Lecture Notes - Chemistry 1110Dr. Luther Giddings

Last Updated: August 22, 2020

Table of ContentsChapter 1: Matter and Measurements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1

What is chemistry? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1Why is chemistry relevant? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2What is matter? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3How is matter classified? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4What kinds of properties does matter have? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6An introduction to the elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7The natural states of the elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9The Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11Chemistry is an empirical science . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13Scientific notation and powers of 10 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13Measurements and significant figures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15SI units . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18Measurements and measured quantities . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19Derived quantities-volume . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 20Derived quantities-density . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 20Energy and heat . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21Dimensional analysis and problem solving . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22Organic or inorganic - a postscript . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 26

Chapter 2: Atoms and the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28The structure of the atom . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28The Nucleus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28Atoms and protons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 29Atoms, neutrons, and isotopes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 30Atoms and electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 34The octet rule . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35Fundamentals of electrons in atoms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37Ground states and excited states . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42Electron configurations: full configurations and the periodic table . . . . . . . . . . . . . . . . 42Electron configurations: Noble Gas configurations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 52Electron configurations: electron spin diagrams (arrow diagrams) . . . . . . . . . . . . . . . . . 53Inner shell, outer shell, and valence electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 58The electron configurations of ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 64Atomic properties and periodic trends . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65

Chapter 3: Ionic Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69Chemical bonds and chemical compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69Lewis electron dot structures (Lewis structures) for atoms . . . . . . . . . . . . . . . . . . . . . . . 72Ionic compound nomenclature: naming ionic compounds . . . . . . . . . . . . . . . . . . . . . . . 73Ionic compounds: cation names . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 74Ionic compounds: monatomic and polyatomic anion names . . . . . . . . . . . . . . . . . . . . . . 75A digression: old ionic compound nomenclature systems . . . . . . . . . . . . . . . . . . . . . . . . 78

The names and molecular formulas of ionic compounds . . . . . . . . . . . . . . . . . . . . . . . . 80Ionic compounds: going from names to molecular formulas . . . . . . . . . . . . . . . . . . . . . 83Ionic compounds: going from molecular formulas to names . . . . . . . . . . . . . . . . . . . . . 86The nomenclature of binary molecular compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 92The nomenclature of acids and bases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 94Common nomenclature problems - some troubleshooting help . . . . . . . . . . . . . . . . . . . 96

Chapter 4: Molecular Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97Covalent bonds and molecular compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97Covalent bonds and the Periodic Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 98Multiple covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 98Coordinate covalent bonds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 99Molecular formulas and structural formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100The Lewis structures of ionic compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103The Lewis structures of covalent compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103Exceptions to the octet rule . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 113The shapes of molecules and VSEPR - Valence Shell Electron-Pair Repulsion . . . . . . 116Two Sets of Electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 119Three Sets of Electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 119Four Sets of Electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 121Five Sets of Electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122Six Sets of Electrons . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 125Polar covalent bonds and polar molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 126Bond polarity and molecular polarity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 128Summary of determining molecular polarity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 128

Chapter 5: Classification and Balancing of Chemical Reactions . . . . . . . . . . . . . . . . . . . . . . . 129Chemical equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 129Balancing chemical equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132Strong electrolytes, weak electrolytes, and non-electrolytes . . . . . . . . . . . . . . . . . . . . . 138Molecular, ionic, and net ionic equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 140Double displacement reactions: precipitation formation and solubility rules . . . . . . . . 143Double displacement reactions and water formation . . . . . . . . . . . . . . . . . . . . . . . . . . . 145Double displacement reactions and gas formation . . . . . . . . . . . . . . . . . . . . . . . . . . . . 148Electron transfer (redox) reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 150Post-script: responses to student questions about Chapter 5 and 6 concepts . . . . . . . . 156

Chapter 6: Chemical Reactions - Mole and Mass Relationships . . . . . . . . . . . . . . . . . . . . . . . . 160Molecular formulas and empirical formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 160Molecular formulas, molecular weights, and formula weights . . . . . . . . . . . . . . . . . . . 161The mole and Avogadro's number . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 162Molecular weight and molar mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 164Stoichiometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 166Theoretical yield and percent yield . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 169

Chapter 7: Chemical Reactions - Energy, Rates, and Equilibrium . . . . . . . . . . . . . . . . . . . . . . 171An introduction to thermodynamics and kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171Energy, work, and other thermodynamic definitions . . . . . . . . . . . . . . . . . . . . . . . . . . 171Internal energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 174Enthalpy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 175Entropy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 178Gibbs free energy and the spontaneity of reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . 180The rates of chemical reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 183Reaction rates and molecular collisions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 184Rate Laws . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 186Activation energy and activated complexes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 186The reversibility of chemical reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 189Catalysts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 190Chemical equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 192The equilibrium constant . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 193Writing equilibrium constant expressions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 197Equilibrium constants and Gibbs free energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 199LeChatelier's principle . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 200

Chapter 8: Gases, Liquids, and Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 204General gas properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 204Ideal gases and the Ideal Gas Law (Universal Gas Law) . . . . . . . . . . . . . . . . . . . . . . . 207The ideal gas law and molar mass calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 209The ideal gas law and stoichiometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 210The combined gas law . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 211Dalton's Law of Partial Pressures and mole fractions . . . . . . . . . . . . . . . . . . . . . . . . . . 216States of matter, phase transitions, and enthalpies of phase transitions . . . . . . . . . . . . 218Intermolecular forces: attractive interactions between molecules . . . . . . . . . . . . . . . . . 222London (dispersion) forces . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 223The magnitude of London forces increases with molecular size and surface area . . . . 225Dipole-dipole interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 226Hydrogen bonds - a special case of dipole-dipole interactions . . . . . . . . . . . . . . . . . . . 227Ion-dipole interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 228Predicting the intermolecular forces in compounds . . . . . . . . . . . . . . . . . . . . . . . . . . . 229Intermolecular forces and solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 233Intermolecular forces and physical properties . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 234

Chapter 9: Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 238General solution properties and definitions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 238Calculating concentration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 240Dilutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 243Some properties of liquids and solutions: surface tension and capillary action . . . . . . 246Some properties of liquids and solutions: vaporization and vapor pressure . . . . . . . . . 247Some properties of liquids and solutions: diffusion . . . . . . . . . . . . . . . . . . . . . . . . . . . 249Some properties of liquids and solutions: osmosis . . . . . . . . . . . . . . . . . . . . . . . . . . . . 250

Chapter 10: Acids and Bases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252General acid-base information: a review . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252The Arrhenius acid-base theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252The Brønsted-Lowry acid-base theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252The Lewis acid-base theory . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 254Strong and weak acids and bases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 257The behavior of acids and conjugate bases in water . . . . . . . . . . . . . . . . . . . . . . . . . . . 258The behavior of bases and conjugate acids in water . . . . . . . . . . . . . . . . . . . . . . . . . . . 262The acid-base behavior of water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 264The pH scale . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 265Calculating [H3O+], [OH-], and pH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 266Buffer solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 270

Chapter 1: Matter and MeasurementsWhat is chemistry?

If you refer to various chemistry text books you will find a variety of definitions ofchemistry. I prefer to define chemistry as the study of things made up of atoms and molecules. Itis the study of why things made up of atoms and molecules behave the way they do. It is also thestudy of how to make them behave more usefully. As everything in the world around us iscomposed from atoms, by gaining a fundamental understanding of chemistry we can begin tounderstand many important and interesting phenomena that affect us on a daily basis.

While one may read of a variety of different types of chemistry such as forensicchemistry, geochemistry, food chemistry, and etc., there are really just five fundamentalbranches of chemistry:

Analytical chemistry is the study of what is in a substance, and how much of a particularthing is in a substance. Most of the things in the world around us are not chemically pure butconsist of dozens or even hundreds or thousands of different chemically distinct components.Gasoline looks like it is simply a liquid, but in fact it is composed of several hundred differentchemicals. The smell of an orange is due to the presence of nearly two hundred differentcompounds. And so on. Analytical chemistry is used all around us, every day, and in a variety ofways. The study of pollutants in the environment, the analysis of the composition of anexpensive perfume, the testing of a patient's blood in a hospital or of an athlete's urine during acompetition, and a check of the composition of the various liquids used during the manufactureof gasoline, are a few examples of how analytical chemistry is used.

Inorganic chemistry is the study of all of the elements and their compounds exceptcarbon and its compounds.

Organic chemistry is the study of carbon and its compounds. Since there are 118 knownelements, it often seems odd that an entire branch of chemistry is devoted to a single element andits compounds while the other 117 elements and their compounds are all lumped together in aseparate discipline, but there is a very good reason for this. There are about 1.5 million knowninorganic compounds. This is a lot of compounds, and you will not be required to know them all forthis class, although it may seem like it before you're finished. The number of known organiccompounds varies, depending on the reference. These sources state that there are from 16 to 40million or more known organic compounds. Carbon is the basic element of life, and livingcreatures have developed an astonishing array of different organic compounds. A traditionaldefinition of organic compounds are those that contain carbon, i.e., compounds in which carbonis found in the molecular formula. But there are a small number of carbon-containing compoundsthat are classified as inorganic, such as the cyanides, carbonates, and bicarbonates (we’ll learnmore about these in Chapter 3). So a better working definition of organic compound is acompound that contains both carbon and hydrogen in its molecular formula, with the exception

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of the bicarbonates (hydrogen carbonates). Inorganic compounds are compounds that do notcontain both carbon and hydrogen. This is a general way of classifying organic compounds and thereare some notable exceptions in the “real world,” but it will be the way by which we make thedistinction between organic and inorganic compounds in this class. I’ll expand on this thought at the endof this chapter. It’s a little thing, but something I expect you to learn.

Biochemistry is the study of the chemistry of life and living things. This is a field of greatimportance to those who practice in the health sciences. The ways the atoms and moleculesbehave in living systems is predicated on the rules of organic chemistry.

Physical chemistry is the study of how the laws of physics affect things as small as atomsand molecules. During the early decades of the 20th Century researchers were astonished todiscover that very small particles, like atoms and molecules, do not obey the laws of physicsexpounded by Sir Isaac Newton. This spurred the development of the field known as quantummechanics, which explains the behavior of things that are very small.

Why is chemistry relevant?

We live in a world that grows more and more technical by the day. Even if you're notinterested in medicine or how the body works, how do you know if the claims made by someoneselling a particular health care product or beauty product are legitimate or bogus? Are youwilling to trust an industry that releases substances into the environment when it says they aresafe, or will you believe those who might be willing to distort or misrepresent the truth in orderto give weight to their argument that the chemicals released are dangerous? Do you trust peopleor government agencies that assure you that exposure to a particular chemical is not dangerous,particularly when their arguments may be based on economic or political considerations and noton pure science? When medicine is prescribed for a sick family member or friend, are you surethey will receive the appropriate medication and not something that may aggravate their illnessor even kill them?

Whether you're interested in science or not, whether you're interested in chemistry or not,you need to know some of the fundamentals of the science in order to be a good consumer, agood citizen, and a good family member and friend. You need to know a little bit aboutchemistry in order to protect yourself. You need to know a little bit about chemistry in order tobe able to think clearly and correctly about many important everyday issues.

To many people chemistry is a dirty word and chemicals are bad and undesirable, but thisattitude is terribly naive and uninformed. Everything around us is chemicals, from the air webreath, the water we drink, and the food we eat, to our very bodies. All living things are made upentirely of chemicals. All dead things, and all things that never have or never will live (such asrocks) are made up of chemicals. The earth on which we live is nothing more, at its mostfundamental, than a great mass of chemicals. The sun on which we rely for light and heat andlife itself is a giant globe of reacting chemicals. And so on. It is true, the incorrect use of

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chemistry is responsible for many important health and environmental issues. But chemistry alsomakes the world a better place.

Compare your life with that of someone living a year ago, ten years ago, fifty years ago,or a hundred or more years ago. What things do you have that they did not? How are medicineand health care different for you? personal hygiene? transportation? communication? the wayyou spend your leisure time? food and nutrition? the clothes you wear? Almost inevitably,anything you have that someone in the past did not have is attributable, either directly orindirectly, to advances made in chemistry.

What do chemists do to make the world a better place? In order to address this questionlet me tell you a little bit about the work of some of my friends who hold masters and doctoraldegrees in chemistry. They are involved in the study of:

• The analysis and identification of DNA and large biological molecules.• The chemistry of blood clotting mechanisms.• The chemistry of cardiovascular disease.• The chemical mechanisms of infectious diseases.• New chemical methods for breast cancer detection.• The development of new and improved antibiotics and other medicinal compounds.• The development of new and improved fertilizers and pesticides.• Food chemistry.• Personal hygiene products & cosmetics.• The chemistry of surfaces, which sounds terribly arcane until one realizes that many

important reactions will only occur on certain types of surfaces.• Nanotechnology• Composite materials, such as graphite fiber products.• The production of gasoline and other petroleum-based products.• Lubricants and fuel additives for automobiles.• The detection of drugs and poisons.This is just a small sample of the sorts of things that chemists do to make our world a betterplace.

What is matter?

As is the case with chemistry, there are a number of different definitions of matter. It iscommon to define matter as anything that has both mass and volume. I prefer to define matter asanything made up of atoms, which are the basic building blocks of the world around us. It is truethat there are types of matter smaller than atoms, such as protons, neutrons, electrons, andquarks, but we leave the study of particles smaller than atoms to courses in physics.

The study of matter is the study of chemistry. This means that the study of matter is thestudy of things made up of atoms and molecules. We will learn more about this in Chapter 2.

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How is matter classified?

Matter can be classified either by its state or by its composition.

There are three common states of matter, solids, liquids, and gases. We do not concernourselves with more exotic states of matter such as plasmas or Bose-Einstein condensates whenwe study elementary chemistry.

We will use the word "particle" as we discuss the three common states of matter, sincesubstances can be made up either of atoms, molecules, or ions. Ions are atoms or molecules with anelectric charge. We’ll discuss ions in the very near future. Particle is used to generically represent thevarious small units of matter we may find in a substance.

There are various ways to make distinctions between the three common states of matter.For our purposes, the most useful comparison at the particle level is as follows:

In the solid state the particles are relatively very close together, usually well-ordered(tidily arranged), and are held together by relatively strong attractive interactions between thesolid particles.

In the liquid state the particles are not quite as close together, not quite as well-ordered,and the attractive interactions are not quite as strong.

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In the gaseous state the particles are relatively far apart, chaotically arranged (i.e., notordered), and attractive interactions between the particles are either negligible or nonexistent.

There are specific names for the transitions from one state of matter to another. Thesetransitions are also called phase changes, or changes of state. These refer to state changes in anymaterial, and not just in water. The names of these changes of state are listed in the graphic at thetop of page 4. You do need to know all six of them, melting, vaporization, sublimation, condensation,freezing, and deposition.

Matter can also be classified based on its composition.

An element is a substance in which all of the atoms are the same. In other words, anelement is atomically pure. Elements may exist as single atoms or in molecular (i.e. more thansingle atom) form, such as Fe, O2, P4 and S8. Notice that the molecular formulas of elementsreflect their composition from one and only one type of atom. A molecular formula is ashorthand way of representing the types of atoms - and how many of each - are found in asubstance. The molecular formula “Fe” tells us our substance is made up exclusively of ironatoms. When we discuss O2, we are talking about two oxygen atoms that have chemicallycombined (i.e, they two atoms have chemically bonded to each other) to form a molecule. Eventhough O2 molecules contain two atoms, this is still representative of an element since both ofthe atoms are of the same element. This is also true of phosphorus, P4, and sulfur, S8.

Compounds are substances made from the atoms of two or more elements that are heldtogether by chemical bonds. There are millions of compounds in the everyday world around us,including water (H2O), carbon dioxide (CO2), and glucose (C6H12O6). Each compound has its ownmolecular formula. What does the molecular formula of water tell you about it’s composition? Howabout the chemical compositions of carbon dioxide and glucose? Notice that the molecular formulasof compounds reflect that they are made up of two or more different types of atoms. Asmentioned, it is essential to understand that the different types of atoms in compounds are heldtogether by chemical bonds, a topic we will examine in Chapter 3. One important thing toremember about compounds is that the atoms in a compound can only be separated from eachother by chemical reactions.

A mixture is a substance composed of two or more components which can be separatedby physical means, i.e., based on the different physical properties of the components. Thosecomponents may be elements, compounds, or other mixtures. As an example, table salt (NaCl,sodium chloride) dissolved in water can be recovered by evaporating the water. The ability toseparate the components of a mixture by taking advantage of differences in physical propertiesmeans that chemical bonds do not exist between the components. We will refine this concept morecorrectly in Chapter 8 but for the time being, this will suffice for our purposes.

There are two types of mixtures. They are classified by how evenly (or uniformly)dispersed the particles are at the molecular level.

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In homogeneous mixtures there is uniformity in composition, properties, and appearancethroughout the mixture down to the particle level. Examples of homogeneous mixtures includetable salt dissolved in water, red wine, gasoline, and alloys such as 18 karat gold.

Heterogeneous mixtures are characterized by discontinuities in composition, properties,and appearance at levels well before the molecular level. In other words, heterogeneous mixturesdo not have uniform composition, properties, and appearance throughout the mixture. Examplesof heterogeneous mixtures include pizza, granite, a nice fizzy soda in a glass filled with ice, anda peanut butter and jelly sandwich. What do you see when you look at a PBJ sandwich? Does it lookthe same throughout, or do you see layers? What does the presence of layers mean about itscomposition?The separation of the components of a mixture can be accomplished using physical changes, i.e.,changes based on the physical properties of the substances. The separation of the components ofa compound can only be achieved by chemical reactions. Elements are pure and cannot beseparated into atoms of any type other those atoms that are unique to the element.

What kinds of properties does matter have?

Matter has physical properties, which are properties that can be observed and measuredwithout changing the composition of the material. The physical properties of a substance includeits color, odor, state of matter, melting point, boiling point, heats of vaporization and fusion,density, solubility, metallic character, electrical and thermal conductivity, magnetic properties,crystal shape, malleability, ductility, and viscosity. Most common physical changes eitherinvolve changes of state (phase transitions) or are the consequence of mechanical processing(e.g. grinding, crushing, slicing, pulverizing, gluing pieces together, etc.)

The chemical properties of a material are the chemical reactivity of the substance, that is,how a substance changes its composition when it chemically reacts with other substances.

There is a difference between changes of state (physical changes) and chemical reactions.Physical changes are those involving physical properties, or, changes in which there is no changeto the chemical composition of the material(s) under observation. When chemical changes (or,chemical reactions) occur, the starting materials are consumed and new materials are formed dueto the breaking and making of chemical bonds. Chemical reactions may be indicated by colorchanges, the absorbing or release of energy (heat, light, sound, electric), or by the formation ofnew materials such as gases, pure liquids, or solids (precipitates). These are general indicationsthat a chemical reaction has occurred and they are not absolutely correct in all instances. Use them tohelp you identify chemical reactions in the world around you, but use them with care!

As an example, The molecular formula of water is H2O. This tells us that each watermolecule consists of two hydrogen atoms and one oxygen atom. If we take water out of ourkitchen faucet, it is in the liquid state. If we take liquid water and cool it by placing it in thefreezer it will eventually change from the liquid to the solid state. But while its state has

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changed, the water remains H2O, even though we call solid water “ice.” If we take solid water(ice) and put it in a pan on the stove, as it warms it will first change from the solid to the liquidstate and then from the liquid to the gaseous state. However, as these changes of state occur, thewater remains chemically unchanged. It remains H2O.

There is another way to think of the distinction between physical and chemical changes.If a substance maintains its same molecular formula throughout a process, it is a physical change.On the other hand, if a process results in changes to the molecular formulas of participatingsubstances, a chemical reaction has taken place. When water freezes, the process is H2O(l) =>H2O(s). The (l) indicates that our substance is in the liquid state, while (s) indicates our substance isnow in the solid state. Even though a change of state has occurred, the molecular formula of waterremains unchanged, indicating this is a physical process.

On the other hand, if we take a glass of liquid water and place a nine volt battery in it,something interesting occurs. This is an experiment you can safely do yourself. Should you elect totry it, I’d suggest you use a clear drinking glass or some other clear glass container. It will notharm the glass. Add about one teaspoon of table salt to one to two pints of tap water, and placethe battery in the water. Streams of bubbles begin to emanate from the battery terminals almostimmediately. As you observe the formation of bubbles, you can see that a far greater volume ofbubbles is being produced at one terminal than at the other. This is because the electrical energyof the battery is being used to break the chemical bonds that hold the hydrogen atoms andoxygen atoms in water molecules together. Hydrogen gas is being produced at the terminalproducing the greater volume of gas, while oxygen gas is being produced at the other terminal.Why would you expect to see more hydrogen gas produced than oxygen gas? Hint: think about themolecular formula of water.

2 H2O(l) => 2 H2 (g) + O2 (g)

This chemical equation described the process. It indicates that while we start with water, thebubbles that are formed consist of hydrogen gas and oxygen gas. Some water has been consumedduring this process and converted into different substances. This is typical of chemical reactions.One or more substances are consumed as one or more new substances are produced.

An introduction to the elements

As we stated earlier, elements are substances in which all of the atoms are the same.There are 118 known elements, of which about 90 are naturally occurring. This is the first ofseveral controversial statements you’re going to encounter in this class. You should know there is somedisagreement from one source to the next as to the number of naturally-occurring versus syntheticelements. Don’t worry about it. This is an early introduction to a reality of chemistry: there are manyareas in which chemists disagree with one another. There are sound reasons for these disagreements. I

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expect you to know what I tell you in these lecture notes and to understand that as you look beyondthem, you will at times find slightly contradictory information, in the text and elsewhere. It is not thatthese other sources are right and I am wrong, or visa versa. It is that we respectfully disagree with oneanother. While these disagreements are common, they are usually also minor. We as chemists aretypically united in the larger, more important points that we teach our students in class.

The twenty most abundant elements in the earth's crust by mass are: Note: I do not expect you tomemorize this table.

element percent composition

oxygen 45.5

silicon 27.2

aluminum 8.3

iron 6.2

calcium 4.7

magnesium 2.8

sodium 2.3

potassium 1.8

titanium 0.6

hydrogen 0.2

phosphorus 0.1

manganese 0.1

fluorine 0.05

barium 0.04

strontium 0.04

sulfur 0.03

carbon 0.02

zirconium 0.02

vanadium 0.01

Chlorine 0.01

all others 0.05

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A table of the fifteen most abundant elements in the human body is found below. It isinteresting to note the difference in the composition of the earth’s crust and that of the humanbody. Again, I do not expect you to memorize this information. It comes from a more detailed tablefound at "Elemental Composition of the Human Body”http://web2.iadfw.net/uthman/elements_of_body.html should you find yourself interested.

element mass of element ina

70 kg person (kg)

percentcomposition

(mass percent)

oxygen 43 61.4

carbon 16 22.9

hydrogen 7.0 10.0

nitrogen 1.8 2.6

calcium 1.0 1.4

phosphorus 0.8 1.1

potassium 0.1 0.2

sulfur 0.1 0.2

sodium 0.1 0.1

chlorine 0.09 0.1

magnesium 0.02 0.02

iron 0.004 0.006

fluorine 0.003 0.004

zinc 0.002 0.003

silicon 0.001 0.001

An important source of information about the elements is found in the periodic table. Wewill we will discuss it in a moment.

The natural states of the elements

How do the elements occur naturally? In other words, when we venture out into the realworld, in what state or condition might we find a sample of a particular element? Knowing thenatural states of the elements is often very important in understanding chemical reactions. You

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should know the following by heart, and soon. Questions regarding this information frequently find theirway into quizzes and exams.

In their elemental state, most elements occur as large numbers of individual atoms of thesubstance, chemically bonded one to another. An element is represented using its chemicalsymbol, which we can find in the periodic table. So, as an example, if we talk about copper wemay represent it as “Cu,” which is its atomic symbol. If we work with potassium in a chemicalreaction we’ll represent it with it’s symbol “K.” And so on. We’re going to talk a great dealabout the periodic table in this class. It is an absolutely essential tool to anyone studyingchemistry. I do not expect you to memorize it, and will always provide you with a copy onquizzes and exams. You can find a copy of the periodic table I will give you on exams at the followingaddress: http://www.slcc-science.org/chem/giddings/chem1110int/info/periodic-table-2016-revised.pdf. Ivery strongly suggest you print a copy of it now and keep it in front of you at all times as you work yourway through this course. Bear in mind that while I do not expect you to memorize the table, you arerequired to know what the table means and represents, how to use it, and to understand a number ofthings about it that we will discuss as we move through class. As an important word of warning: youshould never attempt any quiz, Mastering Chemistry exercise, or exam without a copy of the periodictable immediately in front of you. To do so is an exercise in self-abuse, as well an invitation to trouble.However, you’re the one paying to take this course. You can heed my advise, or you can learn the hardway, through your own mistakes, as you prefer.

Seven elements occur naturally in diatomic (i.e., molecules made up of two atoms) form.These include hydrogen, nitrogen, and oxygen (H2, N2, and O2 respectively) and the halogensfluorine, chlorine, bromine, and iodine (F2, Cl2, Br2, and I2 respectively). You must remember thatthese substances occur as diatomic molecules or you will come to grief in this class repeatedly! Amnemonic device shared with me by a student was: “Have No Fear Of Ice Cold Beer.” In this phrase,the first letter of each word corresponds to an element that exists as diatomic molecules, (H)ave =hydrogen, (N)o = nitrogen, (F)ear = fluorine, (O)f = oxygen, (I)ce = iodine, (C)old = chlorine, and (B)eer= bromine. There are several others. Learn one and remember it.

Phosphorus and sulfur occur naturally as P4 and S8 molecules respectively. However, wewon’t worry much, if at all, about this in this class. In other words, if we deal with elementalphosphorus or sulfur in this class, we’ll simply represent them as P or S respectively unlesswe’re instructed otherwise.

In the real world it is generally more common to find elements chemically combined withatoms of other elements, rather than in their pure form. In other words, in the real world most ofthe elements occur as compounds with other elements. Finding hydrogen and oxygen atomsbonded together in water (H2O), sodium and chlorine bonded together in sodium chloride(NaCl), and aluminum and oxygen bonded together in an aluminum ore called bauxite (Al2O3)

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are three common examples. Note that compounds are represented using molecular formulas,short-hand devices used to indicate the types of atoms - and how many of each - are found in asingle molecule of the compound. H2O, NaCl, and Al2O3 are examples of molecular formulas.H2O tells us that a single water molecule consists of two hydrogen atoms and one oxygen atom.NaCl informs us the there is one sodium atom and one chlorine atom found in a single moleculeof sodium chloride. In a single molecule of bauxite, also known as aluminum oxide, there aretwo aluminum atoms and three oxygen atoms chemically bonded to each other. We’ll discusscompounds at greater length in the near future.

The Periodic Table

The Periodic Table is a table of the elements arranged in rows and columns in sequenceof increasing atomic number (i.e., the number of protons in the nucleus). The rows are calledperiods and the columns are called groups. Periods and groups are vocabulary words you needto know.

The elements are represented by symbols in the periodic table.

In most cases the relationship between the name and the symbol is obvious. For example,O-oxygen, H-hydrogen, N-nitrogen. For a few elements the symbol is derived from anon-English language like Latin or German, e.g., K-potassium (kalium), Fe-iron (ferrum),Au-gold (aurum), W-tungsten (wolfram, Swedish). In these cases there is not an obviouscorrelation between the element's name and its symbol.

I will always give you a periodic table on exams and quizzes. A copy of what I will provide on theexams is at link mentioned previously. If you haven’t already printed yourself a copy, now is the time todo so. As you look at this information, remembering this is what I will give you on exams, it should beapparent that I will not require you to memorize element names and symbols. However, I very strongly

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suggest that you know the names and symbols of the first twenty elements, along with the names andsymbols of all other elements commonly used in this class. It will help you immensely and make yourwork go more quickly if you have committed these names and symbols to memory. You do not need tomemorize other information such as atomic numbers and atomic weights. If pertinent they will beprovided.

The top number in the periodic table entry for each element is the atomic number, thenumber of protons an atom has in its nucleus. The bottom number in the periodic table for eachelement is the atomic weight, the weighted average of the atomic masses of all of the isotopes ofa particular element. The atomic mass is the sum of the protons and neutrons in an isotope or aparticular atom. The masses of individual atoms is expressed in amu, or atomic mass units and isequal to 1/12 the mass of a single 12C isotope (or, 1.661 x 10-27 kg; don’t memorize this number. Ifyou need to use it, I will give it to you). We will discuss atomic number, atomic mass, atomicweight, and isotopes in more detail in Chapter 2 (will you be able to wait, or will the suspense justkill you?).

In the periodic table, groups are "chemical families," because they exhibit similarchemical behavior. We’ll explain why this is so in chapter 2. Some of the groups have specialnames, which you must know:

Group 1A - alkali metals.Group 2A - alkali earth metals.Group 6A - chalcogens.Group 7A - halogens.Group 8A - Noble Gases.Groups 1B-8B: transition metals. This includes the ten series of groups thecontains the elements from scandium to zinc (4th period), yttrium to cadmium (5thperiod), lanthanum to mercury (6th period), and actinium to copernicum (element#112, 7th period).Groups 1A, 2A, 3A-8A: main block elements.

The "island" by itself beneath the main body of the periodic table consists of two periods,the lanthanides and actinides. Collectively these two periods are known as rare earth metals.Interestingly there are a number of rare earth metals far more common than some of the transitionmetals (the 1B-8B elements in the periodic table) such as gold, silver, platinum, and etc. Normally it isgroups that have special names. These two rows are something of an anomaly in this sense, that theyhave their own special names.

The periodic table can also be divided into three families of elements, metals, nonmetals,and metalloids.

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Metals are elements characterized by the tendency to give up electrons in reactions (i.e.form cations), are good thermal and electrical conductors, and are usually lustrous, malleable,and ductile. Metals are found to the left of the "staircase" line on the periodic table. Lustrousmeans shiny. Malleability is the ability to be pounded into a particular shape. Ductility is the ability tobe drawn into a fine wire. One other important note: it’s true that hydrogen is found to the left of thestaircase line. However, hydrogen is a nonmetal. In this sense, it’s position in the periodic table issomething of an anomaly. There are very good reasons for this anomalous position, but I will notexplain them. Take comfort in knowing the hydrogen is the only exception, i.e., the only nonmetal foundto the left of the staircase line.

Nonmetal elements are diametrically opposite to metallic elements in their primaryfeatures. They are characterized by the tendency to gain electrons in reactions (i.e. form anions)and are poor conductors of heat and electricity, dull, and brittle. Nonmetals are found to the rightof the "staircase" line on the periodic table.

Metalloids (or, semi-metals) are elements characterized by the some of the properties ofboth metals and nonmetals. Metalloids straddle the "staircase" line on the periodic table, i.e, theyare the elements found directly on either side of the staircase line. Having mentioned metalloidshere, you can promptly forget about them as they have no relevance to us as we study chemistryin this class. You will need to be able to classify each and every element as a metal or anonmetal, for reasons that will become apparent when we discuss chemical bonds in Chapter 3.In other words, unless I ask you very specifically about metalloids, for the purposes of this class you canforget they exist. This is probably the first, last, and only time we’ll use the word metalloid in class. Forour purposes in this class, every element is either a metal or a nonmetal. It’s that simple.

Examine the periodic table in the front inside cover of your textbook, or the one found atthe link I mention above that you’ll be using on the exams. Can you find the staircase line? Canyou tell which elements are metals? Which are nonmetals? This is an essential fundamental skillyou must develop.

Chemistry is an empirical science

Chemistry is an empirical science. This means it is based on observations andmeasurements made during experiments. This chapter will address measurements andcalculations in chemistry, although we do not actually make any measurements in this course.

Scientific notation and powers of 10

(Note: excluding the sections on “energy and heat” and “density,” nearly everything from this pointthrough the remainder of the chapter is a review of things you should have learned in math 0900,0950, and 0980/1010, the prerequisite math courses for this class. It is your responsibility to know

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these things. If you do not remember discussing them in math, it is your responsibility to do thenecessary review work to bring yourself up to speed! This includes reviewing information in the course'stext book, material in your old math books and notebooks, and online support material such as may befound at Kahn Academy online (https://www.khanacademy.org/). Math is the single biggest challenge tomost students in this class, often far more so than any chemical concepts we discuss. If your math skillsare weak or you are out of practice, you face a long, uphill, challenging climb in class. I’m not trying toscare you off. I do however want you to be fully aware of what awaits you this semester.

It is common to use either very large or very small numbers in science. The size of mostatoms is in the trillionths of an inch (0.000000000001 inch) range. The atoms in a few ounces ofa substance may number as many as one million billion billion(1,000,000,000,000,000,000,000,000) or more. How can we conveniently express numbers thatare either very large or very small?

Scientific notation is a scientific shorthand that makes the expression of large and smallnumbers easier. It is based on our use of a base ten number system. Each digit in a number is aplaceholder representing some multiple of ten. For example, a number greater than 1, such as4321, can also be written as 4321 = (4 x 1000) + ( 3 x 100) + (2 x 10) + (1 x 1), or, as (4 x 10 x10 x 10) + (3 x 10 x 10) + (2 x 10) + (1 x 1). A number less than 1, such as 0.5678, can also bewritten as 0.5678 = (5 x .1) + (6 x .01) + (7 x .001) + (8 x .0001).

4321 is written as 4.321 x 103 in scientific notation. Numbers written using scientificnotation consist of two pieces, the coefficient and the exponent. The coefficient is the part withthe decimal, in this case, 4.321. The exponent is the power to which 10 is raised, in this case, 3.In scientific notation the coefficient is always obtained by moving the decimal until is just oneplace to the right of the first non-zero digit. The magnitude of the exponent depends on howmany places the decimal is moved. The sign of the exponent depends on which direction thedecimal is moved. If the number is greater than one, the decimal will be moved to the left and theexponent will be a positive number. If the number is less than 1 (note: less than one is a smallnumber but not a negative number. Negative numbers are less than 0), the decimal will be movedto the right and the exponent will be negative. The number 0.5678 would be written as 5.678 x10-1.

A few examples:

54405 = 5.4405 x 104

0.000006036 = 6.036 x 10-6

-201,000 = -2.01 x 105

-0.023473 = -2.3473 x 10-2

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Being able to use scientific notation is an essential fundamental skill required in this class. I will oftenrequire you to report your answers on homework, quizzes, and exams using scientific notation. If youdo not remember how scientific notation works, you need to review and bring yourself up to speed.

Measurements and significant figures

First, some definitions. There are many words used in science that have slightly differentmeanings than in everyday use. To most people the words accuracy and precision aresynonymous. To a scientist they have similar but distinct and different meanings. Accuracy is anindication of how close to a correct value a measurement is. Precision is a function of thecloseness of the results of a series of measurements to each other in value.

To the layman this distinction may seem absurd. It is often wondered how the correctvalue of a measurement cannot be known. But in the real world, the absolute correctness of mostmeasurements is seldom known.

As an example, assume someone takes exactly 1.000 pound of finely powdered lead andscatters it evenly across a 10.00 acre field. Now assume a farmer plows the lead deep into theground. Is there still exactly 1.000 pound of lead, or is it possible that perhaps some of the leadstuck to the blades of the plow and the tires of the tractor? Now, assume that the field is inconstant use over a ten year period. Will the field still contain exactly 1.000 pound of lead? Whatprocesses might be responsible for a loss of lead in the 10 acres? The action of wind and water,the integration of lead into growing plants, and the transport of lead by either vehicles or animalsmoving through the field may all contribute to a loss of lead. Lead may also move into the field,especially if there is a copper smelter nearby, since lead is a common by-product of copper andsilver refining. The finely powdered lead that results from metal refining processes can becarried by the wind and deposited in the field, even if the refinery is many miles distant. Also,lead may have existed in the soil before we added our 1.000 pound. If we are asked how muchlead remains in the 10.00 acres, how accurate will our answer be if we state 1.000 pound?

How can we determine exactly how much lead remains in the field? If we want to beexact, we must remove all of the soil to a depth of several feet or more, over all of the ten acres,and measure the total amount of lead. Obviously, this is not a feasible approach.

Alternatively, we may take a number of soil samples scattered across the ten acres andthen extrapolate our test results as being representative for the entire field. This may result in aloss of accuracy, but it is more realistic in terms of time and money since testing can beexpensive. The results are treated statistically to help us estimate the accuracy of our testing. Intesting the samples we hope that the results of our tests are similar in value from one set ofsamples to the next. If the results are in fact similar, we say that the results are precise. Thegreater the difference between the results of various sets of samples, the more imprecise our dataare.

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One very important aspect of testing involves the measurement of standards. Standardsare samples in which the amount of the substance of interest is known exactly because it is addedunder very carefully controlled conditions in a laboratory. Standards are always analyzed alongwith samples from the real world. The results obtained from the analysis of similar standardsshould be very nearly the same, i.e., they should be precise. The greater the precision in themeasurement of a series of standards, the greater the likelihood of the measurements beingaccurate under normal conditions. In other words, precise measurements are also usuallyaccurate measurements, although this might not be entirely true for a number of reasons we willnot discuss here.

Realistically, can we ever know exactly how much lead remains in the field? No. Butwith careful testing we can provide a highly accurate estimation, one that is very close to theactual, correct vaue.

Whenever measurements are made in the real world, there are limitations in themeasurements. For example, if the pound of lead was initially measured on an electronicscientific balance before it was scattered through the field, it might be confidently said that itweighed exactly 1.000 pound. But what if the lead was initially weighed on a bathroom scale,which only indicates in one pound increments? Or worse, what if the lead was weighed on a feedstore scale which only registers 10 pound increments? How certain might we be that the poundof lead weighed exactly 1.000 pound?

Scientists are required to know the limitations of the equipment they use when makingmeasurements during experiments. These limitations are reflected in the manner in whichscientists report their observations. Whenever a scientist reports a measurement, he writes thenumber to as many places as he is certain are correct plus a last place, in which there isuncertainty due to the limitations of the instrument used to make the measurements.

If we define the distance one mile by measuring it with a yardstick, we might report that1 mile is equal to 5280 feet. However, if we measure that same distance using lasers and GPS(global positioning satellite) technology, we might find that 1 mile is equal to 5280.000 feet,since the distance can be measured more exactly with the more sophisticated equipment.

The number of significant figures in a measurement is equal to all of the places in ameasurement we are certain about plus the first uncertain place. In this class we do not makemeasurements. We must learn to recognize how many significant figures are in the numbers reportedto us. Nonzero digits are always significant. Zeros may or may not be significant, as describedbelow.

There are three types of zeros used in numbers.

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Leading zeros are found before (to the left of) the first non-zero digit in a number arenever significant. As an example, the zeros in 0.053 are leading zeros. This value has twosignificant figures.

Captive zeros are found between two non-zero digits and are always significant. Giventhe number 100.01, we have three captive zeros and five total significant figures.

Trailing zeros are found after (to the right of) the last non-zero digit. They may or maynot be significant, depending on the presence of a explicit decimal in a number. The zeros in1300 are trailing zeros but are not counted as significant figures because a decimal is not shownexplicitly in the value. But given 1300. with it’s explicit decimal shown, the two trailing zerosbecome significant and the overall value has four significant figures.

Here are some examples. How many significant figures are found in each value below?

10101 5 sf significant figures, or, sig figs as many lovingly call them; captivezeros are always significant

73,000 2 sf trailing zeros without a decimal are not significant73,000. 5 sf trailing zeros with a decimal are significant0.0006149 4 sf leading zeros are never significant0.00456000 6 sf leading zeros are never significant and trailing zeros with a

decimal are significant 0.002022300 7 sf leading zeros are never significant, captive zeros are always

significant, and trailing zeros with a decimal are significant4321 = 4 sf.5678 = 4 sf0.5678 = 4 sf why isn’t the zero counted as a sig fig?0.56078 5 sf why is the second zero counted as a sig fig when the first one isn’t?93,567,891 = 8 sf.123456789 = 9 sf-23,456 = 5 sf note that the sign of a number does not affect the way its sig figs arecounted-.456 = 3 sf-0.456 = 3 sf why isn’t the zero counted as a sig fig?

Do not confuse the relevance of placeholder zeros with significant figures. Leading ortrailing zeros are often essential in the reporting of a number, but they may or may not besignificant (i.e., they may or may not be counted as sig figs), depending on the type of zero theyare.

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Numbers may need to be rounded if they are to be reported to the correct number of sigfigs. Never round any numbers until all calculations have been completed. The rules forrounding may be found in your text if you are not familiar with them.

SI units

The International System (or SI system) is the measurement system used in science. It isalso known as the metric system. Its basic units are mass in kilogram (kg), length in meter (m),time in seconds (s), temperature in Kelvins (K; note: not degrees Kelvin), and quantity in moles(mol).

Prefixes are used to indicate how many or what part of a base unit is being described. Theprefixes and their abbreviations are as follows. Note: these abbreviations are case-sensitive.

unit abbreviation powers of ten examples

yotta Y 1024

zetta Z 1021

exa E 1018

peta P 1015

tera T 1012 trillions

Every year vegetation around the world emits about 500teragrams of isoprene, a hydrocarbon that plays an importantrole in atmospheric chemistry.

giga G 109billions

A 32 gigabyte thumb drive can hold 32 billion bytes ofinformation.

mega M 106millions

A one megaton nuclear bomb delivers the same amount ofdestructive energy as the detonation as one million tons of TNT.

kilo k 103thousands

One kilogram is exactly equal to one thousand grams.

deci d 10-1tenth

In sound measurements one decibel is equal to one-tenth of abel, a unit of sound pressure.

centi c 10-2hundredths

A centimeter is exactly one one-hundredth of a meter.

milli m 10-3thousandths

A milligram is exactly one one-thousandth of a gram.

micro μ, u, or mc 10-6millionths

Fine human hair has a diameter of about 50 micrometers, 50one-millionths of a meter.

nano n 10-9billionths

A nanosecond is exactly one one-billionth of a second.

pico p 10-12trillionths

The diameters of a single copper atom is about 157 picometers,or 157 trillionths of a meter.

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femto f 10-15thousand-trillionth

There are some very fast chemical reactions that occur on afemtosecond time scale, in thousand-trillionths of seconds.

atto a 10-18million-trillionth

The concentration of certain chemicals in individuals nerve cellshas been measured in the attomolar range.

zepto z 10-21

yocto y 10-24

Work at femto/atto scale is often at the cutting edge of chemistry today. You must know the baseunits and the prefixes from tera to pico in this class. You do not need to know from yotta to peta, orfrom femto to yocto.

Measurements and measured quantities

In chemistry length is commonly reported in meters (m), centimeters (cm), andmillimeters (mm). The sizes of individual atoms and molecules are reported in the picometer(pm) and nanometer (nm) range. A unit called the Angstrom (Å) is occasionally encountered andis equal to 0.1 nm (1 Å = 1 x 10-10m).

Mass is reported in kilograms, although the masses of individual atoms may beexceedingly small. Note that there is a difference between mass and weight. Mass is the amountof matter present in a substance and does not change with respect to local gravitational fields.Mass is constant. Weight depends on gravity. As an example, if a person weighs 180 pounds onearth and travels to the moon, the mass of the person will remain constant. How much of theperson is there will not change. However, since the gravity of the moon is about one-sixth that ofearth, the person's weight would change from 180 pounds to 30 pounds. This is why a trip to themoon might be appealing to some who perceive themselves as overweight.

Temperature is reported using the Fahrenheit scale in the United States but the centigrade(or Celsius) and Kelvin scales are used in science. The Fahrenheit scale is based on the boilingpoint and the melting point of water at 212°F and 32°F respectively. The centigrade scale is alsobased on the melting and boiling points of water, but these are stipulated as occurring at 100°Cand 0°C respectively. Note that 1°C = 1.8°F, i.e., the Celsius degree is larger by nearly a factorof 2. The Kelvin scale is based on 0 K at absolute zero, the point at which all molecular motionstops. On this scale we find the melting point of water at 273 K and the boiling point of water at373 K. Note that 1K = 1°C, i.e., they are the same size. Also note that we refer to degreesFahrenheit and degrees Celsius but never to degrees Kelvin. The word degree is implied in thename Kelvin. So, we state that the melting point of water occurs at 32 degrees Fahrenheit, 0degrees Celsius, and at 273 Kelvin.

The equations used to convert between the temperature scales is as follows:

F to C: (F - 32)/1.8

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C to F: (C x 1.8) + 32

C to K: C + 273

K to C: K - 273

K to F: convert K to C and then convert to F

F to K: convert F to C and then convert to K

Derived quantities-volume

Two of the most important derived quantities are volume and density. Derived quantitiesare based on the measurement of more than one attribute of a material or object.

Volume is the product of the length, width, and height of an object ( for a cube or box-shaped object, V = l x w x h). Since the standard unit of length is the meter, the standard unit ofvolume could be stated in cubic meters (m3), but this is too large a unit to be practical unless thevolume of large objects, such as houses, mountains, or planets, is being discussed. In chemistryvolume is commonly expressed in units of cubic centimeters (cm3). Since 1 m = 100 cm, then 1m3 = (100 cm)3 = 1 x 106 cm3. Another common unit of volume in chemistry is the liter (L),which is equivalent to one-one thousandth of a cubic meter, or, to one deciliter: 0.001 m3 = 1 dm3= 1 L. There are one thousand milliliters in one liter (1000 mL = 1 L), and a milliliter and a cubiccentimeter are exactly the same size (1 mL = 1 cm3 ).

Derived quantities-density

The density (D) of any material is the ratio of the mass (m) of the material to the volume(V) occupied by that mass (D = m/V). Density is an important physical property of substances. Itgives information about a substance at both a macroscopic and a microscopic level. Givensamples of two substances, the more dense material either has more particles per unit volume,heavier particles, or both.

Density is determined by the careful measurement of a substance's mass and volume.Mass can be determined on a scale or balance. Volume can be determined in a variety of waysbut is commonly determined by displacement. To measure the volume of an object bydisplacement, water is placed in a device in which its volume can be carefully measured. Onesuch common device is a thin glass cylinder with volume increments marked on its side called agraduated cylinder (see, for example, Figure 2.3 in your text). After carefully noting the volumeoccupied by just the water, an object may be added. It will be noticed that the level of the waterin the cylinder moves upward, as water is displaced (moved out of the way) by the object. Thedifference between the starting volume and the final volume of water is equal to the volume ofthe object.

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The following are examples of how density may be used in chemical calculations.

Throughout these lecture notes I will work many problems as examples. An essential part of your studyin this class should consist of working these same example problems on your own, along with similarexamples in the text and also the recommended end-of-chapter problems. Like it or not, chemistry isvery problem based. If you don’t practice, practice, practice then you’re going to have some real troublewith the exams. But don’t worry. When you repeat the course because you failed to practice, maybeyou’ll have learned how important this really is to your success in this course.

• If 10.0 g of liquid occupies a volume of 13.5 mL, what is the density of the liquid?

density = m/V = 10.0g/13.5 mL = 0.741 g/mL

• Pure gold (24 K) has a density of 19.32 g/cm3 . A wedding band weighs 3.50 grams. Howmuch water must it displace (i.e., what must its volume be?) if the band is in fact 24 Kgold?

if D = m/V then V = m/D3.50 g/19.32 g/cm3 = 0.181 cm3

• Benzene has a density of 0.880 g/mL. If you need exactly 78.12 grams of benzene, whatvolume would you measure out?

if D = m/V then V = m/D 78.12 g / 0.880 g/mL = 88.77 mL .88.8 mL

Energy and heat

In science, energy is defined as the capacity to do useful work. We speak of potentialenergy, which is energy that is stored and based on position, and kinetic energy which is theenergy of moving objects. Energy can be transferred as heat, light, sound (acoustic), orelectricity by various chemical reactions. Reactions that give off energy are said to beexothermic, while reactions that absorb energy to make them happen are said to be endothermic.We’ll discuss exothermic and endothermic reactions in more detail in Chapter 7.

A few examples of exothermic reactions include fire, explosions, and the flow ofelectricity from a battery. Examples of endothermic reactions include cooking and recharging abattery.

Energy is often used to increase the temperature of a substance. The common unit ofenergy is the calorie (cal). One calorie is the amount of heat required to raise the temperature of

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1 g of water 1°C or 1 K since these temperature units are the same relative size, they are used moreor less interchangeably in this discussion. The SI unit of energy is the Joule (J). There is arelationship between calories and Joules: 4.184 J = 1 cal. We should note that the calories usedto calculate the nutrition value of food are actually equal to one thousand of the calories wediscuss in this chapter. In other words, in nutrition kilocalories are the basis for nutritionaldiscussions.

Let’s do a thought experiment. Imagine taking pieces of gold and iron that are equal inmass and immersing them in boiling water. After allowing them to reach the temperature of thewater, they are removed and dried. The gold will return to room temperature much quicker thanthe iron. This is because of an important physical property called specific heat. The specific heatof a substance is the amount of heat required to raise the temperature of one gram of thesubstance by 1°C. Comparatively speaking things with high specific heats heat more slowly andcool more slowly than things with lower specific heats.

Time for another thought experiment, although this can be attempted at home. Analuminum pan at room temperature is filled with a mass of water equal to the mass of the pan.The pan of water is placed on a stove. Which will heat more quickly, the pan or the water? Onceremoved from the burner, the water is poured out of the pan into another container at roomtemperature. Which will cool more rapidly, the empty pan or the water in its new container?

The pan will heat more rapidly than the water while on the burner, not just because it isin contact with burner, but also because aluminum has a lower specific heat than water. This isconfirmed by the second part of the experiment. The empty pan will return to room temperaturemuch quicker than the water. This is because aluminum in the pan has a lower specific heat thanwater.

Dimensional analysis and problem solving

It is often necessary to convert from one set of units to another while working scienceproblems. This is made possible by a problem-solving technique called dimensional analysis orthe factor label method.

At the risk of belaboring the point I just made above, you absolutely must learn how to use thistechnique. There are times I may require you to answer problems using dimensional analysis and willnot accept an answer, even if it is correct, obtained using other means. Chemistry is difficult, notbecause of the math, which is usually the sort of math most of us learned in junior high school and highschool (although we may not remember it). Chemistry is difficult because the problems are nearlyalways story problems (oh damn, I hear you say) and can only be mastered through practice, practice,and more practice. If you do not practice what follows below you will be extremely unhappy at examtime. Please, do not hesitate to get help from the STEM Center in the Science bulding on the RedwoodRaod campus, other tutors, or your instructor when you have trouble. You have been warned!

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Dimensional analysis works by using conversion factors, which make the correct conversionfrom a value with one set of units to a value with a different set of units possible.

In mathematics there is something called the multiplicative identity, which states that anynumber multiplied by 1 will result in that number i.e., 1 x y = y. A conversion factor permits theconversion from one set of units to another without changing the true value of the numberassociated with the units, because a conversion factor is always exactly equal to one. Conversionfactors are created through equivalent ways of expressing things. As examples:

• 1 mile = 5280 feet If we divide both sides by 5280 feet we obtain (1 mile/5280 feet) = (5280 feet/5280 feet) = 1 The right side of the equation cancels, since the numerator and denominator are equal,and this leaves us with (1 mile/5280 feet) = 1

• 1 foot = 12 inchesIf we divide both sides by 12 inches we obtain (1 foot/ 12 inches) = (12 inches/12 inches) =1The right side of the equation cancels, since the numerator and denominator are equal,and this leaves us with (1 foot/12 inches) = 1

• 1 hour = 3600 secondsIf we divide both sides by 3600 seconds we obtain (1 hour/3600 seconds) = (3600 seconds/3600 seconds) = 1The right side of the equation cancels, since the numerator and denominator are equal,and this leaves us with (1 hour/3600 seconds) = 1

• 1 m = 1000 mm If we divide both sides by 1000 mm we obtain (1 m/1000 mm) = (1000 mm/1000 mm) = 1The right side of the equation cancels, since the numerator and denominator are equal,and this leaves us with (1 m/1000 mm) = 1

Dimensional analysis is the process of problem solving using conversion factors. Howare conversion factors used? While if you know another approach to this technique you’rewelcome to use it, I suggest the following approach, using as an example the conversion of 150lbs to kg:

1. Identify what is given and how should the problem wind up. In this case we are given150 pounds, and expect to wind up with a certain number of kg.

2. Identify which conversion factors are needed. In this example we need to know therelationship between pound and kilograms, which is 1 kg = 2.2 lbs or (1 kg/2.2 lbs) = 1

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3. Starting with what is given, we multiply it by our conversion factor. After using theconversion factor do the units cancel, or is another conversion factor needed?

150 lbs x (1 kg/2.2 lbs) = kg; the units are correct, no additional conversion factors areneeded

4. Do the arithmetic.

150 lbs x (1 kg/2.2 lbs) = 68.2 kg = 68 kg

Since there are only 2 significant figures in 150 lbs (why is the “0" not counted as a sigfig in this number?), we would report our answer as 68 kg.

Note: I strongly suggest that you always check to make sure that your units cancel each otherout and that you wind up with the correct units before you attempt any of the arithmetic in theproblem. I will never require you to do a dimensional analysis problem any one specific way, butI know from experience if you take care of the units in a problem first, then the arithmetic partnearly always turns out correct. If you rush through problems you will ultimately prove to beyour own worst enemy at exam time.

What does this mean physically? That there are 68 kg in 150 pounds of anything.

Some additional examples of dimensional analysis problems are as follows:

• Convert 62.5 cm to inches

conversion factors needed: 1 inch = 2.54 cm

62.5 cm x (1 inch/2.54 cm) = 24.6 inches

• Convert 3.62 ounces to mg

conversion factors needed: 1 lb = 16 oz; 1 lb = 453.6 g; 1 g = 1000 mg

(3.62 oz) x (1 lb/16 oz) x (453.6 g/1 lb) x (1000 mg/1 g) = 102,627 mg = 1.03 x 105 mg

Note that the solution of this problem required the use of several conversion factors. We can use asmany as we need to solve a problem. Also note that when we perform multiplication or division, thenumber of significant figures we report in our answer is determined by the value with the fewest sig figsin our calculations. Conversion factors are usually considered to be exact, and, as such, have an infinitenumber of significant figures. This means that, in this class, when we work problems, we report ouranswers with the correct number of sig figs, and the correct number of sig figs is determined by the

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information we are given in the problem, not the conversion factors. You should also be comforted toread that I will always provide you with all non-metric conversion factors in this class. The metricprefixes are very much your responsibility, and I require you to know them.

• Convert 125 m/s to mph (miles per hour)

conversion factors needed: 60 sec = 1 minute; 60 min = 1 hr; 1 m = 100 cm; 1 in = 2.54 cm; 1 ft = 12 in; 1 mile = 5280 ft

(125 m/s) x (60 s/1 min) x (60 min/1 hr) x (100 cm/m) x (1 in/2.54 cm) x (1 ft/12 in) x (1 mile/5280 ft) = 279.6 mile/hr = 280. mph

• Convert the area of an 8.5" by 11" piece of paper to square centimeters (cm2)

conversion factors needed: 1 in = 2.54 cm

step 1. 8.5 in x 11 in = 93.5 in2

step 2. (93.5 in2) x (2.54 cm/1 in)2 = 603.2 cm2 = 6.0 x 102 cm2

Note that we do not have a conversion factor that permits the conversion from square inches to squarecentimeters. We can, however, create our own, simply by squaring both sides of the conversion factorgiven i.e., if 1.00 in = 2.54 cm then (1.00 in)2 = (2.54 cm)2. If you need to do this you must rememberto square the number as well as the unit during the calculation. You should also note that there is morethan one way to solve this problem and come up with the correct answer. You should also note thatwhile normally we like to do dimensional analysis problems using one continuous string of conversionfactors, sometimes it is easier to break a problem down into pieces as we did here.

• Convert 6 yds3 of cement to cm3

conversion factors needed: 1 yd = 3 ft; 1 ft = 12 in; 1 in = 2.54 cm

(6 yds3) x (3 ft/1 yd)3 x (12 in/1 ft)3 x (2.54 cm/1 in)3 = 4,587,329.1 cm3 = 5 x 106 cm3

Note again the use of conversion factors created, this time, by cubing both sides of a relationship, e.g.(3 ft)3 = (1 yd)3

• Hexane is an organic liquid used in the manufacture of gasoline. If a railroad car contains2.5 x 104 gallons of hexane, how much does the liquid weigh in pounds? The density ofhexane is 0.6594 g/mL

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Ok, this problem is a bit of a tough one but we can do it. You must know that the density is needed toact as a bridging conversion factor between the conversion of mass to volume, or, visa versa, theconversion of volume to mass. It is not possible to correctly convert from mass to volume, or fromvolume to mass, without using the materials density as a conversion factor.

conversion factors needed: 1 gal = 4 qt; 1 L = 1.057 qt; 1000 mL = 1 L; 1 lb = 453.6 g

step 1. (2.5 x 104 gallons) x (4 qt/1 gal) x (1 L/1.057 qt) x (1000 mL/1 L) =94,607,379.376 mL

step 2. (94,607,379.376 mL) x (0.6594 g/mL) x (1 lb/453.6 g) = 137,531.1 lbs =1.4 x 105 lbs

Organic or inorganic - a postscript

I expect you to be able to differentiate between organic and inorganic compounds. You are goingto see questions on quizzes and exams in which you’ll be asked to identify a compound asorganic or inorganic, based strictly on a compound’s molecular formula. It is as simple as this:

If a compound's molecular formula includes both carbon (C) and hydrogen (H), it'sorganic. There are a very small group of exceptions to this rule. In this class the only exceptionsare the bicarbonates (bicarbonates are also known as hydrogen carbonates; either the word"bicarbonate" or the words "hydrogen carbonate" will be part of the name.) Organic compoundsmay also include other types of atoms in addition to carbon and hydrogen atoms. The presence ofatoms other than C and H does not affect the compound’s designation as organic. Organiccompounds may include metal atoms, as well as other types of nonmetal atoms, in addition tocarbon and hydrogen. It is common to find the molecular formulas of organic compounds beginwith carbon, followed by hydrogen, then followed by one or more other elements in a moleculeof the compound which are listed in alphabetical order in its molecular formula. As an example,the molecular formula of sodium gluconate is C6H11NaO7. However, there are times and reasonssome chemists may list its molecular formula as NaC6H11O7. Either way, since it’s molecularformula contains both carbon and hydrogen, it is an organic compound.

If a compound's molecular formula does not contain both carbon and hydrogen (C andH), it's inorganic - end of story.

As a few examples:

The molecular formula of water is H2O. This molecular formula contains H but not C.This is an inorganic compound.

The molecular formula of carbon monoxide is CO. This molecular formula contains Cbut not H. This is an inorganic compound.

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The molecular formula of dinitrogen tetroxide is N2O4. This molecular formula does notcontain either C or H. This is an inorganic compound.

The molecular formula of methane (natural gas) is CH4. This molecular formula containsboth C and H. This is an organic compound.

The molecular formula of strychnine is C21H22N2O2. This molecular formula containsboth C and H. This is an organic compound. The presence of atoms other than C and H does notaffect the designation of strychnine as an organic compound.

The molecular formula of sodium bicarbonate is NaHCO3. While this molecular formulacontains both C and H, the inclusion of the word "bicarbonate" in the name tells us this is anexception to the rule and is an inorganic compound.

The molecular formula of lead(II) acetate is Pb(C2H3O2)2.This molecular formulacontains both C and H. The presence of a metal atom, lead (Pb), does not change the status ofthis substance as an organic compound. A common misconception is that organic compoundscannot contain metal atoms. This is incorrect. As long as a compound contains both carbon andhydrogen in its molecular formula, it is an organic compound (except for bicarbonates!).

It is essential to remember there are 11 elements with symbols that begin with the letter "C."These include chlorine (Cl), cesium (Cs), chromium (Cr), copper (Cu), and etc. None of theseelements are carbon (C), and so their presence in a molecular formula does not make thecompound organic. As examples:

HCl - inorganic (no carbon)

CsH - inorganic (no carbon)

CuH - inorganic (no carbon)

Please, be sure you use a periodic table and a list of the elements and their symbols whenyou're asked to determine if a compound is organic or inorganic. It's easy to inadvertently make amistake if you're not careful.

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Chapter 2: Atoms and the Periodic TableThe structure of the atom

Atoms are made of three types of subatomic particles which are summarized in the tablebelow.

charge relative mass absolute mass (kg)

proton +1 1 1.6726 x 10-27

neutron 0 1 1.6749 x 10-27

electron -1 1/1800 9.109 x 10-31

The Nucleus

Protons and neutrons are held together in the nucleus at the center of the atom. Undernormal circumstances particles with similar electrical charges repel each other, but protons areheld together in the nucleus despite this repulsion by the Strong Force, which is about 100 timesstronger than the electrostatic repulsion experienced by the similarly charged particles. TheStrong Force, while not a particularly clever name, is the strongest physical force in the knownuniverse. Neutrons also help to mitigate the electrical repulsion that arises when positivelycharged protons are jammed tightly together in an atomic nucleus. Electrons orbit the nucleusrelatively far away. As we stated above, the atom is mostly empty space.

To provide an illustration of this statement, the diameter of the nucleus of a hydrogenatom is reported as 1.75 fm (source: Atomic nucleus. (2016, July 12). In Wikipedia, The FreeEncyclopedia. Retrieved 16:53, August 12, 2016, fromhttps://en.wikipedia.org/w/index.php?title=Atomic_nucleus&oldid=729430288). However, thehydrogen atom, the combined size of the hydrogen atom’s nucleus and it’s single electron, havean approximate diameter of 0.106 nm, or 1.06 x 105 fm. In other words, the diameter of the atomis nearly one hundred thousand times larger than that of just the nucleus. To put this anotherway, if the nucleus of a hydrogen atom was represented by a ball exactly one meter in diameterin downtown Salt Lake City, the outer boundary of the hydrogen atom would be found inBrigham City (to the north, or Payson, UT, to the south), as the approximate distance fromdowntown Salt Lake City to Brigham City in the north, and to Payson in the south, is roughly100 km.

The notion that an atom is mostly just empty space becomes even more evident when we think ofatomic sizes in terms of volumes, rather than diameters. The nucleus of a hydrogen atom with adiameter of 1.75 fm occupies a volume of approximately 2.81 fm3. If a hydrogen atom has adiameter of 0.106 nm, the atom’s volume is 6.24 x 1014 fm3. In plain English, the overall volumeof the atom is 100 trillion times greater than the volume of the atom’s nucleus alone. What is it

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that occupies all of this space between the nucleus and the electron? Absolutely nothing, as far aswe know.

Hydrogen atoms are the smallest of any element. How much larger do the nuclei of atomsbecome? The nucleus of a typical uranium atom contains 92 protons and 146 neutrons. In spiteof this large number of protons and neutrons, the nucleus of a uranium atom is about 15 fm indiameter. In other words, even in atoms much larger than hydrogen, the atoms’ nuclei remainextremely small.

An electrostatic attraction, or, electromagnetic attraction, occurs between the positivelycharged protons in the nucleus and the negatively charged electrons as they move around it.While electrons are much smaller than protons, it is the number of electrons possessed by anatom, or, more correctly, the way that the electrons are arranged around the nucleus, thatdetermines that atom's chemical reactivity. This is one of the essential topics of chapter 2 andindeed, of this entire course.

Atoms and protons

It is the number of protons in an atomic nucleus that determines the identity of the atom.The number of protons in the nucleus is unique for the atoms of each element. The number ofprotons in an atom is invariable in chemical reactions. If the number of protons did change in achemical reaction, the atom would be converted into an atom of a different element. These typesof events do occur in nature but they are nuclear events, not chemical events. The process ofconverting an atom of one substance into an atom of another substance is (through the gain orloss of proton) is called transmutation and will be discussed in Chapter 11. To reiterate:transmutation never occurs during chemical reactions.

As we just said, and let us reiterate for emphasis: in chemical reactions, the number of protons inreacting atoms never changes. In nuclear reactions, which are not chemical reactions, the number ofprotons in the nuclei of reacting atoms may or may not change, through transmutation.

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If we examine a periodic table we find three pieces of information for every element.

The top number is the atomic number. In the above figure we see that this number, 29, and is thenumber of protons in the nucleus of all copper atoms. The atomic number is sometimes referredto as the"Z" number of an element although it is a bit archaic to do so. The atomic number of eachelement tells us the number of protons found in the nuclei of atoms of that element.

Atoms, neutrons, and isotopes

While the number of protons for a given element is invariable, the number of neutrons ina nucleus can differ. In other words, it is possible for atoms to have the same numbers of protonsbut different numbers of neutrons. We call these isotopes, atoms with the same atomic number(same number of protons) but with different numbers of neutrons. All elements have at least oneisotope. Most elements have two or more isotopes.

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The sum of the protons and neutrons in any isotope is known as its atomic mass. If wewant or need to be precise we can also include the masses of the electrons when calculating atomicmass but usually we ignore them because the are so small compared to the masses of the protons a