SCI1055 Chapter 9 PDF

1.2 States of Matter (4) Physical properties can be observed or measured without changing the composition of the material. boiling point color melting point odor solubility state of matter 1 ...

1.2 States of Matter (4) Physical properties can be observed or measured without changing the composition of the material. boiling point color melting point odor solubility state of matter 1  1.2 States of Matter (5) A physical change alters the material without changing its composition. 2  1.2 States of Matter (6) Chemical properties determine how a substance can be converted into another substance. Chemical change is the chemical reaction that converts one substance into another 3  1.3 Classification of Matter (1) All matter can be classified as either a pure substance or a mixture. Pure Substances A pure substance is composed of only a single component (atom or molecule). It has a constant composition, regardless of sample size or origin of sample. It cannot be broken down to other pure substances by a physical change. 4  1.3 Classification of Matter (5) A pure substance is classified as an element or a compound. An element is a pure substance that cannot be broken down by a chemical change. 5 a: ©Daniel C. Smith; b: ©Keith Eng, 2008  1.3 Classification of Matter (6) A pure substance is classified as an element or a compound. A compound is a pure substance formed by chemically joining two or more elements. 6 c: ©McGraw-Hill Education/Jill Braaten, photographer; d: ©Daniel C. Smith  1.3 Classification of Matter (2) All matter can be classified as either a pure substance or a mixture. Pure Substances Table sugar(C12H22O11) and water (H2O) are both pure substances: 7  1.3 Classification of Matter (4) All matter can be classified as either a pure substance or a mixture. Mixtures Sugar dissolved in water is a mixture. 8  1.3 Classification of Matter (3) All matter can be classified as either a pure substance or a mixture. Mixtures Mixtures are composed of more than one component. They can have varying composition (any combination of solid, liquid, and gas). Mixtures can be separated into their components by a physical process. 9 1.3 Classification of Matter (7) 10  Atoms and molecules Atom - smallest particle that retains the chemical properties of an element Molecule smallest particle that retains the chemical properties of a compound Examples: oxygen, hydrogen gas - diatomic molecules Ozone - triatomic oxygen molecule Examples of molecules  3.1 Introduction to Bonding (1) Bonding is the joining of two atoms in a stable arrangement. Elements will gain, lose, or share electrons to reach the electron configuration of the noble gas closest to them in the periodic table. There are two different kinds of bonding: Ionic bonds result from the transfer of electrons from one element to another. Covalent bonds result from the sharing of electrons between two atoms. 13  3.1 Introduction to Bonding (2) Ionic bonds form between: A metal on the left side of the periodic table. A nonmetal on the right side of the periodic table. Sodium metal, Chlorine gas: © The McGraw-Hill Companies, Inc./Stephen Frisch, photographer; Sodium chloride crystals: © Dane S. Johnson/Visuals Unlimited 14  3.2 Ions (1) A. Cations and Anions Ions are charged species in which the number of protons and electrons in an atom is unequal. Ionic compounds consist of oppositely charged ions that have a strong electrostatic attraction for each other. There are two types of ions—cations and anions. 15  3.2 Ions (2) A. Cations and Anions Cations are positively charged ions. A cation has fewer electrons (e−) than protons. the sodium atom the sodium ion 16  3.2 Ions (3) A. Cations and Anions By losing one, two, or three e−, an atom forms a cation with a completely filled outer shell of e−. the magnesium atom the magnesium ion 17  3.2 Ions (4) A. Cations and Anions Anions are negatively charged ions. An anion has more e− than protons. the chlorine atom the chloride ion 18  3.2 Ions (5) A. Cations and Anions Metals, like sodium (Na) and magnesium (Mg), form cations. By losing one, two, or three electrons, an atom forms a cation with a completely filled outer shell of electrons. Nonmetals, like chlorine (Cl), form anions. By gaining one, two, or three electrons, an atom forms an anion with a completely filled outer shell of e−. The octet rule: a main group element is especially stable when it possesses an octet of e− in its outer shell. octet = 8 valence e− 19  3.2 Ions (6) B. Relating Group Number to Ionic Charge for Main Group Elements Elements in the same group form ions of similar charge. For metals in groups 1A, 2A, and 3A, the group number = the charge on the cation. For nonmetals in Groups 6A and 7A, the anion charge = 8 − the group number. 20  3.2 Ions (7) B-1 Relating Group Number to Ionic Charge for Groups 1A–3A the cation charge = the group number group 1A: group 2A: group 3A: 21  3.2 Ions (8) B-2 Relating Group Number to Ionic Charge for Groups 6A and 7A the anion charge = 8 − group number group 6A: group 7A: 22 3.2 Ions (9) 23  3.3 Ionic Compounds (1) An ionic bond is formed when a metal transfers one or more electrons to a nonmetal. The sum of the charges in an ionic compound must be zero overall. 24 3.3 Ionic Compounds (2) 25  3.3 Ionic Compounds (3) HOW TO Write a Formula for an Ionic HOW TO Write a Formula for an Ionic Compound Compound Step Step Identify which element is the cation and which is the anion. Metals form cations and nonmetals form anions. Use the group number of a main group element to determine the charge. K Cl  Ca 2 O 2 metal nonmetal metal nonmetal group 1A group 7A group 2A group 6A 26  3.3 Ionic Compounds (4) HOW TO Write a Formula for an Ionic HOW TO Write a Formula for an Ionic Compound Compound Step Step Determine how many of each ion type is needed for an overall charge of zero. When the cation and anion have the same charge, only one of each is needed.   K  Cl KCl zero charge 2 2 Ca  Ο CaO zero charge One of each ion is needed to balance charge. 27  3.3 Ionic Compounds (5) HOW TO Write a Formula for an Ionic HOW TO Write a Formula for an Ionic Compound Compound When the cation and anion have different charges, use the ion charges to determine the number of ions of each needed. Ca 2 Cl  A +2 charge means A −1 charge means 2 Cl  anions are 1Ca 2 cation is needed. needed. Ca 2  Cl  CaCl2 2Cl  for each Ca 2 28  3.3 Ionic Compounds (6) HOW TO Write a Formula for an Ionic HOW TO Write a Formula for an Ionic Compound Compound Step Step To write the formula, place the cation first and then the anion, and omit charges. Examples: KCl CaO CaCl2 Use subscripts to show the number of each ion needed to have a zero overall charge. When no subscript is written, it is assumed to be “1.” 29  3.4 Naming Ionic Compounds (1) A. Naming Cations Main group cations (groups 1—3A) are named for the element from which they are formed. Na  K Ca 2 Mg 2 sodium potassium calcium magnesium 30  3.4 Naming Ionic Compounds (4) B. Naming Anions Anions are named by replacing the ending of the element name by the suffix “-ide.” Table 3.4 Names of Common Anions Element Ion Symbol Name Bromine Br  Bromide Chlorine Cl Chloride Fluorine F Fluoride Iodine I Iodide Nitrogen N 3 Nitride Oxygen O 2 Oxide Phosphorus P3 Phosphide Sulfur S2 Sulfide 31  3.4 Naming Ionic Compounds (5) C. Compounds of Main Group Metals Name the cation and then the anion. Do not specify the charge on the ion. Do not specify how many ions of each type are needed to balance charge. Na  F NaF  sodium fluoride sodiumfluoride Mg  Cl  MgCl2  magnesium chloride magnesiumchloride 32 4.1 Introduction to Covalent Bonding (1) Covalent bonds result from the sharing of electrons between two atoms. A covalent bond is a two-electron bond in which the bonding atoms share valence electrons. A molecule is a discrete group of atoms held together by covalent bonds. 33 4.1 Introduction to Covalent Bonding (2) Unshared electron pairs are called nonbonded electron pairs or lone pairs. Atoms share electrons to attain the electronic configuration of the noble gas closest to them in the periodic table. H shares 2 e−. Other main group elements share e− until they reach an octet of e− in their outer shell. 34  4.1 Introduction to Covalent Bonding (3) A. Covalent Bonding and the Periodic Table Lewis structures are electron-dot structures for molecules. They show the location of all valence e−. 35  4.2 Lewis Structures (8) B. Multiple Bonds One lone pair of e− can be converted into one bonding pair of e− for each 2 e− needed to complete an octet on a Lewis Structure. A double bond contains four electrons in two 2 e− bonds. A triple bond contains six electrons in three 2 e− bonds. 36 4.5 Naming Covalent Compounds (1) HOW TO Name a Covalent HOW TO Name a Covalent Molecule Molecule Exampl Example Name each covalent molecule: e a) NO2 b) N2O4 Step Name the first nonmetal by its element Step name and the second using the suffix “-ide.” a) NO2 b) N2O4 nitrogen oxide nitrogen oxide 37 4.5 Naming Covalent Compounds (2) HOW TO Name a Covalent HOW TO Name a Covalent Molecule Molecule Step Add prefixes to show the number of Step atoms of each element. Use a prefix from Table 4.1 for each element. The prefix “mono-” is omitted when only one atom of the first element is present, but it is retained for the second element. If the combination would place two vowels next to each other, omit the first vowel. mono + oxide = monoxide (not monooxide) 38 4.5 Naming Covalent Compounds (3) HOW TO Name a Covalent HOW TO Name a Covalent Molecule Molecule Table 4.1 Common Prefixes in Nomenclature Number a) NO2 of Atoms Prefix 1 Mono 2 Di nitrogen dioxide 3 Tri 4 Tetra 5 Penta 6 Hexa b) N2O4 7 Hepta 8 Octa dinitrogen tetroxide 9 Nona 10 Deca 39  Chemical Equations (1) Represents chemical reaction using chemical symbols Reactants on left Products on right Arrow indicates change Arrow “yields” or “produces”  Chemical Equations (2) Use chemical symbols and formulas for each reactant and product Correct equation should be balanced obey Law of Conservation of Mass equal numbers of atoms of each element on reactant side as on product side CH CH 4 + O22  4 O CO22 + CO  HH2O 2O(unbalanced) (unbalanced) CHCH 4 + 2Ο22  4 2O CO22 + CO  2H 2H2O 2O(balanced) (balanced) Writing chemical equations One atom of C means carbon O means One atom of oxygen One molecule of oxygen consisting O2 means of two atoms of oxygen One molecule of carbon means monoxide consisting of one atom CO of carbon attached to one atom of oxygen One molecule of carbon dioxide CO2 means consisting of one atom of carbon attached to two atoms of oxygen Three molecules of carbon dioxide, each consisting of one 3 CO2 means atom of carbon attached to two atoms of oxygen  5.3 Types of Reactions (1) The majority of chemical reactions fall into the following categories: oxidation/reduction (redox) Electrons are transferred from one atom to another The below reactions can also be redox combination decomposition single replacement double replacement 43  5.3 Types of Reactions (2) A. Combination and Decomposition A combination reaction is the joining of two or more reactants to form a single product. Table 5.2 Examples of Combination Reactions 1) N 2  3H 2 2 NH 3 2) Ca  Br2 CaBr2 44 3) H 2 C CH 2  Cl 2 ClCH 2 CH 2 Cl  5.3 Types of Reactions (3) A. Combination and Decomposition A decomposition reaction is the conversion of a single reactant to two or more products. Table 5.3 Examples of Combination Reactions 1) 2 NH 3 N 2  3 H 2 2) 2 KClO3 2 KCl  3 O 2 45 3) CH 3 CH 2 Cl H 2 C CH 2  HCl  5.3 Types of Reactions (4) B. Replacement Reactions A single replacement reaction is a reaction in which one element replaces another element in a compound to form a different compound and element as products. 46  5.3 Types of Reactions (5) B. Replacement Reactions A double replacement reaction is a reaction in which two compounds exchange “parts”–atoms or ions—to form two new compounds. Table 5.4 Examples of Single and Double Replacement Reactions Type of Reaction Examples Comment Single replacement 1) 2 NaCl  Br2 2 NaBr  Cl 2 The element Br replaces CI in the compound NaCI. Single replacement 2) Fe  CuSO 4 FeSO 4  Cu The element Fe replaces Cu in the compound CuSO4. Double replacement 1) AgNO3  NaCl AgCl  NaNO3 The ions Ag and Na exchange.   2) HCl  NaOH H 2 O  NaCl Double replacement The species H and Na  exchange. 47