Atoms, Molecules, and Ions

• Chemistry; The Central Science, Fifteenth Global Edition in SI Units covers the study of Atoms, Molecules, and Ions.

Atomic Theory of Matter

• Democritus, a Greek philosopher, proposed that all of nature comprised an indivisible smallest particle termed "atomos."
• Experiments during the 18th and 19th centuries by John Dalton led to the organized atomic theory in the early 1800s including:
• The law of constant composition.
• The law of conservation of mass
• The law of multiple proportions

Law of Constant Composition

• Compounds have a specific and definite composition
• The relative number of atoms of each element in a given compound remains constant across different samples.
• Joseph Proust discovered the law of Constant Composition.
• This law was one of the laws of Dalton's Atomic Theory.

Law of Conservation of Mass

• The total mass of substances at the end of a chemical process equates to the mass of substances present prior to the process, maintaining mass conservation.
• Antoine Lavoisier discovered that mass is conserved.
• This law was one of the laws of Dalton's Atomic Theory.

Law of Multiple Proportions

• When two elements (A and B) combine to form multiple compounds, the masses of B that combine with a fixed mass of A are in ratios of small whole numbers.
• John Dalton discovered this law when he was developing his atomic theory.
• Compounds composed of the same elements cannot have the same relative number of atoms.

Postulates of Dalton’s Atomic Theory

• Each element consists of extremely small particles referred to as atoms.
• Atoms of a given element are identical in mass and other properties, yet they differ from atoms of any other element.
• Chemical reactions involve the rearrangement of atoms; atoms are neither created nor destroyed during chemical reactions.
• Compounds arise from the combination of atoms from multiple elements, with a given compound always having the same relative number and kind of atom.

Discovery of Subatomic Particles

• Dalton viewed the atom as the smallest possible particle.
• Subsequently multiple discoveries showed the atom itself comprises smaller particles including:
• Electrons and cathode rays.
• Nucleus, protons, and neutrons.
• Radioactivity.

The Electron (Cathode Rays)

• Streams of negatively charged particles from cathode tubes cause fluorescence.
• J. J. Thomson discovered the particles in 1897.
• Thomson measured the charge/mass ratio of the electron to be 1.76 × 108 coulombs/gram (C/g).

Millikan Oil-Drop Experiment (Electrons)

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

Radioactivity

• Radioactivity denotes the spontaneous emission of high-energy radiation by an atom.
• Henri Becquerel discovered Radioactivity
• Marie and Pierre Curie further studied radioactivity.
• Radioactivity's discovery implied atoms contain more subatomic particles and energy.
• Ernest Rutherford discovered three types of radiation:
• Alpha (α) particles are positively charged.
• Beta (β) particles are negatively charged, like electrons.
• Gamma (γ) rays carry are uncharged

The Atom, circa 1900

• The prevailing theory was the "plum pudding" model put forward by J.J. Thomson.
• The model featured a positive sphere of matter with negative electrons embedded in it

Discovery of the Nucleus

• Ernest Rutherford shot α particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
• Since some alpha particles were deflected at large angles, Thomson's Plum Pudding model could not be correct.
• This lead to the nuclear view of the atom.

The Nuclear Atom

• Rutherford proposed a very dense, small positive center called the nucleus containing almost all an atom's mass, surrounded by electrons.
• Most of the volume of an atom is empty space.
• Atoms are extremely small (1-5 Å or 100-500 pm).
• Additional subatomic particles (protons and neutrons) were discovered in the nucleus.

Subatomic Particles

• Protons (+1) and electrons (-1) have a charge; neutrons are neutral.
• Protons and neutrons have comparable mass (relative mass 1), while electrons are negligible (relative mass 0).
• Protons and neutrons exist in the nucleus, with electrons orbiting around it.

Atomic Number

• Each element is identified by the number of protons in the nucleus of an atom.
• Since atoms have no net charge because the number of protons equals the number of electrons in an atom

Atoms of an Element

• Elements are abbreviated using one- or two-letter symbols, with the first letter capitalized.
• All atoms of the same element have the same number of protons, known as the atomic number.
• The atomic number sits as subscript before an element's symbol.
• Mass number is the sum of the number of protons and neutrons in an atom's nucleus and is written as a superscript before the element's symbol.

Isotopes

• Isotopes are atoms of the same element that differ in mass.
• Isotopes have different numbers of neutrons, but the identical quantity of protons.
• Almost 99% of the carbon found in nature is Carbon-12, ¹²C.

Atomic Mass Unit (u)

• Atoms possess minuscule masses.
• The most massive known atoms weigh approximately 4 × 10⁻²² g.
• A mass scale on the atomic level is used, where an atomic mass unit (u) is the base unit.
• 1 u = 1.66054 × 10⁻²⁴ g

Atomic Weight

• In practical applications, we use significant amounts of multiple atoms and molecules, and utilize average masses in calculations.
• The average mass utilizes all isotopes of an element proportionally to their relative abundances, called an element's atomic weight.
• Atomic Weight = ∑ [(isotope mass) × (fractional natural abundance)] for all isotopes.
• The masses of any atom are compared to Carbon-12 (six protons and six neutrons), set to 12.

Atomic Weight Measurement

• A mass spectrometer measures atomic and molecular weights.
• The spectrometer can determine the spectrum and abundances of chlorine isotopes.

Periodic Table

• The periodic table is a systematic organization of elements.
• With elements arrange based on atomic number

Reading the Periodic Table

• Boxes on the periodic table show each element's atomic number above the symbol and the element's atomic weight below.
• The table's rows are called periods and the columns are called groups, elements in the same group have similar chemical properties.

Periodicity

• Elements show repeating patterns of chemical properties referred to as periodicity.

Groups

• There are five groups that are known by special names:
• Alkali metals
• Alkaline earth metals
• Chalcogens
• Halogens
• Noble Gases

Metals

• Metals are on the left side of the periodic table.
• Metallic properties include:
  • Shiny Luster
  • Conducting heat and electricity
  • Solid form at room temperature (except mercury)

Nonmetals

• Nonmetals are on the right side of the periodic table (include hydrogen).
• They exist as solids (like carbon), liquids (like bromine), or gases (like neon) - room temperature depending.

Metalloids

• Elements on the steplike line of the periodic table are known as metalloids (except for Al, Po, and At).
• Metalloids possess chemical properties, that are either metallic or nonmetallic.

Chemical Formulas

• The subscript on the element's symbol tells the number of atoms of that number in the chemical compound.
• Molecular compounds consist of molecules that nearly always only contain nonmetals.

Diatomic Molecules

• Seven elements are found in nature as molecules containing two atoms:
  • Hydrogen
  • Nitrogen
  • Oxygen
  • Fluorine
  • Chlorine
  • Bromine
  • Iodine

Types of Formulas

• Empirical Formulas: The lowest ratio of atoms of each element in a compound
• Molecular Formulas: The exact number of atoms of each element in a compound
• If we know the molecular formula of a compound, we can determine its empirical formula, but this is difficult to do in reverse.

Picturing Molecules

• Structural formulas show the order in which atoms are attached, but do not represent a molecule's three-dimensional shape.
• Perspective Drawings, ball-and-stick model and space-filling models are visualizations that display the three-dimensional order of atoms in a compound.

ions

• An ion is formed when an atom/group of atoms gains or loses/loses electrons.
• Cations form when one or more electrons are lost
• Monatomic Cations are formed by metals losing electrons.
• Anions are formed when one or more electron is gained
• Monatomic Anions are formed by nonmetals.

Writing Formulas

• Compounds possess electrical neutrality.
• Determine a compound's formula by:
  • Making the charge of the cation the subscript of the anion.
  • The charge of the anion becomes the subscript of the cation; and,
  • Dividing the subscripts by their greatest common factor if they are not in the most reduced whole-number ratio.

Chemical Nomenclature

• Naming compounds is called chemical nomenclature.
• Naming includes:
  • Ionic Compounds
  • Acids
  • Binary Molecular Compounds
  • Simple Organic Compounds (Alkanes and Alcohols)

Inorganic Nomenclature

• Name the cation first. If a cation could have more than one charge, indicate the charge with Roman numerals in brackets. Polyatomic cations end in -ium.
• Change -ides for monatomic anions
• A ployatomic ion's name stays

Patterns in Oxyanion Nomenclature

• Many elements form more than one oxyanion.
• The one with less oxygen ends in -ite.
• The one with more oxygen ends in -ate.

Acid Nomenclature

• For an anion ends in -ide: Replace the -ide to -ic acid: add hydro- to the front.. Ex: HCl = hydrochloric acid.
• If it ends in -ite, change the ending to -ous acid. Ex: HClO = hypochlorous acid.
• If it ends in -ate, change the ending to -ic acid. Ex: HClO4: perchloric acid.

Nomenclature of Binary Molecular Compounds

• In binary compounds, the most metallic element is listed first.
• Greek prefixes denote the number of formula units of each element, with mono- typically not used on the first nonmetal.
• The suffix -ide is added on the last nonmetal.

Nomenclature of Organic Compounds: Alkanes

• Organic chemistry involves primarily the study of carbon and its chemistry has its system of nomenclature.
• Hydrocarbons, the simplest organic compounds only contain carbon and hydrogen
• The names are composed of the carbon name and the suffix -ane.
• The first parts of the names correspond to the number of carbon atoms in the molecule (-meth = 1 -eth =2 -prop = 3, etc.)

Nomenclature of Organic Compounds: Alcohols

• When hydrogen in an alkane is replaced with other functional group (like -OH), the name is derived from the name of the alkane, but the ending indicates the chemical class.
• An alcohol ends in -ol.
• Isomers are molecules with the same chemical formula in different conformations.