Chapter 3: Atoms Lecture Notes Review
Introduction: Mercury and Small-Scale Gold Mining
The Problem: Small-scale gold mining in gold-rich regions (e.g., South America and East Africa) leads to toxic mercury vapor contamination of soil, air, and water, causing severe health problems for nearby residents.
Adam Kiefer's Work: A chemistry professor from Mercer University, Adam Kiefer, leads a team measuring mercury levels, teaching safer mining techniques, and cleaning up highly polluted regions. His work emphasizes that miners pollute out of necessity to feed families, not disregard for the environment.
The Gold Mining Process (using mercury):
Miners collect coarse gold-containing ore from community mines.
Ore is milled into fine particles at a grinding facility.
Particles are placed in bowls with water and stirred; lighter soil rises and is drained, leaving heavier ore.
Elemental mercury is stirred into the remaining ore.
Gold particles dissolve in mercury, forming small, shiny pellets called an amalgam (a mixture of gold and mercury).
To separate gold, the solid amalgam is heated on a smoldering log.
Heat evaporates the mercury, leaving behind purified gold pellets. The mercury vapor is toxic and released into the environment.
Mercury Toxicity:
Highly toxic, affecting the central nervous system.
Causes symptoms like tremors, anxiety, memory loss, and eventually insanity.
Small-scale gold mining is the largest human activity releasing mercury into the environment.
Similarities and Differences between Gold and Mercury:
Similarities: Both are metals, shiny, melt at fairly low temperatures, and conduct electricity.
Differences: Gold is a solid at room temperature, while mercury is a liquid. Gold itself does not pose health hazards like mercury.
Chapter Focus: This chapter will explore atomic structure to understand why elements behave the way they do.
3.1 Atoms: The Essential Building Blocks
Matter: Everything we can see, touch, taste, hear, or smell is composed of matter.
Elements: All matter is composed of a basic set of elements. Currently, 118 elements are known, with about 98 naturally occurring.
Atom: The fundamental unit of matter. Each element is composed of a different type of atom, which gives it unique properties. Atoms combine to form compounds and mixtures with new behaviors.
Uncovering the Atom: From Democritus to Dalton
Democritus (~400 B.C.E.): A Greek philosopher who proposed the idea of atomos (Greek for "indivisible"). He argued that matter could be divided repeatedly until an indivisible particle was reached.
Antoine Lavoisier (late 1700s): Systematically explored chemical behavior. He formulated the Law of Conservation of Mass: "In chemical reactions, matter is neither created nor destroyed." This means the total mass before a reaction equals the total mass after.
Example: 4 grams of hydrogen react with 32 grams of oxygen to form 36 grams of water.
Example 3.1: If 48.6 grams of magnesium combines with 32.0 grams of oxygen, then 80.6 grams (48.6 ext{ g} + 32.0 ext{ g}) of magnesium oxide will form.
Try It 1: If acetylene ( 26 kg) burns with oxygen (X kg) to produce carbon dioxide ( 88 kg) and water ( 18 kg), then:
26 ext{ kg (acetylene)} + X ext{ kg (oxygen)} = 88 ext{ kg (CO}2) + 18 ext{ kg (H}2 ext{O})
26 ext{ kg} + X ext{ kg} = 106 ext{ kg}
X = 106 ext{ kg} - 26 ext{ kg} = 80 ext{ kg of oxygen}
John Dalton (1808): An English schoolteacher who published a framework for modern atomic theory. Key concepts included:
Elements are made of tiny, indivisible particles called atoms. (This part was later found to be incorrect as atoms have subatomic particles).
The atoms of each element are unique.
Atoms can join together in whole-number ratios to form compounds.
Atoms are unchanged in chemical reactions (not created or destroyed).
Example: When charcoal (carbon) burns, one carbon atom combines with two oxygen atoms to form one carbon dioxide molecule (1 ext{ C} + 2 ext{ O}
ightarrow 1 ext{ CO}_2). The atoms themselves are not altered, but the compound's properties are different.
Three Fundamental Ideas from this History:
All matter is composed of atoms.
The atoms of each element have unique characteristics and properties.
In chemical reactions, atoms are not changed but combine in whole-number ratios to form compounds.
Can We See Atoms?
Historically, atoms were too small to be seen directly with the naked eye or conventional microscopes.
Scanning Tunneling Microscopy (early 1990s): Developed by IBM, this technique allows scientists to visualize atoms on a metal surface and even manipulate them (e.g., forming the letters IBM).
X-ray Crystallography: A common technique where a solid is bombarded with X-rays. The patterns formed as X-rays pass through reveal the arrangement of atoms in the solid (e.g., the structure of aspirin). This helps in studying diseases, developing medicines, and designing new materials.
3.2 The Periodic Table of the Elements
Dmitri Mendeleev (late 1860s): A Russian scientist who organized the then-known 60 elements into a table based on their atomic masses and similar properties. He successfully predicted the existence of undiscovered elements like germanium and gallium.
Modern Periodic Table: Organizes all 118 known elements based on their properties.
"Periodic" meaning: Means "cyclical," referring to its organizational principle, similar to a calendar.
Organization:
Periods: Horizontal rows (1-7) across the table. Moving from left to right, elements exhibit a full spectrum of chemical properties.
Groups (or Families): Vertical columns. Elements within the same group generally share similar chemical and physical properties.
Chemical Symbols: One- or two-letter abbreviations (e.g., H for hydrogen, He for helium). The first letter is capitalized; the second (if present) is lowercase. Some symbols are derived from Latin names (e.g., Fe for iron from ferrum, Au for gold from aurum, Hg for mercury from hydrargyrum).
Caution: Capitalization is crucial (e.g., Co is cobalt, CO is carbon monoxide).
Regions of the Periodic Table
Main-group elements: Found in columns 1-2 and 13-18 (or 1 ext{A}-8 ext{A} in the older system). Many of their properties can be predicted from their location.
Transition elements (Transition Metals): Located in the middle columns (3-12) of the table. They tend to be harder and less reactive than main-group metals (e.g., iron, copper, gold, silver).
Inner transition elements: Two additional rows usually placed at the bottom of the table for readability. They actually belong in periods 6 and 7.
Lanthanide series: The first row; contains heavier, naturally occurring rare earth metals.
Actinide series: The second row; elements up to uranium are naturally occurring; most beyond uranium are human-made.
Metals, Nonmetals, and Metalloids
Metals: Located on the left side of the periodic table.
Typically solid at room temperature (mercury is an exception).
Malleable (can be molded into different shapes).
Good conductors of heat and electricity.
Nonmetals: Located on the upper right side of the periodic table.
Exhibit a wide variety of physical properties (solids like carbon/sulfur, liquid like bromine, gases like nitrogen/oxygen/fluorine).
Form a rich diversity of compounds (from simple gases to biomolecules and plastics).
Provide the atomic basis for life (e.g., carbon, nitrogen, oxygen, hydrogen).
Metalloids: Found along the "stairstep" pattern between metals and nonmetals.
Semiconductors: Conduct electricity, but not as efficiently as metals.
Essential components of modern electronics (e.g., silicon).
Groups (Families) of Elements
Alkali metals (Group 1A): Far left column (1 ext{A}).
Very soft metals.
React violently with air or moisture (e.g., lithium, sodium, potassium).
Alkaline earth metals (Group 2A): Column 2 ext{A}.
Reactive, but less violently than alkali metals.
React slowly with water but burn brightly with oxygen (e.g., beryllium, magnesium, calcium).
Halogens (Group 7A): Column 17 (7 ext{A}) on the upper right side.
Reactive nonmetals.
Exist as diatomic molecules ( ext{F}2, ext{Cl}2, etc.) in their elemental form.
Form many different compounds, found in substances like bleach, Teflon®, fire retardants, antiseptics, and table salt (e.g., fluorine, chlorine, bromine, iodine).
Noble gases (Group 8A): Rightmost column (18 or 8 ext{A}).
Very stable.
Generally do not react with other elements to form compounds.
All are gases at room temperature (e.g., helium, neon, argon).
3.3 Uncovering Atomic Structure
Dalton's atomic theory, while revolutionary, was incomplete. Atoms are not indivisible.
Subatomic particles: Atoms are composed of even smaller components.
Atomic Mass Unit (u or amu): A unit of mass used for atoms and subatomic particles.
1 ext{ u} = 1.66 imes 10^{-27} ext{ kg}
A hydrogen atom (the lightest) has a mass of approximately 1.0 ext{ u}.
Charge: A characteristic property of subatomic particles.
Particles can have a positive charge, a negative charge, or no charge.
Opposite charges attract; like charges repel.
Electricity (electrical energy): Involves the motion of charged particles.
The Discovery of Charged Particles
Alessandro Volta (1800): Built the first battery (electrochemical cell). This device generated an electrical current, the flow of charged particles. The battery's invention was crucial for discovering many elements and understanding charged particle interactions.
J.J. Thomson (1897): An English scientist who discovered the electron—a tiny, negatively charged subatomic particle, approximately 2,000 times smaller than the lightest atom. He showed that electricity involves the flow of electrons.
Plum Pudding Model: Thomson hypothesized that small, negative electrons were spread throughout a positive atomic substance, similar to blueberries in a muffin. This model attempted to balance the negative electron charge with a diffuse positive charge.
The Discovery of the Nucleus
Ernest Rutherford (1909): Thomson's former student, conducted the gold foil experiment, which disproved the plum pudding model.
Experiment: Rutherford and his students fired positively charged alpha particles at a thin gold film, expecting them to pass straight through.
Observation: Most alpha particles passed through as expected, but a small number were deflected back toward the source.
Conclusion: The atom is mostly empty space with a very tiny, incredibly dense, positively charged nucleus at its center. Alpha particles that hit the nucleus were deflected.
Modern Atomic Structure (Rutherford's findings elaborated):
Atoms are mostly empty space.
Nucleus: Contains nearly all the atom's mass, but occupies a tiny volume (like an insect in a football stadium).
Protons: Positively charged (charge +1), mass of about 1 ext{ u}. Located in the nucleus.
Neutrons: No charge (neutral), mass of about 1 ext{ u}. Located in the nucleus.
Electron cloud: The space around the nucleus where tiny, negatively charged electrons (charge -1) reside. Electrons are much lighter than protons or neutrons (about 1/2000 of their mass) but account for nearly the entire volume of the atom.
Neutral Atoms: For an atom to be electrically neutral, the number of protons and electrons must be equal.
3.4 Describing Atoms: Identity and Mass
Atomic Number (Z): The number of protons in an atom. This number determines the identity of an atom. For a neutral atom, the atomic number also equals the number of electrons. It is the integer value typically found above the atomic symbol on the periodic table.
Mass Number (A): The sum of the number of protons and the number of neutrons in an atom. It is always an integer and is typically not shown on the periodic table.
Isotopes: Atoms of the same element (i.e., same atomic number/number of protons) but with different mass numbers (i.e., different numbers of neutrons).
Example: Hydrogen has three isotopes:
Protium: 1 proton, 0 neutrons (mass number 1).
Deuterium: 1 proton, 1 neutron (mass number 2).
Tritium: 1 proton, 2 neutrons (mass number 3).
Symbolic Representation of Isotopes:
Mass number is placed at the upper left of the chemical symbol.
Atomic number is placed at the lower left of the chemical symbol.
Example: Chlorine's two common isotopes are {17}^{35}{ ext{Cl}} (meaning 17 protons, 35-17=18 neutrons) and {17}^{37}{ ext{Cl}} (meaning 17 protons, 37-17=20 neutrons).
Often, only the mass number and symbol are written (e.g., ^{235}{ ext{U}} or U-235).
Example 3.4 Determining Nuclear Structure: For an isotope of cadmium (^{ ext{114}}{ ext{Cd}}).
Cadmium (Cd) has an atomic number of 48 (from the periodic table).
Number of protons = 48.
Number of neutrons = Mass number - Atomic number = 114 - 48 = 66.
For a neutral atom (^{ ext{114}}{ ext{Cd}}), the number of electrons = 48.
Average Atomic Mass
Average Atomic Mass: The weighted average of the masses of an element's naturally occurring isotopes. This is the decimal value typically displayed beneath the atomic symbol on the periodic table.
It is a weighted average because it accounts for the relative abundance of each isotope.
Calculation: Multiply the mass of each isotope by its fractional abundance (percentage expressed as a decimal), and sum the results.
Formula: Average Atomic Mass ( ext{u}) = \sum ( ext{Isotope Mass} imes ext{Fractional Abundance})
Example (Carbon): 98.93 ext{%} is ^{12}{ ext{C}} (mass 12.000 ext{ u}); 1.07 ext{%} is ^{13}{ ext{C}} (mass 13.0034 ext{ u}).
ext{Average mass} = (0.9893 imes 12.000 ext{ u}) + (0.0107 imes 13.0034 ext{ u})
ewline = 11.8716 ext{ u} + 0.1391 ext{ u} = 12.0107 ext{ u}
Difference between Mass Number and Average Atomic Mass: Mass number refers to a specific isotope and is an integer. Average atomic mass refers to the element as it appears in nature, accounting for all isotopes and their abundances, and is usually a decimal.
3.5 Electrons - A Preview
The arrangement of electrons around the nucleus largely determines how atoms combine to form compounds.
The Bohr Model and the Quantum Model
Bohr Model (early 20th century): Developed by Ernest Rutherford and Niels Bohr.
Treated the atom like a tiny solar system, with the nucleus at the center and electrons orbiting in fixed paths, similar to planets orbiting the Sun.
A significant advance over the plum pudding model, explaining some elemental properties.
However, it was incomplete in describing electron behavior.
Quantum Model (1920s): The modern description of electrons.
Recognizes that electrons cannot be simply described as particles; their behavior also requires describing them as waves moving at different energies.
Electrons occupy an electron cloud around the nucleus.
Explains a wide variety of physical and chemical properties and has led to many technological advances (e.g., solar panels, supercomputers).
Summary of Atomic Structure Models:
Dalton's atomic theory (1808): Atoms are indivisible particles.
Plum pudding model (1904): Atoms are solid, with negative electrons spread throughout a positively charged matrix.
Bohr model (1913): Electrons orbit the nucleus like planets orbit the Sun.
Quantum model (1920s): Electrons behave as both particles and waves, occupying an electron cloud around the nucleus.
The Formation of Ions
Electrons, unlike protons and neutrons, occupy the outer volume of the atom, making them more easily gained or lost.
Ions: Atoms or groups of atoms that possess an overall net positive or negative charge due to an unequal number of protons and electrons.
Gaining electrons: Results in a negative charge (anion).
Example: Fluorine atom (9 protons, 9 electrons, charge 0) gains 1 electron to become a fluoride ion (9 protons, 10 electrons, charge -1).
Losing electrons: Results in a positive charge (cation).
Example: Lithium atom (3 protons, 3 electrons, charge 0) loses 1 electron to become a lithium ion (3 protons, 2 electrons, charge +1).
Example 3.6 Finding Ion Charge: A calcium ion has 20 protons and 18 electrons. Its charge is (+20) + (-18) = +2.
Example 3.7 Relating Ion Charge to Subatomic Particles: A sulfide ion has a charge of -2 and sulfur's atomic number is 16.
Neutral sulfur atom: 16 protons, 16 electrons.
Sulfide ion: Since the charge is -2, it has two extra electrons. So, 16 ext{ electrons} + 2 ext{ electrons} = 18 ext{ electrons}.
Capstone Question: Cinnabar (Mercury Sulfide)
In cinnabar, mercury forms an ion with a +2 charge ( ext{Hg}^{2+}), and sulfur forms an ion with a -2 charge ( ext{S}^{2-}$)$.
Atomic Number of Mercury (Hg): 80.
Table Completion for Mercury Isotopes ($_{ ext{80}}^{ ext{200}}{ ext{Hg}} ext{ and } _{ ext{80}}^{ ext{202}}{ ext{Hg}}$):
Isotope
Form
Protons
Neutrons
Electrons
^{ ext{200}}{ ext{Hg}}
Atom (charge =0)
80
120 (200-80)
80
^{ ext{200}}{ ext{Hg}}
Ion (charge =+2)
80
120
78 (80-2)
^{ ext{202}}{ ext{Hg}}
Atom (charge =0)
80
122 (202-80)
80
^{ ext{202}}{ ext{Hg}}
Ion (charge =+2)
80
122
78 (80-2)
Chemical Change: Heating cinnabar in the presence of oxygen gas to produce elemental mercury and sulfur oxides is a chemical change because new substances with different chemical compositions are formed.
Mass Calculation: If a manufacturer starts with 10.00 ext{ kg} of cinnabar and produces 8.62 ext{ kg} of elemental mercury:
Mass of sulfur in original cinnabar = 10.00 ext{ kg (cinnabar)} - 8.62 ext{ kg (mercury)} = 1.38 ext{ kg}.
Percentage of sulfur in cinnabar = (1.38 ext{ kg} / 10.00 ext{ kg}) imes 100 ext{%} = 13.8 ext{%} (by mass).
Chapter Summary - Mercury and Small-Scale Gold Mining Revisited
Historical Development of Atomic Theory:
Democritus first postulated atoms.
Antoine Lavoisier developed the Law of Conservation of Mass (matter is neither created nor destroyed).
John Dalton laid the foundation for modern atomic theory: matter is made of unique atoms that combine in whole-number ratios and are conserved in reactions.
Periodic Table:
Dmitri Mendeleev developed the periodic table, organizing elements by atomic mass and properties. It now organizes elements by atomic number and chemical reactivity.
Periods (rows) show a cycle of chemical behavior.
Groups/Families (columns) contain elements with similar properties.
Atomic Structure Discoveries:
Alessandro Volta's battery enabled element discoveries and the study of charged particles.
J.J. Thomson discovered the electron.
Ernest Rutherford's gold foil experiment revealed the atom's structure: mostly empty space with a tiny, dense, positively charged nucleus composed of protons (positive) and neutrons (neutral). Electrons (negative) form an electron cloud around the nucleus.
Describing Atoms:
Atomic number (Z): Number of protons; defines element identity and equals electrons in a neutral atom.
Isotopes: Atoms of the same element with different numbers of neutrons (thus different mass numbers).
Mass number (A): Sum of protons and neutrons in an isotope's nucleus.
Average atomic mass: Weighted average of isotope masses, found on the periodic table.
Electron Models:
Bohr model: Early model, electrons orbit nucleus like planets.
Quantum mechanical model: Modern model, treats electrons as both particles and waves, occupying an electron cloud.
Ions: Atoms or groups of atoms with an overall charge, formed by gaining or losing electrons.
Mercury and Small-Scale Gold Mining (Continued)
Unique Properties of Mercury: Only naturally occurring metal that is a liquid at room temperature; it is shiny, dense, and flows effortlessly.
Mercury's Role in Mining: Used to form an amalgam with gold, then evaporated to purify gold.
Health Hazards: Gaseous mercury is a powerful toxin. Inhaled mercury atoms enter the bloodstream and lose electrons to form toxic mercury ions ( ext{Hg}^{2+}$$), which damage kidneys and the central nervous system, leading to tremors, dementia, and death.
Dr. Kiefer's Interventions:
Conducts demonstrations using UV light and a fluorescent background to show invisible mercury vapors absorbing light and casting shadows, illustrating evaporation.
Uses atomic absorbance spectrometers linked to GPS to measure and map airborne mercury concentrations in communities.
Designed a device for miners to separate mercury from gold without releasing mercury into the environment, aiming to drastically reduce pollution.