Chapters 1-4. Numbers and Measurements in Chemistry. SI Prefixes. Units. Dimensional Analysis

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1 Chapters 1-4 What is Chemistry? Chemistry is the study of the composition, structure, properties and reactions of matter (the physical material of the universe). A main challenge of chemistry is to bridge the macroscopic world we experience with the microscopic structure of matter (atoms and molecules). Numbers and Measurements in Chemistry Chemists quantify data, expressing collected data with units and significant figures. Units - designate the type of quantity measured. Prefixes - provide scale to a base unit. Significant Figures - indicate the amount of information that is reliable when discussing a measurement. Units SI Prefixes The base unit designates the type of quantity being measured. SI units (from French Système International) are the base units of science. Some units comprise combinations of these base units and are termed derived units 1 J = 1 kg m 2 s -2 Prefixes are based on multiples of 10. Dimensional Analysis It may sound obvious (and yet is often forgotten) many mistakes can be avoided if an answer is checked for: units order of magnitude sign Dimensional Analysis Example What is the mass of a cube of osmium that is 32 mm on each side? Conversion of a result from one system of units to another is called dimensional analysis: 1. To carry out dimensional analysis, we must know the relationship between units (equivalents): e.g. 1 nm = 10-9 m; 2. Use equivalents to determine unit factors: e.g. 1 = 1 nm/10-9 m; 3. Multiply result by appropriate unit factor(s) to convert units. 1

2 Atomic Theory Each element is made up of tiny particles called atoms. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. Chemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. Atomic Theory Atomic theory is the foundation of modern chemistry.supported today by direct visualization (and manipulation) of individual atoms using scanning tunnelling microscopes (STM): IBM spelled out in individual atoms on a surface the first nano-advertisement Fundamental Chemical Laws Law of conservation of mass Mass is neither created nor destroyed Law of definite proportion A given compound always contains exactly the same proportion of elements by mass Law of multiple proportions When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. Atomic Structure and Mass The atom contains: Electrons found outside the nucleus; negatively charged. Protons found in the nucleus; positive charge equal in magnitude to the electron s ec negative e charge. Neutrons found in the nucleus; no charge; virtually same mass as a proton. The nucleus is: Small compared with the overall size of the atom. Extremely dense; accounts for almost all of the atom s mass. The Modern View of Atomic Structure The atom contains: electrons protons: found in the nucleus; positive charge equal in magnitude to the electron s negative charge. neutrons: found in the nucleus; no charge; virtually same mass as a proton. Atomic Number and Mass Number Atomic Number, Z, is the number of protons in a nucleus. identifies the element Mass Number, A, is the sum of the number of protons and number of neutrons in a nucleus. 1 amu = x g Proton charge = 1.602x10-19 C Electron charge = x10-19 C Protons and neutrons are nearly 2000 times more massive than electrons Particle Mass (amu) Charge Proton Neutron Electron Heaviest atom is ~260 amu (4x10-22 g) Largest atom is 500 pm across. Typical C-C bond length 154 pm 2

3 Isotopes Isotopes are atoms of an element that differ in the number of neutrons in their nucleus. same Z but different A Isotopic abundance is the mass percentage of an isotope in a naturally occurring element. Show almost identical chemical properties; chemistry of atom is due to its electrons. In nature most elements contain mixtures of isotopes. Atomic Symbols Information regarding atomic structure is written in scientific shorthand called the atomic symbol. A Z E E is the atomic symbol for element Superscript A is the mass number. Subscript Z is the atomic number, it is redundant and is often left off. Atomic Masses Entry for carbon on the periodic table. Z = 6 Atomic Masses Relative atomic mass for an element is an average of the atomic masses for the naturally occurring isotopes for an element. 12 6C Relative atomic mass = (~99% carbon-12) Element Symbol: C Carbon-12 = x = amu Carbon-13 = x = amu Average mass = = amu Isotopes Two Isotopes of Sodium Mass spectrum showing carbon isotopes. 3

4 Ions Ions are formed when the number of protons and electrons in an atom are not equal. Ions with more protons than electrons are called cations. net positive charge Ions with more electrons that protons are called anions. net negative charge A monatomic ion is derived from a single atom. Ca 2+, Cl - A polyatomic ion is derived from a group of atoms with an overall charge. NH 4+, SO 4 2- Chemical Formulas Chemical formulas describe a compound in terms of the elements the compound contains. The number of atoms for each element is indicated by a subscript to the right of the chemical symbol. Groups of atoms can be designated using parentheses. Subscripts outside these parentheses mean that all atoms enclosed in the parentheses are multiplied by the value indicated by the subscript. Water molecules associated with certain compounds, called hydrates, are indicated separately from the rest of the compound. Fe 3 (PO 4 ) 2 8H 2 O Chemical Formulas Compounds have different properties than their constituent atoms. Ionic compounds contain cations and anions, usually arranged in a lattice. Molecular l formulas indicate the elements and number of atoms of each element actually contained in a discrete unit of a compound. Empirical formulas indicate the smallest whole number ratio between the number of atoms of each element in a molecular formula. Chemical Formulas The molecular formula for ethylene is C 2 H 4. The empirical formula for ethylene is CH 2. Polyethylene l can be written as [CH 2 CH 2 ] n n is used to emphasize that a large number of these units are found in an individual molecule Chemical Bonding The Periodic Table All bonds are created by the exchange or sharing of electrons. The exchange or sharing of electrons results in lower energy for the compound relative to the separate atoms. Ionic: exchange; cations and anions present Metallic: sharing by forming a mobile sea of electrons Covalent: sharing between atoms The Periodic Table is based on periodic law. Periodic law - when arranged properly, the elements display a regular and periodic variation in their chemical properties. Periods are horizontal rows on the periodic table. Groups are vertical columns on the periodic table. 4

5 Periodic Table Periods and Groups Common names of specific groups: Group 1: alkali metals. Group 2: alkaline earth metals. Group 17: halogens. Group 18: noble gases/rare gases. Table Regions: Groups 1-2 and are main group elements Groups 3-12 are transition metals Lanthanides and actinides are below the rest of the table Special Names for Groups in the Periodic Table Metals, Nonmetals, and Metalloids Metals are generally toward the left and bottom of the periodic table. They are shiny, malleable, and ductile. They conduct current and easily form cations. Nonmetals occupy the upper right-hand portion of the periodic table. They are not shiny, malleable, or ductile. They do not conduct current but do easily form anions. Metalloids, or semimetals, have chemical properties intermediate of metals and nonmetals. Metalloids are clustered along a diagonal line on the periodic table between the metals and nonmetals. Inorganic and Organic Chemistry Organic Chemistry Organic chemistry is the study of the compounds of the element carbon, usually with oxygen, nitrogen, and hydrogen. More than 18 million organic compounds exist. Includes biological i l molecules l and nearly all synthetic ti polymers. Isomers: Different organic molecules that have the same formula but are connected differently. Because carbon compounds can become quite large, organic compounds are described simply and unambiguously using line structures, where carbons and hydrogens are not explicitly shown. Each corner or end of a line is a carbon. Hydrogen atoms on carbon atoms are implied. Carbon makes four bonds, missing bonds go to hydrogen atoms. Hydrogen can only make one covalent bond to another atom. Hydrogen atoms on any other element are shown All other elements are shown Inorganic chemistry is the study of all other elements and their compounds. H 2 N CH 2 C OH becomes H 2 N OH O O 5

6 Functional Groups Functional groups are arrangements of atoms that tend to display similar chemical properties. Alkanes Alkanes have the general formula C n H 2n+2 where n is an integer. Chemical formulas are often written to emphasize functional groups. Methanol, an alcohol, is often written CH 3 OH instead of CH 4 O. Hydrocarbons contain only H and C atoms. Addition of functional groups to hydrocarbons results in more complex compounds. Alkanes are hydrocarbons where the carbon atoms are linked together with single bonds. Isomers Functional Groups Isomers are compounds that have the same chemical formula but are connected differently. Three isomers of pentane, C 5 H 12. One straight chain Two branched chains Functional Groups Polymers Polymers are very large molecules made up of many smaller molecules linked together. Monomers - The small molecules linked together in polymers. Polymer backbone - The long chain of bonded atoms formed when monomers link together to form polymers. 6

7 Polyethylene Ethylene monomers are linked together via a free radical mechanism, which converts the carboncarbon double bond to a single bond. Polyethylene The ethylene free radical reacts with another ethylene monomer, extending the polyethylene chain. Free radicals have an unshared single electron and are extremely reactive. An initiator produces the free radical that reacts with ethylene, opening the double bond and transferring the free radical to the ethylene monomer. Polymerization continues until the free radical reacts with another free radical, or terminator, which terminates the growth of the chain. Polyethylene Polymers Polyethylene polymers can be linear chains (highdensity polyethylene, HDPE) or branched chains (low-density polyethylene, LDPE). Ultra-high molecular weight polyethylene, UHMWPE, contains extremely long chains. Polymers are the materials of choice for a host of everyday objects. a) polyethylene b) polystyrene c) poly(vinyl chloride), PVC Polymers Models showing how atoms are arranged in several polymers. Each of these polymers has distinct properties. Chemical Equations A representation of a chemical reaction: C 2 OH + 3O 2 2CO 2 + 3H 2 O reactants products Reactants are only placed on the left side of the arrow, products are only placed on the right side of the arrow. The equation is balanced. All atoms present in the reactants are accounted for in the products. 1 mole of ethanol reacts with 3 moles of oxygen to produce 2 moles of carbon dioxide and 3 moles of water. The balanced equation represents an overall ratio of reactants and products, not what actually happens during a reaction. Use the coefficients in the balanced equation to decide the amount of each reactant that is used, and the amount of each product that is formed. 7

8 Avogadro s Number and the Mole A mole is a means of counting the large number of particles in samples. One mole is the number of atoms in exactly 12 grams of 12 C (carbon-12). 1 mole contains Avogadro s number (6.022 x particles/mole) of particles. The mass of x atoms of any element is the molar mass of that element. Balanced chemical reactions also provide mole ratios between reactants and products. 2H 2 (g) + O 2 (g) 2H 2 O(g) 2 moles H 2 : 1 mole O 2 : 2 moles H 2 O Determining Molar Mass The molar mass of a compound is the sum of the molar masses of all the atoms in a compound. 2 mol H 1.0 g H 1 mol H + 1 mol O 16.0 g 1 mol O = 18.0 g/mol H 2 O Calculations Using Moles and Molar Mass Molar mass allows conversion from mass to number of moles, much like a unit conversion. 1 mol C 7 = g C 7, (TNT) Calculations Using Moles and Molar Mass Avogadro s number functions much like a unit conversion between moles to number of particles. 1 mol C 7 = C 7 molecules How many molecules are in moles of TNT,? 1 mol C g C g C mol C molecules C 7 1 mol C 7 = mol C 7 = molecules C 7 Molar Example Fundamentals of Stoichiometry Potassium-40 is one of the few naturally occurring radioactive isotopes of elements of low atomic number. Its percent natural abundance among K isotopes is 0.012%. How many 40 K atoms do you ingest by drinking one cup of whole milk containing 371 mg of K? Stoichiometry is a term used to describe quantitative relationships in chemistry. How much? of a product is produced or reactant is consumed. Balanced chemical equation needed. Conversion between mass or volume to number of moles frequently needed. 8

9 Ratios from a Balanced Chemical Equation Ratios from a Balanced Chemical Equation This flow diagram illustrates the various steps involved in solving a typical reaction stoichiometry problem. Mole ratios are obtained from the coefficients in the balanced chemical reaction. 1 mol CH 4 : 2 mol O 2 : 1 mol CO 2 : 2 mol H 2 O These ratios can be used in solving problems: No different than unit conversion Usually more than one conversion is necessary Write all quantities with their complete units 1 mol CH 4 or 2 mol H O 2 2 mol O 2 1 mol CH 4 Ratios from a Balanced Chemical Equation Limiting Reactants How many grams of water can be produced if sufficient hydrogen reacts with 26.0 g of oxygen? In many chemical reactions, one reactant is often exhausted before the other reactants. This reactant is the limiting reactant. Limiting iti reactant t is determined d using stoichiometry. t The limiting reactant limits the quantity of product produced. Limiting Reactants Introductory Examples 2H 2 (g) + O 2 (g) 2H 2 O(g) Reaction between 6 H 2 and 6 O 2 will produce 6 H 2 O. 1. How many copper atoms are there on the end surface of 18-gauge wire? 6 H 2 can produce 6 H 2 O. 6 O 2 can produce 12 H 2 O. H 2 is limiting reactant. 3 O 2 left over. 2. Oxygen can be generated by heating a solid if the oxygen is loosely bound in the solid. KClO 3 is an example of such a solid. a. Write a balanced chemical equation for the decomposition of KClO 3 into KCl and O 2. b. How many grams of O 2 can be produced from 10. g of KClO 3? 9