INTRODUCTORY CHEMISTRY ONLINE! 
Paul R. Young
University of Illinois at Chicago
University of Illinois at Chicago Introductory Chemistry Online!
Paul R. Young
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This text was compiled on 09/09/2023
TABLE OF CONTENTS
Licensing
1: Measurements and Atomic Structure
1.1: Why Study Chemistry
1.2: Organization of the Elements - The Periodic Table
1.3: Scientific Notation
1.4: SI and Metric Units
1.5: Unit Conversion with the Metric System
1.6: Significant Figures
1.7: Atomic Structure and Electron Configuration
1.8: Filling Orbitals with Electrons
1.S: Measurements and Atomic Structure (Summary)
2: The Physical and Chemical Properties of Matter 2.1: Pure Substances and Mixtures
2.2: The States of Matter
2.3: Density, Proportion and Dimensional Analysis
2.4: Chemical and Physical Properties and Changes
2.5: Conservation of Mass
2.S: The Physical and Chemical Properties of Matter (Summary)
3: Chemical Bonding and Nomenclature
3.1: Compounds, Lewis Diagrams and Ionic Bonds
3.2: Covalent Bonding
3.3: Lewis Representation of Ionic Compounds
3.4: Identifying Molecular and Ionic Compounds
3.5: Polyatomic Ions
3.6: Resonance
3.7: Electronegativity and the Polar Covalent Bond
3.8: Exceptions to the Octet Rule
3.9: Common Valence States and Ionic Compounds
3.10: Nomenclature of Ionic Compounds
3.11: Nomenclature of Molecular Compounds
3.S: Chemical Bonding and Nomenclature (Summary)
4: The Mole and Measurements in Chemistry 4.1: Measurement and Scale - The Mole Concept
4.2: Molar Mass
4.3: Mole-Mass Conversions
4.4: Percentage Composition
4.5: Empirical and Molecular Formulas
4.S: The Mole and Measurements in Chemistry (Summary)
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5: Chemical Reactions
5.1: Chemical Changes and Chemical Reactions
5.2: Chemical Equations
5.3: Balancing Chemical Equations
5.4: Classifying Chemical Reactions
5.5: Oxidation and Reduction Reactions
5.6: Predicting Products from Chemical Reactions
5.7: Predicting Solubility Trends
5.8: The Energetics of Chemical Reactions
5.S: Chemical Reactions (Summary)
6: Quantitative Relationships in Chemistry 6.1: An Introduction to Stoichiometry
6.2: Molar Stoichiometry in Chemical Equations
6.3: Mass Calculations
6.4: Percentage Yield
6.5: Limiting Reactants
6.S: Quantitative Relationships in Chemistry (Summary)
7: Aqueous Solutions
7.1: Hydrogen Bonding and the Properties of Water
7.2: Molecular Dipoles
7.3: Dissolution of Ionic Compounds
7.4: Concentration and Molarity
7.5: Solution Stoichiometry
7.6: Dilution of Concentrated Solutions
7.S: Aqueous Solutions (Summary)
8: Acids, Bases and pH
8.1: Hydrogen Bonding
8.2: Ionization of Acids in Solution
8.3: Conjugate Acid-Base Pairs
8.4: Acids-Bases Reactions: Neutralization
8.5: The Meaning of Neutrality - The Autoprotolysis of Water 8.6: pH Calculations
8.7: Titrations - Neutralization and Stoichiometry
8.S: Acids, Bases and pH (Summary)
9: The Gaseous State
9.1: Gasses and Atmospheric Pressure
9.5: The Ideal Gas Law
9.6: Combining Stoichiometry and the Ideal Gas Laws
9.S: The Gaseous State (Summary)
9.2: The Pressure-Volume Relationship: Boyle’s Law
9.3: The Temperature-Volume Relationship: Charles’s Law 9.4: The Mole-Volume Relationship - Avogadro’s Law
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10: Principles of Chemical Equilibrium 10.1: The Concept of Equilibrium Reactions
10.2: The Equilibrium Constant
10.3: Calculating Equilibrium Values
10.4: Using Molarity in Equilibrium Calculations
10.5: Equilibria involving Acids and Bases
10.6: The pH of Weak Acid Solutions
10.7: Solubility Equilibria
10.8: Study Points
11: Nuclear Chemistry
11.1: Radioactivity
11.2: The Nuclear Equation
11.3: Beta Particle Emission
11.4: Positron Emission
11.5: Radioactive Half-Life
11.6: Nuclear Fission
11.7: Nuclear Fusion
11.S: Nuclear Chemistry (Summary)
Index
Detailed Licensing
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Licensing
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CHAPTER OVERVIEW
1: Measurements and Atomic Structure
Chemistry is the study of matter and the ways in which different forms of matter combine with each other. You study chemistry because it helps you to understand the world around you. Everything you touch or taste or smell is a chemical, and the interactions of these chemicals with each other define our universe. Chemistry forms the fundamental basis for biology and medicine. From the structure of proteins and nucleic acids, to the design, synthesis and manufacture of drugs, chemistry allows you an insight into how things work. Chapter One in this text will introduce you to matter, atoms and their structure. You will learn the basics of scientific measurement and you will gain an appreciation of the scale of chemistry; from the tiniest atom to the incredibly large numbers dealt with in the “mole concept” (Chapter 4). Chapter One lays the foundation on which we will build our understanding.
1.1: Why Study Chemistry
1.2: Organization of the Elements - The Periodic Table
1.3: Scientific Notation
1.4: SI and Metric Units
1.5: Unit Conversion with the Metric System
1.6: Significant Figures
1.7: Atomic Structure and Electron Configuration
1.8: Filling Orbitals with Electrons
1.S: Measurements and Atomic Structure (Summary)
Thumbnail: Two small test tubes held in spring clamps. (CC BY-SA 3.0 Unported; Amitchell125 via Wikipedia)
This page titled 1: Measurements and Atomic Structure is shared under a CC BY-SA 4.0 license and was authored, remixed, and/or curated by Paul R. Young (ChemistryOnline.com) via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request.
1
1.1: Why Study Chemistry
Chemistry is the branch of science dealing with the structure, composition, properties, and the reactive characteristics of matter. Matter is anything that has mass and occupies space. Thus, chemistry is the study of literally everything around us – the liquids that we drink, the gasses we breathe, the composition of everything from the plastic case on your phone to the earth beneath your feet. Moreover, chemistry is the study of the transformation of matter. Crude oil is transformed into more useful petroleum products such as gasoline and kerosene by the process of refining. Some of these products are further transformed into plastics. Crude metal ores are transformed into metals, that can then be fashioned into everything from foil to automobiles. Potential drugs are identified from natural sources, isolated and then prepared in the laboratory. Their structures are systematically modified to produce the pharmaceuticals that have led to vast advances in modern medicine. Chemistry is at the center of all of these processes and chemists are the people that study the nature of matter and learn to design, predict and control these chemical transformations. Within the branches of chemistry you will find several apparent subdivisions. Inorganic chemistry, historically, focused on minerals and metals found in the earth, while organic chemistry dealt with carbon-containing compounds that were first identified in living things. Biochemistry is an outgrowth of the application of organic chemistry to biology and relates to the chemical basis for living things. In the later chapters of this text we will explore organic and biochemistry in a bit more detail and you will notice examples of organic compounds scattered throughout the text. Today, the lines between the various fields have blurred significantly and a contemporary chemist is expected to have a broad background in all of these areas.
In this chapter we will discuss some of the properties of matter, how chemists measure those properties and we will introduce some of the vocabulary that is used throughout chemistry and the other physical sciences.
Let’s begin with matter. Matter is defined as any substance that has mass. It’s important to distinguish here between weight and mass. Weight is the result of the pull of gravity on an object. On the Moon, an object will weigh less than the same object on Earth because the pull of gravity is less on the Moon. The mass of an object, however, is an inherent property of that object and does not change, regardless of location, gravitational pull, or whatever. It is a property that is solely dependent on the quantity of matter within the object.
Contemporary theories suggests that matter is composed of atoms. Atoms themselves are constructed from neutrons, protons and electrons, along with an ever-increasing array of other subatomic particles. We will focus on the neutron, a particle having no charge, the proton, which carries a positive charge, and the electron, which has a negative charge. Atoms are incredibly small. To give you an idea of the size of an atom, a single copper penny contains approximately 28,000,000,000,000,000,000,000 atoms (that’s 28 sextillion). Because atoms and subatomic particles are so small, their mass is not readily measured using pounds, ounces, grams or any other scale that we would use on larger objects. Instead, the mass of atoms and subatomic particles is measured using atomic mass units (abbreviated amu). The atomic mass unit is based on a scale that relates the mass of different types of atoms to each other (using the most common form of the element carbon as a standard). The amu scale gives us a convenient means to describe the masses of individual atoms and to do quantitative measurements concerning atoms and their reactions. Within an atom, the neutron and proton both have a mass of one amu; the electron has a much smaller mass (about 0.0005 amu).
Figure 1.2: Atoms are incredible small. Atoms are incredibly small. To give you an idea of the size of an atom, a single copper penny contains approximately 28,000,000,000,000,000,000,000 atoms (that’s 28 sextillion).
Atomic theory places the neutron and the proton in the center of the atom in the nucleus. In an atom, the nucleus is very small, very dense, carries a positive charge (from the protons) and contains virtually all of the mass of the atom. Electrons are placed in a diffuse cloud surrounding the nucleus. The electron cloud carries a net negative charge (from the charge on the electrons) and in a neutral atom there are always as many electrons in this cloud as there are protons in the nucleus (the positive charges in the nucleus are balanced by the negative charges of the electrons, making the atom neutral).
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An atom is characterized by the number of neutrons, protons and electrons that it possesses. Today, we recognize at least 116 different types of atoms, each type having a different number of protons in its nucleus. These different types of atoms are called elements. The neutral element hydrogen (the lightest element) will always have one proton in its nucleus and one electron in the cloud surrounding the nucleus. The element helium will always have two protons in its nucleus. It is the number of protons in the nucleus of an atom that defines the identity of an element. Elements can, however, have differing numbers of neutrons in their nucleus. For example, stable helium nuclei exists that contain one, or two neutrons (but they all have two protons). These different types of helium atoms have different masses (3 or 4 amu) and they are called isotopes. For any given isotope, the sum of the numbers of protons and neutrons in the nucleus is called the mass number. All elements exist as a collection of isotopes, and the mass of an element that we use in chemistry, the atomic mass, is the average of the masses of these isotopes. For helium, there is approximately one isotope of Helium-3 for every million isotopes of Helium-4, hence the average atomic mass is very close to 4 (4.002602).
As different elements were discovered and named, abbreviations of their names were developed to allow for a convenient chemical shorthand. The abbreviation for an element is called its chemical symbol. A chemical symbol consists of one or two letters, and the relationship between the symbol and the name of the element is generally apparent. Thus helium has the chemical symbol He, nitrogen is N, and lithium is Li. Sometimes the symbol is less apparent but is decipherable; magnesium is Mg, strontium is Sr, and manganese is Mn. Symbols for elements that have been known since ancient times, however, are often based on Latin or Greek names and appear somewhat obscure from their modern English names. For example, copper is Cu (from cuprum), silver is Ag (from argentum), gold is Au (from aurum), and iron is (Fe from ferrum). Throughout your study of chemistry, you will routinely use chemical symbols and it is important that you begin the process of learning the names and chemical symbols for the common elements. By the time you complete General Chemistry, you will find that you are adept at naming and identifying virtually all of the 116 known elements. Table 1.1 contains a starter list of common elements that you should begin learning now!
Table 1.1: Names and Chemical Symbols for Common Elements
Element Chemical Symbol Element Chemical Symbol
H Phosphorus P
Helium He Sulfur S Lithium Li Chlorine Cl Beryllium Be Argon Ar
Boron B Potassium K Carbon C Calcium Ca Nitrogen N Iron Fe Oxygen O Copper Cu Fluorine F Zinc Zn Neon Ne Bromine Br
Sodium Na Silver Ag Magnesium Mg Iodine I Aluminum Al Gold Au Silicon Si Lead Pb
The chemical symbol for an element is often combined with information regarding the number of protons and neutrons in a particular isotope of that atom to give the atomic symbol. To write an atomic symbol, you begin with the chemical symbol, then write the atomic number for the element (the number of protons in the nucleus) as a subscript, preceding the chemical symbol. Directly above this, as a superscript, now write the mass number for the isotope, that is, the total number of protons and neutrons in the nucleus. Thus, for helium, the atomic number is 2 and there are two neutrons in the nucleus for the most common isotope making the atomic symbol . In the definition of the atomic mass unit, the “most common isotope of carbon”, , is defined as having a mass of exactly 12 amu and the atomic masses of the remaining elements are based on their masses relative to this isotope. Chlorine (chemical symbol Cl) consists of two major isotopes, one with 18 neutrons (the most common, comprising 75.77% of
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natural chlorine atoms) and one with 20 neutrons (the remaining 24.23%). The atomic number of chlorine is 17 (it has 17 protons in its nucleus), therefore the chemical symbols for the two isotopes are and .
When data is available regarding the natural abundance of various isotopes of an element, it is simple to calculate the average atomic mass. In the example above, was the most common isotope with an abundance of 75.77% and had an abundance of the remaining 24.23%. To calculate the average mass, first convert the percentages into fractions; that is, simply divide them by 100. Now, chlorine-35 represents a fraction of natural chlorine of 0.7577 and has a mass of 35 (the mass number). Multiplying these, we get (0.7577 × 35) = 26.51. To this, we need to add the fraction representing chlorine-37, or (0.2423 × 37) = 8.965; adding, (26.51 + 8.965) = 35.48, which is the weighted average atomic mass for chlorine. Whenever we do mass calculations involving elements or compounds (combinations of elements), we always need to use average atomic masses.
This page titled 1.1: Why Study Chemistry is shared under a CC BY-SA 4.0 license and was authored, remixed, and/or curated by Paul R. Young (ChemistryOnline.com) via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request.
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