CHEMISTRY The Central Science 9th Edition

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CHEMISTRY The Central Science 9th Edition

• Chapter 2Atoms, Molecules, and Ions

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The Atomic Theory of Matter

• Daltons law of multiple proportions When two
  elements form different compounds, the mass ratio
  of the elements in one compound is related to the
  mass ratio in the other by a small whole number.
• Atomic theory
• Each element is composed of tiny particles called
  atoms
• All atoms of a given element are identical.
• In chemical reactions, the atoms are not changed.
• Compounds are formed when atoms of more than one
  element combine.

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The Discovery of Atomic Structure

• Atoms are the building blocks of matter.
• The ancient Greeks were the first to postulate
  that matter consists of indivisible constituents.
• Later scientists realized that the atom consisted
  of charged entities.

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The Modern View of Atomic Structure

• The atom consists of positive, negative, and
  neutral entities (protons, electrons, and
  neutrons).
• Protons and neutrons are located in the nucleus
  of the atom, which is small. Most of the mass of
  the atom is due to the nucleus.
• There can be a variable number of neutrons for
  the same number of protons. Isotopes have the
  same number of protons but different numbers of
  neutrons.
• Electrons are located outside of the nucleus.
  Most of the volume of the atom is due to
  electrons.

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The Atom
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Class Practice Problem

• The diameter of a U.S. penny is 19mm. The
  diameter of a copper atom, by comparison, is only
  2.6 angstroms (Å). How many copper atoms could
  be arranged side by side in a straight line
  across the diameter of a penny?

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Atomic Weights

• The Atomic Mass Scale
• 1H weighs 1.6735 x 10-24 g and 16O 2.6560 x 10-23
  g.
• We define mass of 12C exactly 12 amu.
• Using atomic mass units
• 1 amu 1.66054 x 10-24 g
• 1 g 6.02214 x 1023 amu

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Atomic Number, Mass Number, and Isotopes

• Atomic number (Z) number of protons in the
  nucleus.
• Mass number (A) total number of nucleons in the
  nucleus (i.e., protons and neutrons).
• By convention, for element X, we write ZAX.
• Isotopes have the same Z but different A.
• We find Z on the periodic table.

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Class Practice Problem

• How many protons, neutrons, and electrons are in
  an atom of 197Au?
• Hydrogen has three isotopes, with mass numbers 1,
  2, and 3. Write the complete chemical symbol for
  each of them.

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Atomic Weights

• Average Atomic Masses
• Relative atomic mass average masses of isotopes
• Naturally occurring C 98.892 12C 1.108
  13C.
• Average mass of C
• (0.98892)(12 amu) (0.0108)(13.00335) 12.011
  amu.
• Atomic weight (AW) is also known as average
  atomic mass (atomic weight).
• Atomic weights are listed on the periodic table.

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Arrangement of the Periodic Table

• The Periodic Table is used to organize the 114
  elements in a meaningful way.
• As a consequence of this organization, there are
  periodic properties associated with the periodic
  table.

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The Periodic Table
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Reading the Periodic Table

• Columns in the periodic table are called groups
  (numbered from 1A to 8A or 1 to 18).
• Rows in the periodic table are called periods.
• Metals are located on the left hand side of the
  periodic table (most of the elements are metals).
• Non-metals are located in the top right hand side
  of the periodic table.
• Elements with properties similar to both metals
  and non-metals are called metalloids and are
  located at the interface between the metals and
  non-metals.

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Properties of the Periodic Table

• Some of the groups in the periodic table are
  given special names.
• These names indicate the similarities between
  group members
• Group 1A Alkali metals.
• Group 2A Alkaline earth metals.
• Group 6A Chalcogens.
• Group 7A Halogens.
• Group 8A Noble gases.

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Molecules and Molecular Compounds

• Molecules are assemblies of two or more atoms
  bonded together.
• Each molecule has a chemical formula.
• The chemical formula indicates
• which atoms are found in the molecule, and
• in what proportion they are found.
• Compounds formed from molecules are molecular
  compounds.
• Molecules that contain two atoms of the same
  element bonded together are called diatomic
  molecules.

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Molecules and Molecular Compounds
Example of Diatomic Molecules
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Molecules and Molecular Compounds

• Molecular and Empirical Formulas
• Molecular formulas
• give the actual numbers and types of atoms in a
  molecule.
• Examples H2O, CO2, CO, CH4, H2O2, O2, O3, and
  C2H4.

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Molecules and Molecular Compounds

• Most molecular substances that we will study in
  this class contain only nonmetals.

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Molecules and Molecular Compounds

• Molecular and Empirical Formulas
• Empirical formulas
• give the relative numbers and types of atoms in a
  molecule.
• That is, they give the lowest whole number ratio
  of atoms in a molecule.
• Examples H2O, CO2, CO, CH4, HO, CH2.

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Molecules and Molecular Compounds

• Molecular and empirical formulas do not show how
  atoms are arranged when bonded together.

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Molecules and Molecular Compounds

• Picturing Molecules
• Molecules occupy three dimensional space.
• However, we often represent them in two
  dimensions.
• The structural formula gives the connectivity
  between individual atoms in the molecule.
• The structural formula may or may not be used to
  show the three dimensional shape of the molecule.
  
• If the structural formula does show the shape of
  the molecule, then either a perspective drawing,
  ball-and-stick model, or space-filling model is
  used.

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Molecules and Molecular Compounds
Representing Structure in Molecules
Accurately represents the angles at which
molecules are attached.
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Class Practice Exercise
The structural formula of propane and butane
is What is the chemical and empirical
formula for these molecules?
H
H
H
H
H
H
H
C
C
C
H
H
C
C
C
H
H
C
H
H
H
H
H
H
H
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Ions and Ionic Compounds

• When an atom or molecule loses electrons, it
  becomes positively charged.
• For example, when Na loses an electron it becomes
  Na.
• Positively charged ions are called cations.
• When an atom or molecule gains electrons, it
  becomes negatively charged.
• For example when Cl gains an electron it becomes
  Cl-.
• Negatively charged ions are called anions.
• An atom or molecule can lose more than one
  electron.
• When molecules loose electrons, polyatomic ions
  are formed.

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Ions and Ionic Compounds

• In general metal atoms tend to lose electrons to
  become cations nonmetal ions tend to gain
  electrons to form anions.
• Predicting Ionic Charge
• The number of electrons an atom loses is related
  to its position on the periodic table.

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Ions and Ionic Compounds
Predicting Ionic Charge
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Ions and Ionic Compounds

• Element Bonding
• The majority of chemistry involves the transfer
  of electrons between species.
• Example
• To form NaCl, the neutral sodium atom, Na, must
  lose an electron to become a cation Na.
• The electron cannot be lost entirely, so it is
  transferred to a chlorine atom, Cl, which then
  becomes an anion Cl-.
• The Na and Cl- ions are attracted to form an
  ionic NaCl lattice which crystallizes.
• NaCl is an example of an Ionic compound
  (consisting of positive and negatively charged
  atoms)

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Ions and Ionic Compounds
Crystal Structure of NaCl
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Ions and Ionic Compounds

• Ionic Compounds
• Important note that there are no easily
  identified NaCl molecules in the ionic lattice.
  Therefore, we cannot use molecular formulas to
  describe ionic substances.
• Writing the empirical formulas for ionic
  compounds
• you need to know the ions of which it is
  composed.
• The formula must reflect the electrical
  neutrality of the compound
• the total positive charge must equal the total
  negative charge
• Example Consider the formation of Mg3N2
• Mg loses two electrons to become Mg2
• Nitrogen gains three electrons to become N3-.
• For a neutral species, the number of electrons
  lost and gained must be equal.

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Ions and Ionic Compounds

• Writing the Empirical Formula
• However, Mg can only lose electrons in twos and N
  can only accept electrons in threes.
• Therefore, Mg needs to lose 6 electrons (2 ? 3)
  and N gain those 6 electrons (3 ? 2).
• I.e., 3Mg atoms need to form 3Mg2 ions (total 3
  ? 2 charges) and 2 N atoms need to form 2N3-
  ions (total 2 ? 3- charges).
• Therefore, the formula is Mg3N2.

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Controversy in Naming Inorganic/Organic Compounds
Organic compounds contain carbon. Inorganic
compounds don't. This definition is often given
but is no help at all. What do we make of carbon
dioxide, sodium cyanide, baking soda (sodium
bicarbonate), ...? Organic compounds contain
carbon-hydrogen bonds. Inorganic compounds
don't. This is a much better definition,
allowing us to call sodium acetylide "organic"
but calcium carbide "inorganic," but it doesn't
always work. Inorganic compounds contain metal
atoms. Organic compounds don't. This doesn't
really work any too well either. Even leaving the
huge field of organometallic chemistry out of the
running, are we really going to call soap (sodium
salts of fatty acids) or the lipid bilayers
forming cell membranes (again, salts of
long-chain organic acids) "inorganic"??? An
organic compound is whatever an organic chemist
says it is an inorganic compound is whatever an
inorganic chemist says it is.
Reproduced from http//www.madsci.org/posts/archiv
es/dec2000/975719013.Ch.r.html
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Naming Inorganic Compounds

• Naming of compounds, nomenclature, is divided
  into organic compounds (those containing C,
  usually in combination with hydrogen) and
  inorganic compounds (the rest of the periodic
  table).
• Naming Ionic Compounds
• Based on the names of the ions of which they are
  composed.
• Example, NaCl is called sodium chloride (based on
  Na and Cl- ions).
• The cation is written first and the anion is
  written last.
• Ions may be monoatomic or polyatomic.
• Vast majority of monoatomic cations are made from
  metals.
• These ions take the name of the element itself.

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Naming Inorganic Cations

• Cations formed from a metal have the same name as
  the metal.
• Example Na sodium ion.
• If the metal can form more than one cation, then
  the charge is indicated by a Roman numeral in
  parentheses in the name.
• Examples Cu copper(I) Cu2 copper(II)
  (Page 61).
• Most of the elements that can form more than one
  cation are the
• transition metals (3B to 2B).
• Or placing ous or ic at the end of the name to
  indicate the lower
• and higher, respectively, charged cation.
• Cations formed from non-metals (end in -ium).
• Example NH4 ammonium ion.

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Some Common Cations
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Naming Inorganic Anions

• Monoatomic anions (with only one atom) are named
  by dropping the ending of the name and replacing
  with -ide.
• Example Cl- is chloride ion.
• Polyatomic anions (with many atoms) containing
  oxygen end in -ate or -ite. (The one with more
  oxygen is called -ate.)
• Examples NO3- is nitrate, NO2- is nitrite.
• (Exceptions hydroxide (OH-), cyanide (CN-),
  peroxide (O22-).)

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Naming Polyatomic Inorganic ions

• Polyatomic anions containing oxygen with more
  than two members in the series are named as
  follows (in order of decreasing oxygen)
• per-.-ate
• -ate
• -ite
• hypo-.-ite
• Examples
• ClO4- perchlorate ion, ClO3- chlorate, ClO2-
  chlorite, ClO- hypochlorite.

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Road Map to Naming Monoatomic and Polyatomic
Anions
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Naming Inorganic Compounds

• Polyatomic anions containing oxygen with
  additional hydrogens are named by adding hydrogen
  or bi- (one H), dihydrogen (two H), etc., to the
  name as follows
• CO32- is the carbonate anion
• HCO3- is the hydrogen carbonate (or bicarbonate)
  anion.
• H2PO4- is the dihydrogen phosphate anion.

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Some Common Anions
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Names and Formulas for Acids

• The names of acids are related to the names of
  anions.
• Acids containing anions whose names end in
• -ide becomes hydro-.-ic acid
• -ate becomes -ic acid
• -ite becomes -ous acid.

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Names and Formulas of Binary Molecular Compounds

• Binary molecular compounds have two elements.
• The most metallic element is usually written
  first (i.e., the one to the farthest left on the
  periodic table). Exception NH3.
• If both elements are in the same group, the lower
  one is written first.
• Greek prefixes are used to indicate the number of
  atoms.