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Atoms, Molecules and Ions Chapter 2 L.Pillay, 2010 1 Dalton - 1808 • Atoms of different elements could be distinguished by differences in their weights. – – – – – – All matter is composed of atoms. Atoms cannot be made or destroyed. All atoms of the same element are identical. Different elements have different types of atoms. Chemical reactions occur when atoms are rearranged. Compounds are formed from atoms of the constituent elements. 2 Thomson - 1897 • Discovered the electron. • Calculated charge to mass ratio • Plum pudding model Millikan - 1909 • • • • Oil Drop Experiment Measured the charge and mass of the electron Charge: 1.602 x 10-19 C Mass: 9.10 x 10-28 g L.Pillay, 2010 3 Rutherford -1911 & Chadwick 1932 • Experimental work with α- particles • The positive charge must be concentrated in a tiny volume at the centre of the atom, while the electrons orbited. • 1919 – Discovered positive particles (protons). • Chadwick - neutrons L.Pillay, 2010 4 Atoms • Proton + Electron + Neutron (1.602 x 10-19 C) (-1.602 x 10-19 C) +1 -1 • Atomic mass is small so we use atomic mass units (amu) represented by u. L.Pillay, 2010 5 Atomic Mass Charge Proton Electron Neutron Mass Actual in C Relative 1.602 x 10-19 +1 Actual in g 1.673 x 10-24 Relative in u 1.00727 1.602 x 10-19 0 9.109 x 10-28 1.675 x 10-24 0.00054858 1.00866 -1 0 L.Pillay, 2010 6 ↔ L.Pillay, 2010 7 No of atoms in a teaspoon of salt L.Pillay, 2010 8 Differentiating elements • Each element has a characteristic number of PROTONS – called atomic number (Z). • The number of protons = no. electrons • No of neutrons can be different for an element. • The number of protons + neutrons = mass number (A). • Atoms with differing numbers of neutrons are called isotopes. L.Pillay, 2010 9 Mass number = Atomic number = Mass number = protons + neutrons No of protons = No. of electrons = No of neutrons = Mass number = Atomic number = Atomic number = no. of protons/ electrons No of protons = No. of electrons = L.Pillay, 2010 No of neutrons = 10 Isotope examples… Isotope Mass number Hydrogen -1 1H Hydrogen -2 2H Hydrogen -3 3H Protons Electrons Neutrons 11 Cobalt Isotope Mass number Cobalt -56 56Co 59 Cobalt -57 57Co 60 Cobalt -58 58Co 61 Protons L.Pillay, 2010 Electrons Neutrons 12 Atomic Mass • Mass of an atom is measured relative to carbon • 12 u = 12C • e.g. 1H = 1.008 u (12 x lighter than 12C) 16O = 15.9949 u • Since most elements occur as a mixture of isotope – use AVERAGE atomic mass L.Pillay, 2010 13 Average atomic mass e.g. Calculate the average atomic mass of Ag Isotope 107Ag 109Ag Mass / u 106.90509 108.90476 % Abundance 51.84 48.16 From 107Ag = From 109Ag = Atomic mass = L.Pillay, 2010 14 Try 1. 52.0 % of Br has 44 neutrons= mass 79.0 [ 48.0 % of Br has 46 neutrons = mass 81.0 [ Calculate the atomic mass of Br. ] ] 2. Boron (Z = 5) has two naturally occurring isotopes. Find the % abundance of 10B and 11B given the following: 10B 11B Isotopic mass 10.129 u 11.0093 u L.Pillay, 2010 Atomic mass 10.81 u 15 The periodic table 16 The periodic table Atomic number Atomic symbol Atomic weight L.Pillay, 2010 18 Chemical Compounds Chapter 3 L.Pillay, 2010 19 Molecules and Molecular Compounds • A molecule consists of two or more atoms bound together • Each molecule has a chemical formula which indicates – Which atoms are found in the molecule, and – In what proportion they are found. • A molecule made up of two of the same atoms is called a diatomic molecule. (Halogens – e.g. Cl, Br) • Compounds composed of molecules are molecular compounds. • Most molecular substances contain only nonmetals. L.Pillay, 2010 20 Examples H2 H–H diatomic H2O H–O–H NH3 polyatomic CH4 L.Pillay, 2010 21 Molecular and Empirical Formulas • Molecular formulas: Give the actual numbers and types of atoms in a molecule. • Empirical formulas : Give the relative numbers and types of atoms in a molecule (they give the lowest whole-number ratio of atoms in a molecule). Molecular Formula Empirical Formula H2O H2O H2O2 HO C2H4 CH2 C6H12O6 CH2O 22 Ions and Ionic Compounds • If electrons are added to or removed from a neutral atom, an ion is formed. • When an atom or molecule loses electrons it becomes positively charged. – Positively charged ions are called cations (M+) 11 e11 p+ Na atom 10 e11 p+ Na+ ion 23 When an atom or molecule gains electrons it becomes negatively charged. – Negatively charged ions are called anions (M-). 17 e- 18 e- 17 p+ 17 p+ Cl atom Cl- ion • Generally, metal atoms tend to lose electrons and non-metal atoms gain electrons. L.Pillay, 2010 24 Predicting Ionic Charges • An atom or molecule can gain or lose more than one electron. • Many atoms gain or lose enough electrons to have the same number of electrons as the nearest noble gas (group 8A). L.Pillay, 2010 25 Ionic Charges +3 ±4 -3 +1 +2 L.Pillay, 2010 -2 -1 0 26 Examples • • • • P Al Ca Si L.Pillay, 2010 27 Oxidation state • Used to determine the ratio in which the elements forming a compound will exist. • Use the ionic charge. • Follow a number of rules L.Pillay, 2010 28 Oxidation State 1. Free elements are assigned an oxidation state of 0. 2. The oxidation state for any simple one-atom ion is equal to its charge. 3. The alkali metals (Li, Na, K, Rb, Cs and Fr) in compounds are always assigned an oxidation state of +1. 4. Fluorine in compounds is always assigned an oxidation state of -1. L.Pillay, 2010 29 5. The alkaline earth metals (Be, Mg, Ca, Sr, Ba, and Ra) and also Zn and Cd in compounds are always assigned an oxidation state of +2. Similarly, Al & Ga are always +3. 6. Hydrogen in compounds is assigned an oxidation state of +1. 7. Oxygen in compounds is assigned an oxidation state of -2. Exception - Peroxide, e.g. H2O2 (O = -1). L.Pillay, 2010 30 8. The sum of the oxidation states of all the atoms in a species must be equal to net charge on the species. e.g. Net Charge of CrO42- = -2, For Cr: x + (- 4)x 2 = -2, →x = +6, so the oxidation state of Cr is +6. • K2MnO4 K = +1 + 2 contribution O = -2 Mn = [2 x (+2)] + [ 4 x (-2)] = +6 • B4O72B=? O = -2 -14 contribution BOS + (-14) = -2 B4 = 12 and each B = +3 L.Pillay, 2010 -8 contribution Net charge = -2 31 Try Determine the oxidation state of 1. P in PCl5 2. Cl in Ca(ClO4)2 3. S in S2O324. S in Na2SO3 L.Pillay, 2010 32 Charges for common polyatomic anions • You will be taught HOW to determine these charges in a few weeks. Anion OH NO3 2SO4 CO32PO42- Charge -1 -1 -2 -2 -2 L.Pillay, 2010 33 Ionic compounds • Form when cations and anions attract each other to form a neutral compound. • Ratio of ions can differ. • An ionic compound must have a metal present. – H classified as nonmetal. e.g. Na+ + Cl- → NaCl (+1 from Na and -1 from Cl = 0 net charge) e.g. Al3+ and O2- → Al2O3 ( 2x +3 from Al and 3 x -2 from O = 0 net charge L.Pillay, 2010 34 Al2O3 … • What does this mean? • 2 Al3+ ions and 3 O2- ions make up the molecule. Try : Ca2+ and NO3Ca: 2+ NO3: -1 1 Ca + 2 NO3 = +2 + (-1)2 = 0 Ca(NO3)2 () implies the subscript applies to ALL the atoms within. 35 Molecular Ionic H + non-metal Metal + non-metal HCl, C2H4 CaCl2, LiF L.Pillay, 2010 Ionic bonds (transferring electrons) Covalent bonds (sharing electrons) Compounds 36 Naming inorganic ionic compounds Cations a) Cations formed from metal ions have the same name as the metal b) The positive charge of the metal is indicated by a roman numeral in brackets. e.g. Arsenic can exist with a +3 charge and a + 5 charge. This is represented as As(III) and As(IV) L.Pillay, 2010 37 c) Cations formed from non-metal atoms have names that end in –ium e.g. NH4+ ammonium ion H3O+ hydronium ion Anions a) Monoatomic ions have names ending in –ide e.g. I Iodine iodide BrBromine bromide ** OHhydroxide CNcyanide O22peroxide L.Pillay, 2010 38 b) Polyatomic anions with O have names ending in –ate or –ite Same charge but one less O Most common oxy-anion SO42- Sulfate ion SO32- Sulfite ion c) When an H+ is added to an oxyanion add hydrogen (1H) or dihydrogen (2H) to the name. e.g. H2SO4 dihydrogen sulfate H2O dihydrogen oxide L.Pillay, 2010 39 ** Some compounds can have multiple oxyanion forms ClO- ClO2- ClO3- ClO4- The number of O's relative to the Chlorine is changing, but the ionic charge is not (-1 for all). The -ite and -ate suffixes are still used, but to distinguish between the four forms we use prefixes: ClOClO2ClO3ClO4- hypochlorite ion chlorite ion chlorate ion perchlorate ion L.Pillay, 2010 40 e.g. CuBr2 NH4ClO3 KHSO4 NaClO2 Lithium sulfite Barium phosphate Try 1. Sodium persulphate 2. Magnesium(II) bromide 3. Potassium cyanide L.Pillay, 2010 41 Names and formulas of Acids Acids • An acid is a substance whose molecules yield hydrogen (H+) ions when dissolved in water. • Acid = anion + H+ • The name of the acid is related to the name of the anion L.Pillay, 2010 42 • Anions whose names end in -ide have associated acids that have the hydro- prefix and an -ic suffix: Cl- chloride anion HCl hydrochloric acid S2- sulfide anion H2S hydrosulfuric acid • For naming the acids of oxyanions: – If the anion has an -ate ending, the corresponding acid is given an -ic ending. – If the anion has an -ite ending, the corresponding acid has an -ous ending. • Prefixes in the name of the anion are kept in naming the acid ClO- hypochlorite ion HClO hypochlorous acid ClO2- chlorite ion HClO2 chlorous acid ClO3- chlorate ion HClO3 chloric acid 43 ClO4- perchlorate ion HClO4 perchloric acid Rules for Binary Molecular Compounds • The naming system is for compounds composed of two non-metallic elements. • The element furthest left is named first. (except when O and a halogen!) Prefixes • The first element gets a prefix if it has a 1 – mono subscript in the formula. – The second element gets the –ide suffix (ending) – The second element ALWAYS gets a prefix L.Pillay, 2010 2 – di 3 – tri 4 – tetra 5 – penta 6 – hexa 7 - hepta44 Examples 1. N2O4 2. P4O10 3. PBr5 4. nitrogen trichloride 5. disulfur hexafluoride L.Pillay, 2010 Prefixes 1 – mono 2 – di 3 – tri 4 – tetra 5 – penta 6 – hexa 7 - hepta Simple Organic Compounds Alkanes • Contain only C and H • Name according to number of carbons, end in –ane. • After 4 carbons names as for other compounds. L.Pillay, 2010 46 Alcohols • When an H is replaced by –OH group. • -ol added of end of name. Methanol ethanol 1-propanol 1- butanol 1-pentanol 47 You should be able to … • understand the concept of the atom. • differentiate elements according to mass number. • use the periodic table to determine the number of protons, electrons and neutrons an element has. • explain what an isotope is. • calculate the average atomic mass of an element. • identify basic groups in the periodic table. • understand and differentiate between ions, molecules and their compounds. • predict the charge for ions. • Assign oxidation states to elements in a charged or neutral compound • Determine the name or the chemical formula for ionic and binary molecular compounds, acids and simple alkanes. KNOW • The charges for common polyatomic anions. • The rules for naming ionic and binary molecular compounds, acids and simple alkanes. L.Pillay, 2010 48