Atoms, Molecules, and Ions - Detailed Notes

Atomic Theory of Matter

• Some Greek philosophers, like Democritus, proposed the existence of smallest particles called "atomos" (uncuttable).
• Experiments in the 18th and 19th centuries led to John Dalton's atomic theory in the early 1800s, based on:
  • The law of constant composition
  • The law of conservation of mass
  • The law of multiple proportions

Law of Constant Composition

• Compounds have a definite composition, meaning the relative number of atoms of each element in the compound is the same in any sample.
• Discovered by Joseph Proust.
• One of the laws on which Dalton’s atomic theory was based.

Law of Conservation of Mass

• The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.
• Discovered by Antoine Lavoisier.
• One of the laws on which Dalton’s atomic theory was based.

Law of Multiple Proportions

• If two elements, A and B, form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers.
• Discovered by John Dalton.
• When two or more compounds exist from the same elements, they cannot have the same relative number of atoms.

Postulates of Dalton’s Atomic Theory

• Each element is composed of extremely small particles called atoms.
• All atoms of a given element are identical in mass and other properties, but differ from atoms of other elements.
• Atoms are not changed into different atoms during chemical reactions; atoms are neither created nor destroyed during chemical reactions.
• Atoms combine to form compounds with a specific relative number and kind of atoms.

Discovery of Subatomic Particles

• Dalton viewed the atom as the smallest particle, but later discoveries showed that atoms are composed of smaller particles.
  • Electrons and cathode rays
  • Radioactivity
  • Nucleus, protons, and neutrons

The Electron (Cathode Rays)

• Streams of negatively charged particles emanate from cathode tubes, causing fluorescence.
• J.J. Thomson is credited with their discovery (1897).

The Electron

• Thomson measured the charge/mass ratio of the electron to be $1.76 \times 10^8$ coulombs/gram (C/g).

Millikan Oil-Drop Experiment (Electrons)

• Robert Millikan determined the charge on the electron in 1909.
• Knowing the charge/mass ratio, determining either the charge or the mass of an electron would yield the other.

Radioactivity

• Radioactivity is the spontaneous emission of high-energy radiation by an atom.
• First observed by Henri Becquerel.
• Marie and Pierre Curie further studied it.
• Its discovery showed that the atom comprises more subatomic particles and energy.

Radioactivity Types

• Three types of radiation were discovered by Ernest Rutherford:
  • α particles (positively charged)
  • β particles (negatively charged, like electrons)
  • γ rays (uncharged)

The Atom, circa 1900

• The “plum pudding” model, proposed by J.J. Thomson, was the prevailing theory.
• It described a positive sphere of matter with negative electrons embedded within it.

Discovery of the Nucleus

• Ernest Rutherford shot α particles at a thin sheet of gold foil and observed the scattering pattern of the particles.

The Nuclear Atom

• Since some particles were deflected at large angles, Thomson’s model was incorrect, leading to the nuclear view of the atom.

The Nuclear Atom

• Rutherford proposed a very small, dense, positive center (nucleus) with electrons around the outside.
• Most of the atom is space.
• Atoms are very small; 1–5 Å or 100–500 pm.
• Other subatomic particles (protons and neutrons in the nucleus) were discovered.

Subatomic Particles

• Protons (+1) and electrons (–1) have a charge; neutrons are neutral.
• Protons and neutrons have essentially the same mass (relative mass 1). The mass of an electron is so small we ignore it (relative mass 0).
• Protons and neutrons are found in the nucleus; electrons travel around the nucleus.

Atomic Number

• Atomic Number: the number of protons in the nucleus of an atom.
• Since atoms have no overall charge, the number of protons equals the number of electrons in an atom.

Atoms of an Element

• Elements are represented by one or two letter symbols, with the first letter always capitalized (e.g., C for carbon).
• All atoms of the same element have the same number of protons, known as the atomic number.
• The mass number is the total number of protons and neutrons in the nucleus of an atom.

Isotopes

• Isotopes are atoms of the same element with different masses.
• Isotopes have different numbers of neutrons, but the same number of protons.

Atomic Mass Unit (u)

• Atoms have extremely small masses. Heaviest known atoms have a mass of approximately $4 \times 10^{-22}$ g.
• A mass scale on the atomic level uses the atomic mass unit (u) as its base unit: $1 u = 1.66054 \times 10^{-24}$ g

Atomic Weight

• Using average masses in calculations is essential because we deal with large amounts of atoms and molecules.
• An average mass is found using all isotopes of an element, weighted by their relative abundances, resulting in the element’s atomic weight.
• Atomic Weight = $\sum$ [(isotope mass) × (fractional natural abundance)] for ALL isotopes.
• The masses of any atom is compared to C-12 (6 protons and 6 neutrons) being exactly 12.

Atomic Weight Measurement

• Atomic and molecular weights can be measured using a mass spectrometer.

Periodic Table

• The periodic table is a systematic organization of the elements.
• Elements are arranged in order of atomic number.

Reading the Periodic Table

• Boxes on the periodic table list the atomic number ABOVE the symbol.
• The atomic weight of an element is listed below the symbol on the periodic table.

Organization of the Periodic Table

• Rows are called periods.
• Columns are called groups.
• Elements in the same group have similar chemical properties.

Periodicity

• A repeating pattern of properties and reactivity is observed when examining the chemical properties of elements.

Groups

• Group 1: Alkali metals (Li, Na, K, Rb, Cs, Fr)
• Group 2: Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra)
• Group 16: Chalcogens (O, S, Se, Te, Po)
• Group 17: Halogens (F, Cl, Br, I, At)
• Group 18: Noble gases (He, Ne, Ar, Kr, Xe, Rn)

Periodic Table - Metals

• Metals are on the left side of the periodic table.
  • Shiny luster
  • Conduct heat and electricity
  • Solids (except mercury)

Periodic Table - Nonmetals

• Nonmetals are on the right side of the periodic table (including H).
• They can be solid (like carbon), liquid (like bromine), or gas (like neon) at room temperature.

Periodic Table - Metalloids

• Elements on the steplike line are metalloids (except Al, Po, and At).
• Their properties are sometimes like metals and sometimes like nonmetals.

Chemical Formulas

• The subscript to the right of the symbol of an element indicates the number of atoms of that element in one molecule of the compound.
• Molecular compounds are composed of molecules and almost always contain only nonmetals.

Diatomic Molecules

• The following seven elements occur naturally as diatomic molecules:
  • Hydrogen
  • Nitrogen
  • Oxygen
  • Fluorine
  • Chlorine
  • Bromine
  • Iodine

Types of Formulas

• Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.
• Molecular formulas give the exact number of atoms of each element in a compound.
• Knowing the molecular formula allows determination of the empirical formula, but the reverse is not true without more information.

Picturing Molecules

• Structural formulas show the order in which atoms are attached but do NOT depict the three-dimensional shape of molecules.
• Perspective drawings, ball-and-stick models, and space-filling models show the three-dimensional order of the atoms in a compound.

Ions

• When an atom or group of atoms loses or gains electrons, it becomes an ion.
• Cations are formed when at least one electron is lost (metals).
• Anions are formed when at least one electron is gained (nonmetals, except noble gases).

Common Cations

• Examples:
  • H⁺ (hydrogen ion)
  • Li⁺ (lithium ion)
  • Na⁺ (sodium ion)
  • K⁺ (potassium ion)
  • Ag⁺ (silver ion)
  • Mg²⁺ (magnesium ion)
  • Ca²⁺ (calcium ion)
  • Ba²⁺ (barium ion)
  • Al³⁺ (aluminum ion)
  • NH₄⁺ (ammonium ion)
  • Cu⁺/Cu²⁺ (copper(I) or cuprous/copper(II) or cupric ion)
  • Fe²⁺/Fe³⁺ (iron(II) or ferrous/iron(III) or ferric ion)

Common Anions

• Examples:
  • F⁻ (fluoride ion)
  • Cl⁻ (chloride ion)
  • Br⁻ (bromide ion)
  • I⁻ (iodide ion)
  • O²⁻ (oxide ion)
  • S²⁻ (sulfide ion)
  • N³⁻ (nitride ion)
  • OH⁻ (hydroxide ion)
  • CN⁻ (cyanide ion)
  • NO₃⁻ (nitrate ion)
  • SO₄²⁻ (sulfate ion)
  • PO₄³⁻ (phosphate ion)
  • CH₃COO⁻ (acetate ion)

Polyatomic Ions

• Groups of atoms that gain or lose electrons.
  • Ammonium (NH₄⁺) is a polyatomic cation.
  • Sulfate (SO₄²⁻) is a polyatomic anion.

Ionic Compounds

• Ionic compounds (e.g., NaCl) are typically formed between metals and nonmetals.
• Electrons are transferred from the metal to the nonmetal, resulting in oppositely charged ions that attract each other.
• Only empirical formulas are written for ionic compounds.

Writing Formulas

• Because compounds are electrically neutral:
  • The charge on the cation becomes the subscript on the anion.
  • The charge on the anion becomes the subscript on the cation.
  • If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

Chemical Nomenclature

• The system of naming compounds is called chemical nomenclature.
• Naming conventions will be covered for:
  • Ionic compounds
  • Acids
  • Binary molecular compounds
  • Simple organic compounds (alkanes and alcohols)

Inorganic Nomenclature

• Name the cation first. If the cation has more than one possible charge, indicate the charge using Roman numerals in parentheses. Polyatomic cations end in -ium.
• For monatomic anions, change the ending to -ide. For polyatomic ions, simply state the name of the polyatomic ion.

Patterns in Oxyanion Nomenclature

• For two oxyanions involving the same element:
  • The one with fewer oxygens ends in -ite (e.g., nitrite).
  • The one with more oxygens ends in -ate (e.g., nitrate).
  • NO₂⁻: nitrite; NO₃⁻: nitrate
  • SO₃²⁻: sulfite; SO₄²⁻: sulfate

Patterns in Oxyanion Nomenclature

• Central atoms on the second row bond to a maximum of three oxygens; those on the third row can bond to up to four.
• Charges increase as you move from right to left in the periodic table.

Patterns in Oxyanion Nomenclature

• The oxyanion with the second fewest oxygens ends in -ite (e.g., chlorite).
• The one with the second most oxygens ends in -ate (e.g., chlorate).
• The one with the fewest oxygens has the prefix hypo- and ends in -ite (e.g., hypochlorite).
• The one with the most oxygens has the prefix per- and ends in -ate (e.g., perchlorate).
  • ClO⁻ is hypochlorite.
  • ClO₂⁻ is chlorite.
  • ClO₃⁻ is chlorate.
  • ClO₄⁻ is perchlorate.

Acid Nomenclature

• If the anion ends in -ite, change the ending to -ous acid (e.g., chlorous acid).
  • HClO: hypochlorous acid
  • HClO2: chlorous acid
• If the anion ends in -ate, change the ending to -ic acid (e.g., chloric acid).
  • HClO3: chloric acid
  • HClO4: perchloric acid
• If the anion ends in -ide, change the ending to -ic acid and add the prefix hydro- (e.g., hydrochloric acid).
  • HCl: hydrochloric acid
  • HBr: hydrobromic acid
  • HI: hydroiodic acid

Nomenclature of Binary Molecular Compounds

• The element farther to the left in the periodic table (closer to metals) or lower in the same group is usually written first.
• A prefix denotes the number of atoms of each element in the compound (mono- is not used on the first element listed).

Nomenclature of Binary Compounds

• The ending of the second element is changed to -ide.
  • CO₂: carbon dioxide
  • CCl₄: carbon tetrachloride
• If the prefix ends with a or o, and the element name begins with a vowel, the two successive vowels are often elided into one.
  • N₂O₅: dinitrogen pentoxide
  • CO: carbon monoxide

Nomenclature of Organic Compounds: Alkanes

• Organic chemistry is the study of carbon.
• Organic chemistry has its own system of nomenclature.
• The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes.
• The first part of the names corresponds to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.), followed by -ane.

Nomenclature of Organic Compounds: Alcohols

• When a hydrogen in an alkane is replaced with a functional group (like –OH), the name is derived from the alkane name.
• The ending denotes the type of compound (an alcohol ends in -ol).

Nomenclature of Organic Compounds: Alcohols

• Isomers are molecules with the same chemical formula but different structures.
• 1-Propanol and 2-propanol have the oxygen atom connected to different carbon atoms but both have the formula C3H8O.