Chapter 2 Atoms, Molecules, and Ions

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Chapter 2Atoms, Molecules, and Ions
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Preview

• Reviewing the fundamental chemical laws and their
  historical impacts.
• Understanding the modern view of atomic
  structure, molecules, and ions.
• Being familiar with the periodic table.
• Naming simple compounds, ionic compounds, and
  getting the chemical formulas from the names.

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The Early History of Chemistry
Chapter 2 Section 1

• Greeks were the first to pointed out to the
  concept of that matter is composed from atoms.
• The 4 fundamental substances.
• They didnt carry out experiments.
• Alchemists Changing coal into gold,
• They discovered some elements.
• Modern Chemistry
• 16th Century Extraction of metals from ores
  (metallurgy).
• 17th Century Robert Boyle performed quantitative
  experiments, introduced definitions for the term
  element.
• 18th Century CO2 by Stahl and O2 by J.
  Priestley. gt The concept of combustion. N2 and
  H2 gases were also identified.

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Fundamental Chemical Laws
Chapter 2 Section 2

• Late 18th century Lavoisier characterized
  quantitatively the combustion process. He named
  the O2 gas with its name.
• He carried out experiments (weighing products
  and reactants)
• gt conservation of mass.
• In the 19th century
• Law of conservation of mass by Lavoisier Mass is
  neither created nor destroyed.
• Law of definite proportion by Proust A compound
  always contains exactly the same proportion of
  elements by mass (copper carbonate). For example,
  he found that in CuCO3, it is always 5.3 parts by
  mass copper, 4 parts by mass oxygen, and 1 part
  by mass carbon.
• Law of multiple proportions by Dalton When two
  elements form a series of compounds, the ratios
  of the masses of the second element that combine
  with 1 gram of the first element can always be
  reduced to small whole numbers.

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What Did Dalton Observe in CO Molecules?
Chapter 2 Section 2
Mass of oxygen that combines with 1g of carbon
Ratio of mass of oxygen that combines with 1g of
carbon
1.33g
or
etc.
2.66g
or
etc.
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Sample Exercise 2.1
Chapter 2 Section 2

• For several compounds of nitrogen (N) and oxygen
  (O), we have the following data

Study this example carefully. It is described in
details on pages 42-43.
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Sample Exercise 2.1
Chapter 2 Section 2
A
B
C
The mass ratios shown can be readily described on
basis of the ratios of number of atoms.
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Exercise 27 on page 70
Chapter 2 Section 2
Then C could be NH A could be NH6 B could be NH9
But we know Ammonia NH3
So C could be N3H A could be NH2 B could be NH3
Another possibility C could be NH2 A could be
NH12 B could be NH18
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One More Example about the Law of Multiple
Proportions
Chapter 2 Section 2
This can be thought of as atom ratios of F atoms
instead of mass ratios.
If compound (i) has an actual formula of SF2 ,
what would be the actual formulas of (ii) and
(iii) assuming there is only one sulfur atom in
each?
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Daltons Atomic Theory
Chapter 2 Section 3

• Early in the 19th century Dalton presented his
  theory.
• 1. An element is made up from atoms.
• 2. Atoms of a given element are identical, but
  are different for different elements.
• 3. Elements combine to form chemical compounds.
• 4. Chemical reactions involve reorganization of
  atoms, but atoms themselves dont change.
• Atoms Elements Molecules (Compounds)

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Daltons Atomic Theory
Chapter 2 Section 3

• Dalton started assigning masses (relative to
  hydrogen) for different substances.

Chemical formula for water proposed by Dalton was
OH
All of these masses were proved later to be
incorrect. However, it was a great step forward
to modern chemistry.
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Daltons Atomic Theory
Chapter 2 Section 3

• Dalton showed that water is composed of the
  elements oxygen and hydrogen with 8g of oxygen
  for every 1g of hydrogen.

OR

• Dalton assumed that nature would tend to be as
  simple as possible. Thus, he gave water the
  formula OH. As a result, he assigned hydrogen
  mass of 1g and oxygen mass of 8g.

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Joseph Gay-Lussacs Experiment
Chapter 2 Section 3

• J. Gay-Lussac (mid 19th century)

A representation of some of Gay-Lussac's
experimental results on combining gas volumes at
the same conditions of temperature and pressure.
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Avogadros Hypothesis (1811)
Chapter 2 Section 3

• At the same temperature and pressure, equal
  volumes of different gases contain the same
  number of particles.
• The volume of the gas is determined NOT by the
  size of the individual particles, BUT by the
  number of molecules present. In the gas phase,
  the distances between molecules are much larger
  than the size of the molecules themselves.
• 2 Volumes Hydrogen 1 Volume Oxygen ?
  2 Volumes Water
• With Avogadros hypothesis
• 2 Molecules Hydrogen 1 Molecule Oxygen ?
  2 Molecules Water

Each balloon has the same number of particles
(molecules)
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Avogadros Hypothesis (1811)
Chapter 2 Section 3

• Interpreting Gay-Lussacs experiment on the light
  of Avogadros hypothesis

The spheres represent atoms in the molecules. The
molecules are considered to be diatomic.
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Avogadros Hypothesis
Chapter 2 Section 3

• Avogadros hypothesis had not been accepted for
  half a century.
• During the 19th century
• More elements were discovered.
• It was proven that the correct atomic mass of
  carbon is 12 amu.
• More relative atomic masses were listed.
• Chemistry started to make sense. However, atoms
  were not yet characterized or observed.

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Early Experiments to Characterize the Atom
Chapter 2 Section 4

• The attempt to explain the concept of electrons.
• Thomsons Experiment called the cathode-ray
  tube.

with high voltage
Cathode
Anode
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Thomsons Experiment
Chapter 2 Section 4

• Thomson concluded important points about the
  structure of the atom.
• The ray is a stream of negatively charged
  particles (later on called electrons).
• All atoms must contain electrons.
• Since the atom is neutral overall, it must have a
  positively charged component.
• So he came up with the plum pudding model for
  the atom.

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Thomsons Experiment
Chapter 2 Section 4

• The Deflection of the ray by the negative pole of
  an applied electric field.

He measured the charge-to-mass ratio as
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Mass of the Electron
Chapter 2 Section 4

• Millikans Experiment.

Oil droplets
Mass of the electron is 9.1110-31 Kg
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Radioactivity
Chapter 2 Section 4

• Types of spontaneous radioactive emission
• a particles have 2 charge and have mass that is
  7300 time the mass of electron
• ß particles high-speed electrons.
• ? particles high-energy light.

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The Nuclear Atom
Chapter 2 Section 4

• Rutherford Experiment.

http//www.mhhe.com/physsci/chemistry/essentialche
mistry/flash/ruther14.swf
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The Nuclear Atom
Chapter 2 Section 4

• Rutherfords Model vs. Thomsons Model.

Thomsons model
Rutherfords model (The nuclear atom)
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The Modern View of Atomic Structure
Chapter 2 Section 5

• Main components of atoms
• Outside the nucleus
• Electrons are responsible for the chemistry of
  the atom.
• Inside the nucleus
• Protons are positively charged particles whose
  charge is equal in magnitude to that for
  electrons.
• Neutrons have the same mass as protons but have
  no charge.

The simplest view of the atom
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The Modern View of Atomic Structure
Chapter 2 Section 5

• Nucleus is very tiny in terms of size, though
  almost all the atomic mass is concentrated in it
  (very dense)!!
• If a nucleus were to have the size of a pea, it
  would weigh 250,000,000,000 kg!
• Chemical properties of various atoms differ
  based on 1) The numbers of electrons, protons
  and neutrons are in the atom , and 2) The type of
  arrangement the electrons possess within the atom.

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The Atom Symbol Sodium as an Example
Chapter 2 Section 5
Element symbol (Na) Sodium Mass number (A)
of protons of neutrons Atomic number (Z)
of protons
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Atomic Symbols from the Periodic Table
Chapter 2 Section 5
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Example The Sodium Atom
Chapter 2 Section 5
Mass number
Isotopes show almost identical properties
Atomic number
Ion (Cation)
Isotopes
of protons 11
of neutrons 12
of electrons 10
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Exercises
Chapter 2 Section 5

• Q47 (a) Atomic number of protons 63 gt
  Eu
• Atomic mass of protons neutrons
  63 88
• Atomic charge 63 60 3
• The symbol is
• Q50 For
• of protons 26
• of neutrons 53 26
• Net charge 2
• of electrons 26 2

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Molecules and Ions
Chapter 2 Section 6

• Atoms are held together to produce molecules.
• The force holds atoms together is called a
  chemical bond.
• Some types of chemical bonds are
• Covalent bonds Two atoms can form a bond by
  sharing electrons.
• Ionic bonds Two oppositely charged ions (a
  cation and an anion) can form a bond by
  attraction.

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Chemical Formulas and Covalent bonds
Chapter 2 Section 6
Covalent-bonded Molecules
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Ionic Bonds
Chapter 2 Section 6
Na
Cl
17 protons
11 protons
17 electrons
11 electrons
Cl-
Na
e-

Na
Cl-
In form of crystals, called ionic solid or
commonly known as salt
11 protons
17 protons
10 electrons
18 electrons
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An Introduction to the Periodic Table
Chapter 2 Section 7
Groups (Families) similar chemical properties
Periods (rows) based on the electron
configurations
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An Introduction to the Periodic Table
Chapter 2 Section 7

• Metals compose most of the periodic table. They
  have characteristic physical properties e.g.
• High heat and electric conduction.
• Malleability (hammered to sheet)
• Ductility (pulled into wires)
• Chemically metals tend to lose electrons to form
  ve ions. Fe2 , Fe3 , Na ,K , Ca2 .
• Nonmetals lack the physical properties of
  metals.
• tend to gain electrons to become ve ions, like
  Cl-, F-, O2-, S2-.
• tend to bond with each other by forming covalent
  bonds, such as Cl2, HCl, N2O, CO2 etc.
• react with metals to form salt (ionic bonds)
  NaCl, KI2, etc.

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An Introduction to the Periodic Table
Chapter 2 Section 7
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Naming Simple Compounds
Chapter 2 Section 8

• There are two types of names
• Common names sugar, table salt, alcohol, etc.
• Systematic names on the basis of well-defined
  rules.
• We are going to study how to name binary
  compounds.
• Ionic (Type I and Type II).
• Covalent (Type III).
• Acids.

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Binary Ionic Compounds (Type I)
Chapter 2 Section 8
-

1- Cations named first then anions. 2- Cation
element has the same name without change. 3-Use
ide root to the anion name. 4-Double check the
ionic charges to have the correct chemical
formula. 5-You will need to practice this
table. 6- You will need to be able to get names
from formulas and vise versa.
It contains a ve ion and a ve ion.
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Binary Ionic Compounds (Type I)
Chapter 2 Section 8
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Exercise 55 on Page 72
Chapter 2 Section 8

• Rb2O
• Rubidium oxide.
• CaS
• Calcium sulfide.
• AlI3
• Aluminum iodide.

• Strontium fluoride.
• SrF2
• Aluminum selenide.
• Al2Se3
• Magnesium phosphide.
• Mg3P2

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Binary Ionic Compounds (Type II)
Chapter 2 Section 8

• In type II, metals (normally positive ions) can
  form more than one type of cations.
• In this case the charge must be specified using
  Roman numerals (only with Type II)
• Examples
• CuCl
• Copper(I) chloride.
• CuCl2
• Copper(II) chloride.

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Common Cations and Anions Found in Types I and II
binary Compounds
Chapter 2 Section 8
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Ionic Compounds with Polyatomic Ions
Chapter 2 Section 8

• They work exactly like type I ionic compounds.
• They must be memorized!!

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Ionic Compounds with Polyatomic Ions
Chapter 2 Section 8
oxyanions
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Exercises 59 and 60 on Page 72
Chapter 2 Section 8

• BaSO3
• Barium sulfite
• K2Cr2O7
• Potassium dichromate
• KMnO4
• Potassium permanganate
• NaNO2
• Sodium nitrite

• Chromium(III) hydroxide
• Cr(OH)3
• Magnesium cyanide
• Mg(CN)2
• Lead(IV) carbonate
• Pb(CO3)2
• Ammonium acetate
• NH4C2H3O2

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Binary Covalent Compounds (Type III)
Chapter 2 Section 8

• With bonding formed between two nonmetals.
• The rule for Type I compounds is applied here.

• N2O Dinitrogen monoxide
• NO Nitrogen monoxide
• N2O5 Dinitrogen pentoxide
• SO2 Sulfur dioxide
• PCl3 Phosphorus trichloride

Only used for the second element
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Naming Binary Compounds
Chapter 2 Section 8
?Practice Sample Exercise 2.9 on pages
65-66.
Some compounds are never referred to their
systematic names, such as H2O and NH3.
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Naming Acids
Chapter 2 Section 8

• Acids when are dissolved in water, they give H
  ions (protons) in the solutions.
• Examples are HCl, HBr, HNO3, H2SO3, and HC2H3O2.
• How to recognize an acid?

H
X-
Where X is an anion
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Naming Acids
Chapter 2 Section 8

• HNO2.
• HNO3.
• H2SO3.
• HC2H3O2.

• Hydrochloric acid (HCl).
• Hydrobromic acid (HBr).

• Nitric acid.
• Acetic acid.

• Nitrous acid.
• Sulfurous acid.

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Naming Acids
Chapter 2 Section 8
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Exercise 75 on Page 73

• Lead(II) acetate
• Copper(II) sulfate
• Calcium oxide
• Magnesium sulfate
• Magnesium hydroxide
• Calcium sulfate
• Dinitrogen monoxide