Law of Multiple Proportions Elements are composed of small particles,

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1 CHAPTER 2 Atoms, Molecules, and Ions 2.1 Atomic Theory Democritus (5th century BC) suggested that matter was made of small, indivisible particles, atomos John Dalton (1808) formulated modern atomic theory Dalton s hypotheses Law of Multiple Proportions Elements are composed of small particles, Law of Conservation of Mass atoms. All atoms of an element are identical. The atoms of one element are different from all other elements. Compounds are composed of atoms of more than one element in simple integral ratios. A chemical reaction involves the separation, combination, or rearrangement of atoms, not their creation or destruction. 16 X + 8 Y d 8 X 2 Y 2.2 Structure of the Atom Atoms consist of three subatomic particles: Electron Proton Neutron A. Magnetic Field B. No Field C. Electric Field Cathode rays = electrons J. J. Thomson, charge/ mass of e = 1.76 x 10 8 C/g (1906 Nobel Prize in physics) 1

2 Electrons Thomson s experiment showed that cathode rays are made of charged particles that interact with electric and magnetic field when moving Particles are negatively charged (repelled by the negative plate, attracted toward the positive plate) R. A. Millkan measured charge and mass of e (1923 Nobel Prize in physics) Negatively charged oil droplets suspended by gravity (down) vs. charge repulsion (up); smallest charge = single electron e charge = 1.60 x C Thomson s charge/mass of e = 1.76 x 10 8 C/g e mass = 9.10 x g RADIOACTIVITY: Spontaneous emission of particles and/or radiation Alpha (a) rays = positively charged α particles (helium nuclei) Gamma (g) rays = highenergy radiation (photons), with no charge Beta (b) rays = free electrons (negatively charged particles) Atomic structure model Atoms are neutral and electrons are negative. Where is the positive charge? Thomson s plum pudding model: atom is a uniform positive sphere with electrons embedded in it Rutherford s αscattering, 1910 Bombard gold foil with positive αparticles Most α penetrated foil undeflectedor slightly deflected Small number of α deflected at large angles back to source Atom is mostly empty space w/ positive charge concentrated in a central core (NOT Thomson s model) Positive core deflects α strongly Proton Positive central core of atom = nucleus Positive particles in nucleus = protons Proton charge is equal but opposite electron charge Proton mass = x g = 1840 x electron mass Nucleus occupies volume of atom Basketball in center of New Circle Road 2

3 Neutron 1 Problem: Hydrogen has 1 proton Helium has 2 protons The relative mass H:He is 1:4, not 1:2! What accounts for extra mass? Neutron 2 Something in an atom besides protons and electrons contributes to mass of an atom Chadwick (1932) detected neutral particles from αbombardment of Be Neutral particles = neutrons Charge = 0 Mass slightly greater than proton s Neutron 3 Hydrogen atom has 1 proton, no neutron Helium atom has 2 protons, 2 neutrons Relative mass = 1:4 Neutrons are also located in the nucleus More in Chap Isotopes Atomic number (Z) = number of protons Z identifies the element (see periodic table) Mass number (A) = total number of neutrons and protons Number of neutrons (N) = A Z Using A and Z What is the atomic number of chlorine? 17 How many protons does a Cl atom have? 17 How many neutrons does a Cl atom with mass number 37 have? 20 How many electrons does a Cl atom have? 17 Isotopes Atoms with the same atomic number but different mass number Most elements have two or more isotopes. Symbols distinguish different isotopes: Mass number = # protons + # neutrons Atomic number = # protons A Z X Element symbol determined by Z 3

4 11 B 5 Hydrogen isotopes How many protons does this atom have? 5 How many neutrons does this atom have? 6 Is it necessary to include the 5? No Each hydrogen isotope has its own name 1 H is hydrogen (protium) = H 2 H is deuterium = D 3 H is tritium = T Large relative mass difference between H isotopes leads to measurably different properties compare 235 U and 237 U Alkali Metals 2.4 Periodic Table Halogens Main group elements Transition metals Note categories Lanthanides and actinides Alkaline earth metals Noble Gases Having flunked three consecutive chemistry tests, Brad got home one day to discover that his parents had wallpapered his room with the periodic chart. Section 2.5 Molecules and Ions Monatomic single atom Noble gases are the only substances that exist as single atom gases under standard conditions Molecule two or more atoms in a definite arrangement held together by chemical bonds. Elements with all the same atom Cl 2 S 8 Compounds with different atoms H 2 O CH 4 Molecules Diatomic molecules made up of two atoms Typical elements: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 Typical compounds: HCl, NO, CO Polyatomic molecules made up of more than two atoms Typical elements: O 3, P 4 Typical compounds: BF 3, CO 2, C 6 H 14 4

5 Ions Charged species formed from neutral atom or molecule by gain or loss of electron(s) Cation = positive ion Electron(s) lost from neutral atom or molecule Anion = negative ion Electron(s) gained from neutral atom or molecule Monatomic ion ion of a single atom Examples: Cl, Fe 3+ Polyatomic ion ion containing more than one atom Examples: OH, CO 3, NH 4 + Ionic Compound a neutral compound formed fromcations and anions Examples: KI, MgCO 3, NH 4 Cl Section 2.6 Chemical Formulas Express composition of molecules and ionic compounds in chemical symbols H 2 O, NaCl Molecular Formula shows number of each element in the smallest unit of a substance H 2 O, C 6 H 6 (Why not just CH?) Allotropes distinct forms of an element Carbon: graphite (planar), diamond (3 dimensional), fullerenes (molecular) Oxygen: dioxygen (O 2 ) and ozone (O 3 ) Empirical vs. Molecular Formula Empirical Formula simplest wholenumber ratio of atoms in a compound Empirical formula CH Acetylene (C 2 H 2 ), benzene (C 6 H 6 ), cubane (C 8 H 8 ) Different properties more in Chap. 3 Formulas of ionic compounds are usually same as empirical formulas Sum of charges of anions and cations is zero NaCl is array of Na + and Cl ions (next slide) Exception: Na 2 S 2 is array of Na + and S 2 2 ions 5

6 Crystal of typical ionic solid Section 2.7 Nomenclature Organic compounds Contain carbon, usually with H, O, N, etc. Structures, therefore names, are very complicated and not covered here (Chap. 24) Inorganic compounds Ionic (include hydrates) Binary (2 elements) Ternary (3 elements) Molecular (include acids and bases) Common monatomic ions Hg 2 2+ is a diatomic ion Ionic Compounds Cations (positive ions), usually metals Group 13 metals, charge = group number Na + = sodium (cat)ion, Al 3+ = aluminum (cat)ion Metals with multiple charge states, e. g. Fe 2+ = ferrous ion (low) Fe 3+ = ferric ion (high) Anions (negative ions), usually nonmetals O = oxide, F = fluoride Charge typically group number 18 Ionic Compounds Name as cation (charge), anion, (# hydrate) Examples: CaBr 2 calcium bromide (no need to indicate # Br) SnCl 2 tin(ii) chloride [Stock] or stannous chloride MgSO 4 7H 2 O magnesium sulfate heptahydrate Polyatomic ions Table 2.3. Memorize names and charges of polyatomic ions in table Significant examples follow 6

7 Important polyatomic ions Name Formula Ammonium + NH 4 Carbonate CO 3 Hydroge n carbonate HCO 3 (bicarbonate) Chlorate ClO 3 Perchlorate ClO 4 Chromate CrO 4 Dichromate Cr 2 O 7 Permanganate MnO 4 Important polyatomic ions, contd. Cyanid e CN Hydroxide OH Peroxi de O 2 Sulf ate SO 4 Thiocyanate SCN Nitri te NO 2 Nitr ate NO 3 Phosphate 3 PO 4 Number of oxygens d Name Oxygen anions O = oxide O 2 = peroxide O = superoxide Oxoanions of chlorine ClO 4 = perchlorate ClO 3 = chlorate ClO = chlorite ClO = hypochlorite Add hydrogen ion (H + ) to anions Raise charge by +1 Add hydrogen to name H 2 PO 4 = dihydrogen phosphate HPO 4 = hydrogen phosphate PO 4 3 = phosphate Name from formula & vice versa Nickel(III) oxide Ni 2 O 3 Cuprous cyanide or copper(i) cyanide CuCN Zn(NO 3 ) 2 Zinc nitrate Al(OH) 3 Aluminum hydroxide Ammonium fluoride NH 4 F Naming Molecular Compounds Discrete molecular units Often composed of nonmetallic elements Name similarly to ionic compounds Simple binary compounds HF = Hydrogen fluoride BN = Boron nitride 7

8 Greek prefixes give number of atoms Naming Molecular Compounds Mono 1 Di 2 Tri 3 Tetra 4 Penta 5 Hexa 6 Hepta 7 Octa 8 Nona 9 Deca 10 PCl 5 Phosphorus pentachloride mono is usually omitted for first element Phosphorus(V) chloride is acceptable N 2 O 4 Dinitrogen tetroxide a ending of prefix is (tetra) dropped for oxide Carbon dioxide CO 2 Digermanium hexachloride (orhexachloro digermane) Ge 2 Cl 6 SeF 4 Selenium tetrafluoride Cl 2 O 7 Dichlorine heptoxide Common (trivial) names of hydrides B 2 H 6 diborane CH 4 methane SiH 4 silane NH 3 ammonia PH 3 phosphine H 2 O water Acids and Bases (pp 58 60) Postpone until Chapter 4 Acids (Chap. 4) Acid = Proton (H + ) donor in H 2 O ide anion forms ic acid Br = bromide; HBr = hydrobromic acid (in water) or hydrogen bromide (gas) 8

9 Oxoacids and oxoanions Oxoacids contain hydrogen, oxygen and a third (central) element NO 3 = nitrate; HNO 3 = nitric acid SO 4 = sulfate; H 2 SO 4 = sulfuric acid Variable number of oxygens ClO 4 = perchlorate; HClO 4 = perchloric acid ClO 3 = chlorate; HClO 3 = chloric acid ClO = chlorite; HClO 2 = chlorous acid ClO = hypochlorite; HClO = hypochlorous acid H partly removed H 2 PO 4 = dihydrogen phosphate; HPO 4 = hydrogen phosphate; PO 4 3 = phosphate Bases (Chap. 4) Base = Hydroxide (OH ) donor in water Named as simple hydroxide salts KOH = potassium hydroxide Mg(OH) 2 = magnesium hydroxide Some substances yield hydroxide in water NH 3 = ammonia Reacts with H 2 O to give NH 4 + (ammonium) and OH (hydroxide) Naming Hydrates Hydrates compounds with a specific number of water molecules included Use Greek prefix to tell how many waters, followed by hydrate Example: CuSO 4. 5H 2 O Copper (II) sulfate pentahydrate 9