of 2 or more atoms in a definite arrangement held together by chemical forces. The 7 diatomic elements shown in blue are always found bonded to each other (or another element). *Know these diatomic elements A polyatomic molecule contains more than 2 atoms: H2O, NH3, CH4, P4, P8 Allotropes – multiple forms of an element Representation: Formulas and Models 25 Example 2.2 Write the molecular formula of methanol, an organic solvent and antifreeze, from its ball-and-stick model, shown below. There are four H atoms, one C atom, and one O atom. Therefore, the molecular formula is CH4O. However, the standard way of writing the molecular formula for methanol is CH3OH because it shows how the atoms are joined in the molecule. A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance. An empirical formula shows the simplest wholenumber ratio of the atoms in a substance. molecular empirical N2H4 NH2 C6H14N2O2 O3 H2O C3H7NO O H2O Empirical data: knowledge based on observation or experience rather than theory or pure logic. These ratios were first experimentally observed from decomposition of substances 27 Example 2.3 Empirical formulas Write the empirical formulas for the following molecules: (a) acetylene (C2H2), which is used in welding torches There are two carbon atoms and two hydrogen atoms in acetylene. Dividing the subscripts by 2, we obtain the empirical formula CH. (b) glucose (C6H12O6), a substance known as blood sugar In glucose there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms. Dividing the subscripts by 6, we obtain the empirical formula CH2O. (c) nitrous oxide (N2O), used as an anesthetic gas (“laughing gas”) and as an aerosol propellant for whipped creams. Because the subscripts in N2O are already the smallest possible whole numbers, the empirical formula for nitrous oxide is the same as its molecular formula. A monatomic ion contains only one atom: Na+, Cl-, Ca2+, O2-, N3-, Fe2+, Fe3+ Some elements can have multiple charge states (oxidation states), especially transition metals. Their charges can also be represented as Roman Numerals (Iron II vs Iron III) A polyatomic ion contains more than one atom: They are molecules that have a net charge Hydroxide: OH-1 Carbonate: CO3-2 Sulfate: SO4-2 Phosphate: PO4-3 Nitrate: NO3-1 Chlorate: ClO3-1 Ammonium: NH4+ Memorize these common seven Nitrate (NO3-1) +1 Common Ions Shown on the Periodic Table “Memorize these” +2 Transition metals can have Various charge states +3 0 -3 -2 -1 Charges only if found in an ionic compound, But NOT if neutral or found in a molecule Group 4A doesn’t often form ionic compounds 30 An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive (+) charge If a neutral atom loses one or more electrons it becomes a cation (oxidation). Na 11 protons 11 electrons Na+ 11 protons 10 electrons anion – ion with a negative (-) charge If a neutral atom gains one or more electrons it becomes an anion (reduction). Cl 17 protons 17 electrons Cl- 17 protons 18 electrons Ionic compounds form from alternating cations(+) and anions(-) (Metal+ and nonmetal-) The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero. MgO, CaCl2, CoF2, K2S The formula is the same Mg2O2 as the empirical formula. Is NOT A MOLECULE, no fixed quantity of atoms 32 Ionic Compounds form a Crystal Lattice • In NaCl, there are 6 Cl- interacting with 1 Na+. – Each Cl- is interacting with 6 Na+. • Therefore, they are in a 1:1 ratio. • Na1Cl1 gives the element ratio, but is not a molecule NaCl is the formula unit Cl- Cl- ClNa+ Cl- Cl- Cl- Crystal Lattice Molecule or Ionic Compound? • C6H12O6 • Molecule *Remember • NaCl • Ionic Compound Molecules are made of only non-metals • CO2 • Molecule Ionic compounds have a metal and non-metal • O2 • Molecule • LiOH • Ionic Compound • Fe • Neither, just an atom • Fe(NO3)2 • Ionic Compound “Compound” • Term only refers to a chemical that consists of multiple types of elements. • Ionic compounds are always compounds: (Metal/Non-metal) NaCl, FeBr2, Al(NO3)3 • Molecules are normally compounds, but not always. • Compound: H2O, SF6, CO2, C6H12O6 • Non-compound: O2, S8, Br2, P4 Polyatomic ions • Covalently bonded molecules that possess a net charge • They can then participate in ionic bonds, but act as single units – – – – – – – Ammonium: NH4+ Carbonate: CO3-2 Phosphate: PO4-3 Sulfate: SO4-2 Hydroxide: OHNitrate: NO3Chlorate: ClO3-1 Most common elements in Ionic Compounds The most reactive metals (green) and the most reactive nonmetals (blue) combine to form ionic compounds. 2Na (s) + Cl2 (g) → 2NaCl (s) + Energy! Example 2.4 Write the formula of Calcium carbonate (limestone), containing the Ca2+ and CO32− ions. Solution To satisfy electrical neutrality, the following relationship must hold: (+2)x + (−2)y = 0 The simplest ratio is 1:1 giving the formula unit: CaCO3 Check The subscripts are reduced to the smallest whole- number ratio of the atoms because the chemical formula of an ionic compound is usually its empirical formula. Formulas of Ionic Compounds: Must equal 0 2 x +3 = +6 Al3+ 3 x -2 = -6 Al2O3 1 x +2 = +2 Mg+2 & N-3 O21 x -2 = -2 CaS Ca2+ 2 x +1 = +2 Na+ S21 x -2 = -2 Na2CO3 “Crisscross and reduce” CO3239 Ionic bond formation practice 1st determine ionized form • • • • • • • Mg and Br Al and OHRb and N Sr and PO4-3 Co (II) and O Ca and CO3-2 NH4+ and O 2nd Cross charges and reduce Mg+2 Br-1 Al+3 OHRb+ N-3 Sr+2 PO4-3 MgBr2 *Al(OH)3 Rb3N *Sr3(PO4)2 Co+2 O-2 CoO Ca+2 CO3-2 CaCO3 NH4+ O-2 *(NH4)2O *Polyatomic ions are in parentheses if needing multiple charges Chemical Nomenclature Naming Ionic Compounds • Ionic Compounds – Often a metal + nonmetal (or metal + polyatomic ion) – Name metal first (left unchanged) – Then Anion (nonmetal), but add “-ide” to element name BaCl2 barium chloride K2O potassium oxide Mg(OH)2 magnesium hydroxide KNO3 potassium nitrate If anion is a polyatomic ion, simply name it 42 • Transition metals in ionic compounds – Must indicate charge on metal using Roman numerals FeCl2 2 Cl- -2 so Fe is +2 Iron (II) chloride “Ferrous chloride” FeCl3 3 Cl- -3 so Fe is +3 Iron (III) chloride “Ferric chloride” Cr2S3 3 S-2 -6 so Cr is +3 (6/2) Chromium (III) sulfide Tungsten Carbide (WC) drill bits *Miner Carbide head lamps (CaC2) 3x stiffer than steel Reacts with H2O produce Acetylene + charge - charge 45 Example 2.5 Formulas to Names Name the following compounds: (a) Cu(NO3)2 Copper (II) nitrate (b) K3PO4 Potassium phosphate. (c) (NH4)2CO3 Ammonium carbonate (d) Ca3P2 Calcium Phosphide (e) Cr(OH)3 Chromium (III) Hydroxide (f) NiS Nickel (II) Sulfide Example 2.5 Solution (a)The nitrate ion (NO3-1) bears one negative charge, so the copper ion must have two positive charges. Copper (II) nitrate. (b)The cation is K+ and the anion is PO4-3 (phosphate). Because potassium only forms one type of ion (K+), there is no need to use potassium (I) in the name. The compound is potassium phosphate. (c) The cation is NH4+1 (ammonium ion) and the anion is CO3-2 . The compound is ammonium carbonate. (d) Calcium Phosphide, Ca is not a transition metal so we do not need to specify the charge with Roman numerals. (e) Chromium (III) Hydroxide Example 2.6 Name to Formulas Write chemical formulas for the following compounds: (a) chromium (III) nitrate Cr(NO3)3 (b) cesium sulfide Cs2S (c) calcium phosphate Ca3(PO4)2 (d) lead (II) nitrate (e) Calcium Carbonate Pb(NO3)2 CaCO3 Example 2.6 Solutions (a) We can cross the +3 charge of Chromium (III) with the -1 charge of Nitrate to yield Cr(NO3)3 using parentheses to signify we need 3 x NO3-1 to 1 x Cr+3 (b) Each sulfide ion bears two negative charges, and each cesium ion bears one positive charge. Therefore, the formula is Cs2S. (c) Each calcium ion (Ca2+) bears two positive charges, and each phosphate ion (PO4-3) bears three negative charges. 3(+2) + 2(−3) = 0 The formula is Ca3(PO4)2 (d) Pb(NO3)2 (e) Crossing the charges would give Ca2(CO3)2, but this reduces to CaCO3 • Molecular compounds − Nonmetals or nonmetals + metalloids − Some have common names (small) − H2O, NH3, CH4 , PH3, O3 − Element furthest to the left in a period and closest to the bottom of a group on periodic table is placed first in formula − If more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom − Last element name ends in -ide 50 Molecules can arrange the same elements in various arrangements. Necessitates the usage of prefixes to distinguish NO: Nitrogen Monoxide Dinitrogen trioxide: N2O3 Dinitrogen pentaoxide N2O5 Nitrogen dioxide NO2 Dinitrogen monoxide N2O Molecular Compounds P4O10 tetraphosphorus decoxide NF3 nitrogen trifluoride SF6 sulfur hexafluoride N2Cl4 dinitrogen tetrachloride NO2 nitrogen dioxide *The prefix mono- is not normally used if on first element 52 53 Example 2.8 Molecule Name to Formula Write chemical formulas for the following molecular compounds: (a) carbon disulfide (a) Because there are two sulfur atoms and one carbon atom present, the formula is CS2. (b) There are two silicon atoms and (b) disilicon hexabromide six bromine atoms present, so the formula is Si2Br6. (c) B2O3 (c) Diboron Trioxide 55 Hydrates are ionic compounds that have a specific number of water molecules attached to them. BaCl2•2H2O barium chloride dihydrate LiCl•H2O lithium chloride monohydrate MgSO4•7H2O magnesium sulfate heptahydrate Sr(NO3)2•4H2O strontium nitrate tetrahydrate CuSO4•5H2O Copper (II) sulfate pentahydrate Copper (II) sulfate “anhydrous” CuSO4 56 An acid can be defined as a substance that yields hydrogen ions (H+) when dissolved in water. For example: HCl gas and HCl in water • Pure substance, hydrogen chloride • Dissolved in water (H3O+ and Cl−), hydrochloric acid 57 Simple acids: Hydrogen bonded to single element (often a Halogen) Specifically does not have oxygen present 58 An oxoacid is an acid that contains hydrogen, oxygen, and another element. HNO3 nitric acid H2CO3 carbonic acid H3PO4 phosphoric acid 59 Naming Oxoacids and Oxoanions 60 Prefix and Suffix change as number of oxygen atoms change (oxidation state) 61 The rules for naming oxoanions, anions of oxoacids 1. When all the H ions are removed from the “-ic” acid, the anion’s name ends with “-ate.” Nitric Acid (HNO3) vs Nitrate (NO3-1) 2. When all the H ions are removed from the “-ous” acid, the anion’s name ends with “-ite.” Nitrous Acid (HNO2) vs Nitrite (NO2-1) 3. The names of anions must indicate the number of H ions present. –H PO - dihydrogen phosphate 2 4 –HPO4 2- hydrogen phosphate –PO43- phosphate 62 Example 2.9 Name the following oxoacid and oxoanion: (a)HNO3 and NO3-1 (b)H2CO3 and CO3-2 (c)H3PO3 and PO3-3 (d)HClO4 and ClO4-1 Nitric Acid and Nitrate Carbonic Acid and Carbonate Our reference acid is phosphoric acid (H3PO4). Because H3PO3 has one less O atom, it is called phosphorous acid. PO3-3 has one less O atom than phosphate so it is called phosphite. Our reference acid is Chloric Acid (HClO3). Because HClO4 has one more O atom, it is called perchloric acid. ClO4-1 is called perchlorate. A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water. As an ionic compound, it follows the same nomenclature, the anion is always hydroxide NaOH KOH Ba(OH)2 sodium hydroxide potassium hydroxide barium hydroxide 64 Organic Chemistry Carbon Molecules If the molecule has Carbon and Hydrogen it follows organic nomenclature instead Single bonds: end with -ane Double bonds: end with -ene Triple bonds: end with -yne Ethane Ethene Ethyne “Acetylene” Organic chemicals are often named by their Functional Groups that have similar atomic arrangements for similar properties. H H Alcohol C OH H methanol H H Amine C NH2 H methylamine Carboxylic Acid H O H C C OH H acetic acid 66 Organic Structure Representation • Octonol: C8H18O or C8H17OH • CH3CH2CH2CH2CH2CH2CH2CH2OH • CH3(CH2)7OH • Condensed • Space-filled