Chapter 2: Atoms and Elements PDF

1 Chapter 2 : Atoms and Elements Chemical Laws 1. Law of Conservation of Mass In a chemical reaction, matter is neither created nor destroyed 2. Law of Definite Proportions All samples of a given compound have the same proportions of their constituent elements 3. Law of Multi...

1 Chapter 2 : Atoms and Elements Chemical Laws 1. Law of Conservation of Mass In a chemical reaction, matter is neither created nor destroyed 2. Law of Definite Proportions All samples of a given compound have the same proportions of their constituent elements 3. Law of Multiple Proportions (Dalton’s Law) When two elements (A & B) form 2 different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers 2 Discovery of the Electron (Thomson) cathode rays are streams of electrons electrons are particles found in all atoms the electron has a charge of -1.60 x 10-19 C & has a mass of 9.1 x 10-28 g (Millikan) electron Atoms are neutral so something must be balancing out the charge Plum pudding model of the atom (1904) : Soup of positive charge 3 Radioactivity and the Nuclear Atom Henri Becquerel, French physicist (1896) Some materials produce invisible radiation, consisting of charged particles. Radioactivity (Marie Curie, 1897) Discovery of Radium Spontaneous emission of high energy radiation and particles Beta particles (, high energy electrons) Alpha particles (, +2 charge, mass = helium nucleus) (Rutherford, McGill University) 4  2.2 Rutherford’s Experiments Testing models for make-up of atom – Can test using alpha particles (from radioactive source – courtesy of M. Curie!!) Hypothesis if atom was like a Plum Pudding plum pudding, all Atom Model the particles should go straight through 5 Rutherford’s Gold Foil Experiment Most α particles pass through with little or no deflection A few α particles are deflected through large angles 6  Rutherford’s Interpretation of the Atom A tiny dense center called the nucleus - space taken by the nucleus is only about 1/10,000 the size of the atom The nucleus has essentially the entire mass of the atom The nucleus is positively charged - the amount of positive charge balances the negative charge of the electrons The electrons are dispersed in the empty space of the atom surrounding the nucleus 7 Problems with mass ?!?!? Rutherford realized that the total mass of the atom was not equal to the sum of the masses of its protons and electrons. ie. Hydrogen (with 1 proton and 1 electron) = ¼ the mass of Helium 2 protons and 2 electrons In 1932 Chadwick : missing mass came from particles that were the same size as protons but were neutral. = Neutron 8 Elements unique name, symbol & number of protons in its nucleus the number of protons in the nucleus of an atom is called the atomic number, Z. 2.3 Isotopes all isotopes of an element : are chemically identical i.e. they undergo the exact same chemical reactions. have the same number of protons have different masses (due to different numbers of neutrons) isotopes are identified by their mass number, A. Different isotopes are called nuclides 9 2.3 Symbols of Elements Mass number (A) – total number of nucleons (protons & neutrons) in the nucleus (unique to each isotope) Charge (if any) – not shown if 0 A Y X Elemental symbol – a one- or two- Z letter symbol to identify the type of atom Atomic number (Z) – the number of protons in the nucleus; determines the identity of the element 10 Isotopes (cont. 1) A Y X Z Most elements have two or more isotopes, atoms that have the same atomic number (Z) but different mass numbers (A). 238 235 92 U 92 U 238−92 = 146 neutrons 235−92 = 143 neutrons 11 Atomic Numbers (a) * What are the atomic number, mass number, and symbol of the chlorine isotope with 18 neutrons (b) How many protons, electrons, and neutrons are present in an atom of 52 24 Cr 2+ ? 12  2.4 Atomic Mass We arbitrarily assign the 12C atom a mass of exactly 12 atomic mass units (12 amu) and then scale the mass of every other atom to that. Sometimes the term daltons (Da) is used instead of amu  Atomic Mass (cont) based on the mass of all isotopes of a given element we generally use the average mass of all an element’s isotopes found in a sample in calculations - average must take into account the abundance of each isotope in the sample we call the average mass the atomic mass Atomic Mass = ∑ (fractional abundance of isotope)n x (mass of isotope)n Atomic Mass Example Neon has three naturally occurring isotopes. What is the average atomic mass of neon ? (19.9924 amu x 0.904838) + (20.9940 amu x 0.002696) + (21.9914 amu x 0.092465) 20.1799 amu 15  Ex 2.2 : Calculating Atomic Mass (a) An element has two naturally occurring isotopes. Isotope 1 has a mass of 120.9038 amu and a relative abundance of 57.4%, and isotope 2 has a mass of 122.9042 amu. What is the element?  2.5 The Periodic Table of Elements Mendeleev (1834-1907) ordered elements by increasing mass Saw a repeating pattern of properties Periodic Law – When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically Broad Categories of Elements Metals (left side and bottom of the table) Shiny solids; conduct heat and electricity; are malleable and ductile, variable charge on ions Nonmetals (right side and top of the table) Solids, liquids, and gases; nonconductors; solids are brittle, constant charge on ions Metalloids (between metals/nonmetals) Shiny solids (like metals); brittle (like nonmetals); semiconductors, constant charge on ions 18  Nonmetals Metalloids Metals 19 1A 8A 2A 3A 4A 5A 6A 7A 3B 4B 5B 6B 7B 8B 1B 2B 20 2.6 Trends in Compound Formation Elements combine together in characteristic ratios These characteristic ratios were the basis of Mendeleev’s periodic patterns Dalton had first seen this : Dalton’s Law of Multiple Proportions - when elements A & B react together to form two distinct compounds, the amounts of A that react with a fixed amount of B can be expressed as a ratio of small whole numbers 21 Molecular compounds : NO2, NO, CO, CO2 are all examples of molecular compounds Composed of atoms held together by covalent bonds - sharing of one or more pairs of electrons Compounds formed between nonmetal elements Molecular formula : show the number and types of atoms present in one molecule of the compound Empirical formula : the smallest whole number ratio of the elements in a compound 22  Example : Ethylene Glycol (1,2 - ethanediol) Molecular Formula – C 2H6O2 Empirical Formula – CH3O Structural Formulas OH OH H2C CH2 CH2(OH) CH2(OH) Structura Space Filling l model model Condensed structural Ball and Stick model model Ionic compounds: Look primarily at binary (two element) ionic compounds Formed between cations (positive charge) of metals and anions (negative charge) of nonmetals Na lose e- Na+ + Cl- Cl gain e- - Charge is normally dependent on location on periodic table (next) 24  Ionic compounds (cont): Ionic compounds are held together by the electrostatic charges This means that one ion of Cl- actually interacts with 6 atoms of Na+ (cubic packing) Formula Unit Formula unit – smallest electrically neutral unit within crystal 25  Predicting Likely Charges of Main Group Ions Non-metals gain electrons to give same number as nearest noble gas. Main Group: tend to lose electrons to give same H+ number of electrons as nearest noble gas. Li+ Mg2+ N3- O2- F- Other metals : more than one ion possible. Na+ Ca2+ Al3+ P3- S2- Cl- V3+ Cr2+ Mn2+ Fe2+ Co2+ Ni2+ Cu+ K+ Sr2+ Sc3+ Ti4+ V5+ Cr3+ Mn4+ Fe3+ Co3+ Ni3+ Cu2+ Zn2+ Ga3+ Se2- Br- Sn2+ Rb+ Ba2+ Y3+ Zr4+ Ag+ Cd2+ In3+ Sn4+ Te2- I- Hg 2+ Tl+ Pb2+ Cs+ 2 Hg 2+ Tl3+ Pb4+ 26 2.7 Naming compounds and writing formula Compounds are given names so that we can unambiguously identify the compounds For simple molecules, the name will consist of a two- word name More complicated molecules get more complicated names, with specialized names to represent larger groups 27 Molecular compounds 1. First word is the name of the first element 2. Second word is the name of the second element, changing the end to -ide 3. Use prefixes to indicate the number of atom of each element in compound (no mono in name of first element in name) 28 Examples Name the following compounds SO3 P4O10 29 Ionic compounds Very similar naming for binary ionic compounds 1. First word is the name of the cation 2. Second word is the name of the anion (ending with ide) Don’t use prefixes as the number of atoms is predefined by the charge on the ions 30 Examples Name the following compounds CaBr2 K2S What is the molecular formula for : Zinc nitride 31 H+ Li+ Mg2+ N3- O2- F- Na+ Ca2+ Al3+ P3- S2- Cl- V3+ 2+ Mn2+ Fe2+ Co2+ Ni2+ Cu+ K+ Sr2+ Sc3+ Ti4+ V5+ Cr3+ Mn4+ 3+ Co3+ 3+ 2+ Zn2+ Ga3+ Se2- Br- Cr Fe Ni Cu Sn2+ Rb+ Ba2+ Y3+ Zr4+ Ag+ Cd2+ In3+ Sn4+ Te2- I- Hg 2+ Tl+ Pb2+ Cs+ 2 Hg 2+ Tl3+ Pb4+ 32 Examples Name the following compounds CaBr2 K2S What is the molecular formula for : Zinc nitride 33 What about metals with different charges ? Many of the transition metals can have different charges on their ions Iron : Fe2+, Fe3+ Old naming system – different endings depending on the “state of oxidation” eg : ferrous, ferric New system : Stock system : use element name plus a roman numeral for charge : Iron (II), Iron(III) 34 Polyatomic ions Many ionic compounds consist of bigger groups of atoms as the ions These groups are given their own special names and treated as a single ion. The ones to know : Carbonate CO32- Ammonium NH4+ Bicarbonate HCO3- Hydroxide OH- (HO-) Hypochlorite ClO- Chlorite ClO2- Nitrate NO3- Chlorate ClO3- Nitrite NO2- Perchlorate ClO4- Phosphate PO43- Sulfate SO42- Hydrogen phosphate HPO42- Hydrogen sulfate HSO4- Dihydrogen phosphate H2PO4- Sulfite SO32- Hydrogen sulfite HSO3- 35 Take a closer look : There are relationships in these names: For groups that contain O’s (oxyanions) – changing the number of O’s results in a name change: Named specially (examples in Table 2.3 and 2.4) per- > “-ate” > “-ite” > hypo- NOTE : oxygen isn’t in the names Hypochlorite ClO- Chlorite ClO2- Nitrate NO3- Chlorate ClO3- Nitrite NO2- Perchlorate ClO4- Sulfate SO42- Sulfite SO32- 36 Still looking Groups can have different numbers of hydrogens – add “hydrogen” to front (more than one, use same numbering as earlier NOTE : Can also use “bi” or “bis” instead of hydrogen Carbonate CO32- Bicarbonate HCO3- Phosphate PO43- Hydrogen phosphate HPO42- Dihydrogen phosphate H2PO4- Some groups can have both : Sulfate SO42- Hydrogen sulfate HSO4- Sulfite SO32- Hydrogen sulfite HSO3- 37 Helpful memory aid For the oxyanions : learn the “ates” Pneumonic: NICK the CAMEL ate a CLAM for SUPPER in PHOENIX “ate” reminds you that we’re looking at the –ates Capital words only First two letters = Ion name NI = nitrate, CA = carbonate, CL = chlorate, SU = sulphate, PH = phosphate Consonants = # of oxygens eg. NICK has 3 consonants: NO3 Vowels = # of negative charges eg. NICK has 1 vowel: NO3- 38 Examples ** If there is a multiple number of the polyatomic ion, it is placed inside brackets Ammonium sulfate Aluminum nitrate 39 Acids For common acids (Hydrogen plus Halogen) : eg: HCl Add prefix – hydro to name of second element Replace end of name with –ic acid Hydrochloric acid 40  Oxoacids Acids can be made using the oxoanions discussed earlier. Nomenclature: Don’t use hydro as a prefix Oxoanions with names ending in –ate become acid named –ic acid ie: Sulfate (SO42-) becomes Sulfuric acid (H2SO4) Oxoanions with names ending in –ite become acids names –ous acid ie: Hypochlorite (ClO-) becomes Hypochlorous acid (HClO) 41