Atoms, Molecules and Ions

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Atoms, Molecules and Ions

• Lecture 2
• ???

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Atoms, Molecules and Ions

• 2.1 The Early History of Chemistry
• 2.2 Fundamental Chemical Laws
• 2.3 Daltons Atomic Theory
• 2.4 Cannizzaros Interpretation
• 2.5 Early Experiments to Characterize the Atom
• 2.6 The Modern View of Atomic Structure
• 2.7 Molecules and Ions
• 2.8 An Introduction to the Periodic Table
• 2.9 Naming Simple Compounds

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Early history of chemistry

• 400 B.C., the Greeks
• all matter was composed of 4 fundamental
  substances fire, earth, water and air
• Fundamentals of modern chemistry
• laid in the 16th century
• development of systematic metallurgy (extraction
  of metals from ores)
• Irish scientist Robert Boyle (1627-1691)
• first chemist who performed truly quantitative
  experiments the relationship between P and V of
  gases (the book the Sceptical Chemist in 1661)

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Early history of chemistry

• Combustion evokes intense interests in 17 18
  centuries
• German chemist Georg Stahl (1660-1734) a
  substance burning in a closed container
  eventually stopped burning because the air in the
  container became saturated with phlogiston.
• English scientist Joseph Priestley (1733-1804)
  discovered oxygen gas (originally called
  dephlogisticated air)

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Priestley Medal
Source Roald Hoffman, Cornell University
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Fundamental chemical laws

• A French chemist Antoine Lavoisier (1743-1794)
• named oxygen
• carefully weighted the reactant and products of
  various reactions
• mass is neither created nor destroyed ? LAW of
  CONSERVATION of MASS
• published Elementary Treatise on Chemistry in
  1789, the first modern chemistry textbook

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Fundamental chemical laws

• An English schoolteacher John Dalton (1766-1844)
• Published A New System of Chemical Philosophy
  in 1808
• Atomic theory
• Each element is made up of tiny particles called
  atoms
• The atoms of a given element are identical the
  atoms of different elements are different in some
  fundamental way(s)
• Chemical compounds are formed when atoms combine
  with each other. A given compound always has the
  same relative numbers and types of atoms
• Chemical reactions involve reorganization of the
  atoms changes in the way they are bound
  together. The atoms themselves are not changed in
  a chemical reaction.

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Fundamental chemical laws

• Determining absolute formulas for compounds
• French chemist Joseph Gay-Lussac (1778-1850)
• Italian chemist Amedeo Avogadro (1776-1856)
• Avogadros hypothesis at the same T and P, equal
  V of different gases contain the same number of
  particles.
• No general agreement exists concerning the
  formula for elements
• Chaos reigned in the first half of the
  nineteenth century
• Mess till the leadership of the Italian chemist
  Stanislao Cannizzaro (1826-1910)

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Cannizzaros Interpretation

• Italian chemist Stanislao Cannizzaro
• Guided by two beliefs
• Compounds contained whole numbers of atoms as
  Dalton postulated
• Avogadros hypothesis was correct-equal volumes
  of gases under the same conditions contain the
  same number of molecules
• Relative atomic mass to hydrogen (arbitrary
  assigned to be 2) was determined

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Early Experiments to Characterize the Atom

• English physicist J.J. Thomson (1856-1940)
  studied electrical discharges in partially
  evacuated tube cathode-ray tubes during the
  period from 1898 to 1903

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Cathode-ray tubes

• Thomson postulated that the ray was a stream of
  negatively charged particles, now called
  electrons.
• Deflection of the beam in a magnetic field

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Thomsons Plum pudding model

• An atom consisted of a diffuse cloud of positive
  charge with the negative electrons embedded
  randomly in it.

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Radioactivity

• 3 types of radioactive emission (early 20th
  century)
• gamma (?) rays, beta (ß) particles, and alpha (a)
  particles.
• A ?ray is high-energy light a ß particle is a
  high-speed electron and an a particle has a 2
  charge (i.e., a charge twice that of the electron
  and with the opposite sign). The mass of an a
  particle is 7300 times that of the electron.

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Nuclear atom

• In 1911 Ernest Rutherford carried out an
  experiment to test Thomsons plum pudding model.

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Nuclear atom

• Most of the a particles passed straight through,
  many of the particles were deflected at large
  angles

Thomsons plum pudding model ? X
Actual result ? proposed model
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The Modern View of Atomic Structure An
Introduction

• The simplest view of the atom is that it consists
  of a tiny nucleus with a diameter of about 10-13
  cm and electrons that move about the nucleus at
  an average distance of about 10-8 cm away from it

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• If all atoms are composed of these same
  components, why do different atoms have different
  chemical properties?
• The number and the arrangement of the electrons
• The electrons constitute most of the atomic
  volume
• The number of electrons possessed by a given atom
  greatly affects its ability to interact with
  other atoms

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Isotopes

• Atoms with the same number of protons but
  different numbers of neutrons

Sodium-23 is the only naturally occurring form of
sodium. Sodium-24 does not occur naturally but
can be made artificially.
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A
Z

• where the atomic number Z (number of protons) is
  written as a subscript and the mass number A (the
  total number of protons and neutrons) is written
  as a superscript. (The particular atom
  represented here is called sodium-23. It has 11
  electrons, 11 protons, and 12 neutrons.)

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Molecules and Ions

• Chemical bonds
• Atoms have electrons and these electrons
  participate in the bonding of one atom to another
• The forces that hold atoms together in compounds
• Covalent bonds bonds formed by sharing electrons
  
• Molecule
• Collection of atoms
• The simplest method to present a molecule is the
  chemical formula
• CO2 each molecule contains 1 atom of carbon and
  2 atoms of oxygen

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Structural formula
Solid lines in the plane of the paper Dashed
lines behind the plane of the paper Wedges in
front of the plane of the page
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• Space-filling model
• shows both the relative sizes of the atoms in the
  molecule and their spatial relationships

• Ball-and-stick models

Methane
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Ion

• An atom or group of atoms that has a net positive
  or negative charge

cation
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anion

• Because anions and cations have opposite charges,
  they attract each other. This force of attraction
  between oppositely charged ions is called ionic
  bonding.

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Ionic solid (also named salt)

• A solid consisting of oppositely charged ions
• can consist of simple ions, as in sodium chloride
  (NaCl), or of polyatomic (many-atom) ions as
  ammonium nitrate (NH4NO3)

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Periodic Table
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Periodic Table

• Shows all the known elements
• the letter is the symbol for that element
• the number is the atomic number (number of
  protons) for that element
• Horizontal rows in the table are called periods
• first period contains H and He
• second period contains elements Li through Ne

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Periodic Table

• Most of the elements are metals
• Shown in the center of the periodic table
• Efficient conduction of heat and electricity
• Malleability (can be hammered into thin sheets)
• Ductility (can be pulled into wire)
• (often) a lustrous appearance
• Relatively few nonmetals appear in the upper
  right-hand corner of the table
• tend to gain electrons to form anions in the
  reactions with metals
• often bond to each other by forming covalent bonds

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Periodic Table

• Elements in the same vertical columns (called
  groups or families) have similar chemical
  properties
• Ex alkali metals (Group 1A) Li, Na, K, Rb, Cs
  and Fr - are very active elements that readily
  form ions with a 1 charge when they react with
  nonmetals
• Ex alkali earth metals (Group 2A) Be, Mg, Ca,
  Sr, Ba and Ra - form ions with a 2 charge when
  they react with nonmetals
• Ex halogens (Group 7A) F, Cl, Br, I and At -
  form diatomic molecules. React with metals to
  form salts containing ions with a 1- charge.
• Ex Noble gases (Group 8A) He, Ne, Ar, Kr, Xe
  and Rn - form monatomic gases and have little
  chemical reactivity

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Naming Simple Compounds Binary compounds

• Ionic compound (Type I)
• The metal involved forms only a single type of
  cation
• The cation is always named first and the anion
  second
• A cation takes its name from the name of the
  element
• Ex Na is called sodium
• An anion is named by taking the first part of the
  element name and adding ide.
• Ex Cl- ion is called chloride.

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Ionic compound (Type II)

• Most commonly occurs for compounds containing
  transition metals, which often form more than one
  cation
• Ex The compound FeCl2 contains Fe2 ions, and
  the compound FeCl3 contains Fe3 ions.
• Naming system (I)
• iron (II) chloride (FeCl2) and iron (III)
  chloride (FeCl3)
• Roman numeral indicates the charge of the cation
• Naming system (II)
• Ferrous chloride (FeCl2) and ferric chloride
  (FeCl3)
• The ion with the higher charge has a name ending
  in ic, and the one with the lower charge has a
  name ending in ous
• common transition metals that form only one ion
  are zinc (Zn2) and silver (Ag)

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Naming binary ionic compounds
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Example 2.2

• Give the systematic name of each of the following
  compounds.
• a. CoBr2 b. CaCl2 c. Al2O3 d. CrCl3
• Solution

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• Polyatomic ions are assigned special names that
  must be memorized to name the compounds
  containing them.
• Ex the compound ammonium nitrate (NH4NO3)
  contains the polyatomic ions NH4 and NO3-
• Oxyanions several series of anions contain an
  atom of a given element and different numbers of
  oxygen atoms
• When there are two members in such a series, the
  name of the one with the smaller number of oxygen
  atoms ends in ite, and the name of the one with
  the larger number ends in ate
• Ex sulfite (SO32-) and sulfate (SO42-).
• When more than two oxyanions make up a series,
  hypo- ( less than) and per-(more than) are used
  as prefixed to name the members of the series
  with the fewest and the most oxygen atoms,
  respectively.

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Ionic compounds with Polyatomic Ions
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Binary covalent compound (Type III)

• Containing two nonmetals (i.e., do not contain
  ions)
• Naming rules
• The first element in the formula is named first,
  using the full element name.
• The second element is named as if it were an
  anion.
• Prefixes are used to denote the numbers of atoms
  present
• The prefix mono- is never used for naming the
  first element.
• Ex CO is carbon monoxide, not monocarbon monoxide

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• avoid awkward pronunciations, we often drop the
  final o or a of the prefix when the element
  begins with a vowel. For example, N2O4 is called
  dinitrogen tetroxide, not dinitrogen tetraoxide
  and CO is called carbon monoxide, not carbon
  monooxide.

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Naming compounds
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Naming compounds
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Example 2.3

• Give the systematic name of each of the following
  compounds.
• a. Na2SO4 b. KH2PO4 c. Fe(NO3)3 d.
  Mn(OH)2e. Na2SO3 f. Na2CO3 g. NaHCO3
  h. CsClO4i. NaOCl j. Na2SeO4 k.KBrO3
• Solution

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Example 2.4

• Given the following systematic names, write the
  formula for each compound.
• a. ammonium sulfate
• b. vanadium (V) fluoride
• c. dioxygen difluoride
• d. rubidium peroxide
• e. gallium oxide

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Naming Acids

• Acid
• When dissolved in water, certain molecules
  produce a solution containing free H ions
  (protons).
• The rules for naming acids depend on whether the
  anion contains oxygen
• If the anion does not contain oxygen, the acid is
  named with the prefix hydro- and the suffix ic.
• Ex HCl called hydrochloric acid, HCN called
  hydrocyanic acid and H2S called hydrosulfuric
  acid
• When the anion contains oxygen, the acid name is
  formed from the root name of the anion with a
  suffix of ic or ous.

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Anion containing oxygen

• If the anion name ends in ate, the acid name
  ends with ic (or sometimes ric).
• Ex H2SO4 contains the sulfate anion (SO4 2-) and
  is called sulfuric acid H3PO4 contains the
  phosphate anion (PO4 3-) and is called phosphoric
  acid and HC2H3O2 contains the acetate ion
  (C2H3O2-) and is called acetic acid
• If the anion has an ite ending, the acid name
  ends with ous.
• Ex H2SO3, which contains sulfite (SO3 2-), is
  named sulfurous acid and HNO2, which contains
  nitrite (NO2-), is named nitrous acid.

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Acid naming