Atoms, Molecules, and Ions: Chapter 2

• Chapter 2 is about atoms, molecules, and ions.

Atomic Theory of Matter

• Democritus, a Greek philosopher, theorized that all of nature is composed of a smallest particle called "atomos", meaning uncuttable.
• Experiments in the 18th and 19th centuries resulted in John Dalton's organized atomic theory in the early 1800s.
• Dalton's atomic theory is 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, the relative number of atoms of each element in the compound is the same in any sample.
• Joseph Proust discovered it.
• This contributed to Dalton’s atomic theory.

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.
• Antoine Lavoisier discovered it.
• This contributed to Dalton’s atomic theory.

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.
• John Dalton discovered this while developing his atomic theory.
• Two or more compounds from the same elements cannot have the same relative number of atoms.

Dalton’s Atomic Theory Postulates

• Each element is composed of extremely small particles called atoms.
• All atoms of a given element are identical in mass and other properties, but atoms of one element are different from atoms of all other elements.
• Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
• Atoms of more than one element combine to form compounds, and a given compound always has the same relative number and kind of atoms.

Discovery of Subatomic Particles

• Dalton believed the atom was the smallest particle possible, but discoveries revealed that atoms are made of smaller particles.
• These smaller particles are electrons and cathode rays, radioactivity, and the nucleus, protons, and neutrons

The Electron (Cathode Rays)

• Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence.
• J. J. Thomson discovered this in 1897.
• The charge/mass ratio of the electron measured by Thomson is 1.76 × 108 coulombs/gram (C/g).

Millikan Oil-Drop Experiment (Electrons)

• Once the charge/mass ratio of the electron was known, the determination of either the charge or the mass of an electron would yield the other.
• Robert Millikan is credited with determining the charge on the electron in 1909.

Radioactivity

• Radioactivity is the spontaneous emission of high-energy radiation by an atom.
• Henri Becquerel first observed this.
• Marie and Pierre Curie also studied it.
• Radioactivity's discovery showed that the atom had more subatomic particles and energy associated with it.
• Ernest Rutherford discovered three types of radiation:
  • Alpha (α) particles are positively charged.
  • Beta (β) particles are negatively charged, similar to electrons.
  • Gamma (γ) rays are uncharged.

The Atom circa 1900

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

Discovery of the Nucleus

• Ernest Rutherford shot alpha particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
• Thomson's plum pudding model was thus incorrect.
• From this the nuclear model of the atom was theorized.

The Nuclear Atom

• Rutherford’s nuclear model of the atom states that:
  • There is a very small, dense positive center with electrons around the outside.
  • Most of the atom is space.
  • Atoms are very small, measuring 1–5 Å or 100–500 pm.
  • The nucleus contains other subatomic particles (protons and neutrons).
• Protons (+1) and electrons (-1) have a charge, while neutrons are neutral.
• Protons and neutrons have essentially the same mass (relative mass 1).
• The mass of an electron is so small it's often ignored (relative mass 0).
• Protons and neutrons are in the nucleus; electrons travel around it.

Atomic Number

• The atomic number is 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.

Atoms of an Element

• Elements are represented by a one or two letter symbol, with the first letter always capitalized. C is the symbol for carbon.
• For all atoms of the same element the number of protons is the same; this is the atomic number, written as a subscript BEFORE the element symbol.
• The mass number is the total number of protons and neutrons in the nucleus of an atom, written as a superscript BEFORE the element symbol.

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, with the heaviest known having a mass of approximately 4 × 10⁻²² g.
• A mass scale on the atomic level uses the atomic mass unit (u) as the base unit. 1 u = 1.66054 × 10⁻²⁴ g

Atomic Weight

• Average masses are used in calculations to measure quantities of atoms and molecules.
• Average mass is found using all isotopes of an element weighted by relative abundances.
• This average is the element’s atomic weight.
• Atomic Weight = ∑ [(isotope mass) × (fractional natural abundance)] for ALL isotopes.
• The masses of any atom are compared to carbon-12 (6 protons and 6 neutrons) being exactly 12.

Atomic Weight Measurement

• Atomic and molecular weight can be measured using a mass spectrometer.
• The spectrum of chlorine shows two isotopes; abundances can also be determined this way.

Periodic Table

• The periodic table is a systematic organization of the elements, arranged in order of atomic number.
• Boxes list the atomic number above the symbol.
• The atomic weight of an element is listed below the symbol.
• Rows are called periods.
• Columns are called groups.
• Elements in the same group have similar chemical properties.
• Properties of elements are noticed to display chemical reactivity.

Groups

• Alkali metals (Group 1)
• Alkaline metals (Group 2)
• Chalcogens metals (Group 16)
• Halogens metals (Group 17)
• Noble Gases metals (Group 18)
• Metals are on the left side of the periodic table
• Properties of metals include:
  • Shiny luster
  • Conducting heat and electricity
  • Solids (except mercury)
• Nonmetals are on the right side of the periodic table (Hydrogen is also a nonmetal, despite not being on the right side).
• Nonmetals can be solid (carbon), liquid (bromine), or gas (neon) at room temperature.
• Elements on the steplike line in the periodic table are metalloids (except Al, Po, and At).
• Metalloids' properties are sometimes those of metals, and sometimes nonmetals.

Chemical Formulas

• The subscript to the right of the symbol of an element tells 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.
• Seven elements occur naturally as molecules containing two atoms; they are known as diatomic molecules:
  • Hydrogen
  • Nitrogen
  • Oxygen
  • Fluorine
  • Chlorine
  • Bromine
  • Iodine

Types of Chemical 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.
• If the molecular formula of a compound is known, we can determine its empirical formula but the opposite is not true.

Picturing Molecules

• Structural formulas show the order in which atoms are attached, but do NOT depict 3D shape.
• Perspective drawings, ball-and-stick models, and space-filling models show the three-dimensional order of the atoms in a compound.

Ions

• An ion is an atom or group of atoms that has lost or gained electrons
• Cations are formed when losing at least one electron.
• Monatomic cations are formed by metals.
• Anions are formed when gaining at least one electron
• Monatomic anions are formed by nonmetals, except the noble gases.

Polyatomic Ions

• These are ions which are multiple atoms losing/gaining electrons
• Ammonium ≡ NH₄⁺ Polytomic Anion:
• Sulfate ≡ SO₄²⁻

Ionic Compounds

• Ionic compounds, such as NaCl, are generally formed between metals and nonmetals
• Electrons are transferred from the metal to the nonmetal and the oppositely charged ions attract each other
• Only empirical formulas are written for ionic compounds

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. If these subscripts are to the lowest whole-number ratio, divide them by the greatest common factor

Chemical Nomenclature

• Chemical nomenclature is naming of compounds.
• This inludes ionic compounds, acids, binary molecular compounds and simple organic compounds (alkanes & alcohols.)

Inorganic Nomenclature

• Write the name of the cation.
  • If the cation has more than one possible charge, write the charge as a Roman numeral in parentheses.
  • If it is a polyatomic cation, it will end in -ium.
• If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.

Patterns in Oxyanion Nomenclature

• When two oxyanions involve the same element:
  • The one with fewer oxygens ends in -ite
  • The one with more oxygens ends in -ate
  • NO₂⁻ is nitrite; NO₃⁻ is nitrate
  • SO₃²⁻ is sulfite; SO₄²⁻ is sulfate Central atoms on the second row have a bond to, at most, three oxygens.
• Atoms on the third row take up to four.
• Charges increase as you go from right to left.

Acid Nomenclature

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

Nomenclature of Binary Molecular Compounds

• The name of the element farther to the left in the periodic table (closer to the metals) or lower in the same group is usually written first.
• Prefixes are used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however).
• The ending on the second element is changed to -ide.
  • CO₂: carbon dioxide
  • CCl₄: carbon tetrachloride
• If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are elided into one.
  • N₂O₅: dinitrogen pentoxide
  • CO: carbon monoxide

Nomenclature of Organic Compounds: Alkanes

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

Nomenclature of Organic Compounds: Alcohols

• When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane.
• The ending denotes the type of compound (-ol for alcohols).
• Molecules with the same chemical formula but different structure are isomers.
• 1-Propanol and 2- propanol have the oxygen atom connected connected to different carbon atoms, but both have the formula C3H8O.