Atoms, Molecules, and Ions - Notes

Atomic Theory of Matter

• Democritus: Believed in smallest particle "atomos" (uncuttable).

• John Dalton's Atomic Theory (early 1800s):

  • Law of constant composition

  • Law of conservation of mass

  • Law of multiple proportions

Laws

• Law of Constant Composition: Compounds have a definite composition; the relative number of atoms of each element is the same in any sample.

  • Discovered by Joseph Proust.

• Law of Conservation of Mass: Total mass of substances at the end of a chemical process equals the mass before the process.

  • Discovered by Antoine Lavoisier.

• 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.

  • Compounds from the same elements cannot have the same relative number of atoms.

Dalton’s Atomic Theory

• Each element consists of extremely small particles called atoms.

• All atoms of a given element are identical in mass and properties, but differ from atoms of other elements.

• Atoms are not changed into different elements by chemical reactions; atoms are neither created nor destroyed.

• Atoms combine to form compounds with a specific relative number and kind of atoms.

Discovery of Subatomic Particles

• Dalton's view: Atom was the smallest particle.

• Discoveries leading to smaller particles:

  • Electrons and cathode rays

  • Radioactivity

  • Nucleus, protons, and neutrons

Electron (Cathode Rays)

• Streams of negatively charged particles emanate from cathode tubes.

• J. J. Thomson credited with discovery (1897).

• Thomson measured charge/mass ratio: $1.76 × 10^8$ coulombs/gram (C/g).

Millikan Oil-Drop Experiment

• Robert Millikan determined the charge on the electron in 1909.

Radioactivity

• Spontaneous emission of high-energy radiation by an atom.

• Observed by Henri Becquerel; studied by Marie and Pierre Curie.

• Showed atoms have subatomic particles and associated energy.

• Three types of radiation (Ernest Rutherford):

  • α particles (positively charged)

  • β particles (negatively charged, like electrons)

  • γ rays (uncharged)

The Atom (circa 1900)

• Plum pudding model (J. J. Thomson): Positive sphere with negative electrons embedded.

Discovery of the Nucleus

• Ernest Rutherford shot α particles at gold foil.

Nuclear Atom

• Some α particles deflected at large angles.

• Postulates:

  • Small, dense, positive center (nucleus) with electrons around the outside.

  • Most of the atom is space.

  • Atoms are very small ranging from 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); electron mass is negligible (relative mass 0).

• Protons and neutrons are in the nucleus; electrons travel around the nucleus.

Atomic Number

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

Atoms of an Element

• Elements have one or two-letter symbols (first letter capitalized). C is the symbol for carbon.

• All atoms of the same element have the same number of protons, which is called the atomic number. It is written as a subscript BEFORE the symbol. $^{6}C$ is the atomic number for carbon.

• The mass number is the total number of protons and neutrons in the nucleus of an atom and is written as a superscript BEFORE the symbol.

Isotopes

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

Atomic Mass Unit (amu)

• Atoms have extremely small masses.

• The heaviest known atoms have a mass of approximately $4 × 10^{–22}$ g.

• 1 amu = $1.66054 × 10^{–24}$ g. (atomic mass unit)

Atomic Weight

• Average mass using all isotopes of an element weighted by their relative abundances.

• Atomic Weight = Ʃ [(isotope mass) × (fractional natural abundance)] for ALL isotopes.

• Masses are compared to C-12 (6 protons and 6 neutrons) being exactly 12.

Atomic Weight Measurement

• Measured using a mass spectrometer.

Periodic Table

• Organization of elements by atomic number.

• Periods: horizontal rows.

• Groups: vertical columns with similar properties.

Reading the Periodic Table

• Atomic number ABOVE the symbol.

• Atomic weight BELOW the symbol.

Organization of the Periodic Table

• Rows: periods, Columns: groups

• Elements in the same group have similar chemical properties.

Periodicity

• Repeating pattern of properties and reactivity.

Groups

• Alkali metals (Group 1A)

• Alkaline earth metals (Group 2A)

• Chalcogens (Group 6A)

• Halogens (Group 7A)

• Noble gases (Group 8A)

Periodic Table

• Metals: left side (shiny, conduct heat/electricity, solids except mercury).

• Nonmetals: right side (can be solid, liquid, or gas).

• Metalloids: on steplike line (properties like metals and nonmetals).

Chemical Formulas

• Subscript indicates the number of atoms of that element.

• Molecular compounds: composed of molecules and almost always contain only nonmetals.

Diatomic Molecules

• Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, and Iodine.

Types of Formulas

• Empirical: lowest whole-number ratio of atoms.

• Molecular: exact number of atoms of each element.

Picturing Molecules

• Structural formulas show the order in which atoms are attached (do NOT depict 3D shape).

• Perspective drawings, ball-and-stick models, and space-filling models show 3D order.

Ions

• Atoms gain/lose electrons.

• Cations: formed when electrons are lost (metals).

• Anions: formed when electrons are gained (nonmetals except noble gases).

Polyatomic Ions

• Group of atoms gain/lose electrons.

• Polyatomic cation: Ammonium ($NH_4 ^+$).

• Polyatomic anion: Sulfate ($SO_4 ^{2-}$).

Ionic Compounds

• Formed between metals and nonmetals.

• Electrons transferred from metal to nonmetal; oppositely charged ions attract.

• Only empirical formulas are written.

Writing Formulas

• 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.

  • Divide them by the greatest common factor.

Chemical Nomenclature

• System of naming compounds.

  1. Ionic compounds

  2. Acids

  3. Binary Molecular Compounds

  4. Simple Organic Compounds

  • Alkanes

  • Alcohols

Inorganic Nomenclature

• Name cation (Roman numeral for charge if multiple possibilities).

• Anion:

  • Element: change ending to -ide.

  • Polyatomic ion: write the name of the polyatomic ion.

Patterns in Oxyanion Nomenclature

• Two oxyanions with the same element:

  • Fewer oxygens: -ite.

  • More oxygens: -ate.

• ex: $NO<em>2 ^−$: nitrite; $NO</em>3 ^−$: nitrate; $SO<em>3 ^{2−}$: sulfite; $SO</em>4 ^{2−}$: sulfate

• Elements on the second row have a bond to, at most, three oxygens; those on the third row take up to four.

Patterns in Oxyanion Nomenclature

• The one with the second fewest oxygens ends in -ite: $ClO_2 ^−$ is chlorite.

• The one with the second most oxygens ends in -ate: $ClO_3 ^−$ is chlorate.

• The one with the fewest oxygens has the prefix hypo- and ends in -ite: ClO– is hypochlorite.

• The one with the most oxygens has the prefix per- and ends in -ate: $ClO_4 ^−$ is perchlorate.

Acid Nomenclature

• Anion ends in -ite: change to -ous acid (e.g., HClO: hypochlorous acid, $HClO_2$: chlorous acid).

• Anion ends in -ate: change to -ic acid (e.g., $HClO<em>3$: chloric acid, $HClO</em>4$: perchloric acid).

• Anion ends in -ide: change to -ic acid and add hydro- prefix (e.g., HCl: hydrochloric acid, HBr: hydrobromic acid, HI: hydroiodic acid).

Nomenclature of Binary Molecular Compounds

• The element farther to the left (closer to the metals) in the periodic table is written first.

• A prefix indicates the number of atoms (mono- not used on the first element).

• The ending of the second element changes to -ide.

• ex: $CO<em>2$: carbon dioxide, $CCl</em>4$: carbon tetrachloride

• If prefix ends with a or o and element begins with a vowel, the vowels are often combined as one. – $N<em>2O</em>5$: dinitrogen pentoxide – CO: carbon monoxide

Nomenclature of Organic Compounds: Alkanes

• Organic chemistry is the study of carbon.

• The simplest hydrocarbons contain only carbon and hydrogen are alkanes.

Nomenclature of Organic Compounds: Alcohols

• Replacing a hydrogen in an alkane with –OH (alcohol) derives the name from the alkane, ending in -ol.

Nomenclature of Organic Compounds: Alcohols

• Isomers: molecules with the same chemical formula but different structures.

  • ex: 1-Propanol and 2-propanol ($C<em>3H</em>8O$).