Chapter 2: Atoms, Molecules and Ions

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Chapter 2Atoms, Molecules and Ions
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The Language of Chemistry

• Atoms
• Composed of electrons, protons and neutrons
• Molecules
• Combinations of atoms
• Ions
• Charged particles

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Laws of Chemical Composition

• 1790 Antoine Lavoisier, The Father of Modern
  Chemistry
• Law of Conservation of Matter
• Total mass remains constant during a chemical
  reaction or
• Total mass of reactants total mass of products.

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Law of Conservation of Mass A Conceptual Example

• Jan Baptista van Helmont (15791644) first
  measured the mass of a young willow tree and,
  separately, the mass of a bucket of soil and then
  planted the tree in the bucket. After five years,
  he found that the tree had gained 75 kg in mass
  even though the soil had lost only 0.057 kg. He
  had added only water to the bucket, and so he
  concluded that all the mass gained by the tree
  had come from the water. Explain and criticize
  his conclusion.

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Laws of Chemical Composition

• Joseph Proust, Law of Constant Composition (Law
  of Definite Composition, or Definite Proportions)
• All samples of a compound have the same
  composition, or all samples have the same
  proportion by mass of the elements present.

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Law of Constant Composition Example

• Example CuHCO3 is ALWAYS
• 57.48 Cu, 5.43 C, 0.91 H and
• 36.18 O by mass

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John Dalton and the Atomic Theory of Matter

• Importance
• Explained Laws of Conservation of Mass and
  Constant Composition and
• extended them to cover another law.

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Main ideas of Daltons model

• 1. All matter consists of of small, indivisible
  particles called atoms.
• 2. All atoms of a given element are alike but
  atoms of any one element are different from the
  atoms of every other element.
• 3. Compounds are formed when atoms of different
  elements unite in small, whole-number ratios.
• 4. Chemical reactions involve rearrangement of
  atoms no atoms are created, destroyed or broken
  apart in a chemical reaction.
• According to Dalton, atoms are
• indivisible and indestructible.

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Daltons Atomic Theory Conservation of Mass and
Definite Proportions
six fluorine atoms and four hydrogen atoms
after reaction. Mass is conserved.
Six fluorine atoms and four hydrogen atoms before
reaction
HF always has one H atom and one F atom always
has the same proportions (119) by mass.
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Another Important Law

• Law of Multiple Proportions
• A given set of elements may combine to produce
  two or more different compounds, each with a
  unique composition.
• Example H2O (water) and
• H2O2 (hydrogen peroxide)

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Law of Multiple Proportions (contd)

• Four different oxides of nitrogen can be formed
  by combining 28 g of nitrogen with
• 16 g oxygen, forming Compound I
• 48 g oxygen, forming Compound II
• 64 g oxygen, forming Compound III
• 80 g oxygen, forming Compound IV

What is the ratio 16486480 expressed as small
whole numbers?

• Compounds IIV are N2O, N2O3, N2O4, N2O5

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Daltons Model of the Atom

• NO subatomic particles!

In modern atomic theory, the atom is divided into
protons, neutrons and electrons
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1897 JJ Thomson
Cause stream of negative particles that are
always the same, no matter what gas is used
Thomson experimented with CATHODE RAY TUBES
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1897 JJ Thomson
Mass to charge ratio for an electron m/c 5.69
x 10-9g/coulomb
Known as discoverer of the ELECTRONled to the
plum pudding model of the atom
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Millikan

• Obtained the charge of an electron, which coupled
  with Thomsons work, allowed the calculation of
  the mass of an electron.

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Millikans Conclusions

• Measured the charge of an electron
• 1.602 x 10-19 coulomb (C)
• Calculated the mass of an electron
• 9.109 x 10-31 kg

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• The modern view of the atom was developed by
  Ernest Rutherford of New Zealand (1871-1937).

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Ernest Rutherford
Canterbury University in Christchurch, NZ
Rutherford laboratory
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• Gold Foil Experiment

Screen 2.9
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(No Transcript)
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Rutherfords Main Conclusions

• 1. The atom is mostly empty space.
• All of the positive charge, and most of the mass,
  is concentrated in a very small volume
• THE NUCLEUS
• 3. Electrons are outside the nucleus.

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Protons
Mass of proton about same as an H atom (1 atomic
mass unit) Positive charge negative charge
from electrons in a neutral atom.
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Neutrons (Chadwick, 1932)

• the nucleus also contains neutrons particles
  with masses almost identical to protons but with
  no charge
• neutrons also help disperse the strong repulsion
  of positive charges

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Summary
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Atomic Symbols

• An atomic symbol represents the element.

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Al
26.981
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Mass Number, A

• The Mass Number (A)
• protons neutrons
• A boron atom can have A 5 p 5 n
  10 amu

Named as boron-10
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Atomic Number, Z

• Atomic number, Z, is the number of protons in the
  nucleus. (same for every atom of that element)

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Al
26.981
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Isotopes

• Atoms of the same element (same Z) but different
  mass number (A).
• Boron-10 has 5 p and 5 n 105B
• Boron-11 has 5 p and 6 n 115B

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Hydrogen Isotopes
Hydrogen has 3 isotopes
1 proton and 0 neutrons, protium
11H
1 proton and 1 neutron, deuterium
21H
1 proton and 2 neutrons, tritium radioactive
31H
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Isotopes Their Uses
Heart scans with radioactive technetium-99.
9943Tc Emits gamma rays
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Sample Problem

• Example 2.1 Write the atomic symbols for the
  following species
• a. the isotope of carbon with a mass of 13
• b. the nuclear symbol when Z 92 and the number
  of neutrons 146.

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Solution to Problem

• 13 C
• 6
• 238 U
• 92

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Ions

• Definition
• Atoms GAIN electrons to become negative ions, or
  anions.
• Atoms LOSE electrons to become positive ions, or
  cations.
• How are ions represented?
• Charges are always shown to upper right of
  symbol.

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Sample Problem

• Example 2.2 Write the atomic symbols for the
  following
• a. a species having 16 protons, 16 neutrons and
  18 electrons
• b. the phosphide ion (P) with an overall charge
  of -3

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Solution

• 32 S 2-
• 16
• 31 P 3-
• 15

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Atomic Mass

• F. An atomic mass unit (amu or u) is defined as
  exactly one-twelfth the mass of a carbon-12 atom
• 1 u 1.66054 1024 g
• The atomic mass of an element is the relative
  mass of an atom compared to a standard
  (carbon-12). It is NOT equal to the mass number!

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Atomic Mass Is Not Equal to Mass Number!!

• The atomic mass is a weighted average of the
  masses of the naturally occurring isotopes.
• (also called atomic weight)

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Al
26.981
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Atomic Mass

• Weighted average is the addition of the
  contributions from each isotope
• Isotopic Abundance is the percent or fraction of
  each isotope found in nature.

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Most Abundant Isotope
Usually can round atomic mass on p.t. to nearest
whole number
13
Al
26.981
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Atomic Mass

• Example 2.3 Determine the average atomic mass of
  magnesium which has three isotopes with the
  following masses 23.98 (78.6), 24.98 (10.1),
  25.98 (11.3).

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Radioactivity

• Radioactive isotopes are unstable
• These isotopes decay over time
• Emit other particles and are transformed into
  other elements
• Radioactive decay is not a chemical process!
• Particles emitted
• High speed electrons ß (beta) particles
• Alpha (a) particles helium nuclei
• Gamma (?) rays high energy light

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Nuclear Stability

• depends on the neutron/proton ratio
• For light elements, n/p is approximately 1
• For heavier elements, n/p is approximately 1.4/1

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Figure 2.5 The Nuclear Belt of Stability
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The Periodic Table Elements Organized

• Know location and description of
• groups or families
• periods or series
• metals, metalloids, nonmetals and their
  properties
• main group elements
• transition metals
• lanthanides and actinides

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Groups or Families

• Vertical columns are groups
• Numbered as 1-18 (new)
• Old system uses Roman numerals and A,B

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Periods or Series

• Horizontal rows are periods
• 7 periods
• total
• First period is H and He
• Second period is Li to Ne
• Etc.

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Group Names to Memorize

• - Group 1 (IA) alkali metals.
• - Group 2 (IIA) alkaline earth metals.
• - Group 17(VIIA) halogens.
• - Group 18 (VIIIA) noble gases

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Group 1A Alkali MetalsLi, Na, K, Rb, Cs
Reaction of potassium H2O
https//www.youtube.com/watch?voqMN3y8k9So
https//www.youtube.com/watch?vJy1DC6Euqj4
Cutting sodium metal
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Group 2A Alkaline Earth Metals Be, Mg, Ca, Sr,
Ba, Ra
Magnesium
Magnesium oxide
https//www.youtube.com/watch?vqSr39UwpELo
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Group 7A HalogensF, Cl, Br, I, At
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Group 8A Noble Gases He, Ne, Ar, Kr, Xe, Rn
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Regions of the Periodic Table
Metals are on the left of stair step
line NON-METALS are on the right of stair step
line
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ExceptionGroup 1A Hydrogen is a Non-metal!

• Shuttle main engines use H2 and O2

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Properties of Metals/Non-metals/Metalloids

• Metals-shiny,smooth, solid at room temperature,
  good conductors of heat and electricity,
  malleable and ductile.
• Metalloids (along stair step line) physical and
  chemical properties of both metals and nonmetals-
  B, Si, Ge, As, Sb, Te
• Nonmetals-low melting and boiling points,
  brittle, dull-looking solids, poor conductors of
  heat and electricity.

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The Periodic Table Elements Organized

• Main group elements -tall columns (Groups
  1,2,13,14,15,16,17,18)
• Transition metals- short columns (10)
• Lanthanides and actinides- long rows below main
  part of table.

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Transition Elements

• Lanthanides and actinides

Iron in air gives iron(III) oxide
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Periodic Table

• Dmitri Mendeleev developed the modern periodic
  table. Argued that element properties are
  periodic functions of their atomic weights.

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Periodic Table

• Periodic Law
• We now know that element properties are periodic
  functions of their ATOMIC NUMBERS.

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GermaniumPrediction vs. Observation
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Henry Moseley

• A student of Rutherfords
• Arranged the periodic table in order of
  increasing atomic number

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Molecules

• A molecule is a group of two or more atoms held
  together in a definite shape by covalent bonds.

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Empirical and Molecular Formulas

• Empirical formula the simplest whole number
  ratio of elements in a compound
• Molecular formula gives the ACTUAL number of
  each kind of atom in a molecule.
• Example
• Molecular formula of glucose C6H12O6
• Can divide all subscripts by 6, so the empirical
  formula is CH2O

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Structural Formulas

• Structural formulas show how atoms are attached
  to one another.

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Ions Atoms with a Charge

• Definition
• Cations positive ions
• Anions negative ions
• Polyatomic ion A group of atoms with a
  charge
• You must memorize all the polyatomic ions
  (structure, name and charge) found on your purple
  flashcard sheet!

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Ionic Compounds

• Ionic Compounds are cations and anions held
  together by electrostatic attraction.
• Their formulas are the simplest ratio of numbers
  of atoms (called an empirical formula) and
  represent one formula unit.

There is NO net charge in an ionic compound!
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Solutions of Ionic Compounds

• Solutions of Ionic Compounds are strong
  electrolytes their solutions conduct
  electricity.
• Non-electrolytes do not conduct electricity in
  water solution. (sugar, molecular compounds)

There is NO net charge in an ionic compound!
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Charge Balance of Ionic Compounds

• See Handout and practice worksheets.

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Monatomic Ions

• Group IA metals form ions of 1 charge.
• Group IIA metals form ions of 2 charge.
• Aluminum, a group IIIA metal, forms ions with a
  3 charge.
• Nonmetal ions of groups V, VI, and VII usually
  have charges of
• VA -3
• VIA -2
• VIIA -1

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• Atoms that are close to a noble gas (group 18)
  form ions that contain the same number of
  electrons as the neighboring noble gas atom
• Applies to Groups 1, 2, 16 and 17, plus Grp 13
  metals (e.g., Al 3) and Grp 15
  non-metals/metalloids (e.g., N 3-)

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• Some metal ions have gt one possible charge. A
  Roman numeral are used for the charge.
• If a metal only has ONE charge, a Roman numeral
  is NOT used.

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Symbols and Periodic Table Locations of Some
Monatomic Ions
Copper forms either copper(I) or copper(II) ions.
Titanium forms both titanium(II) and titanium(IV)
ions.
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Polyatomic Ions

• See handouts MUST MEMORIZE!!!

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Polyatomic Ions

• Oxyanions the anions are composed of oxygen and
  one other element

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Nitrate and Sulfate
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Polyatomic Ions

• Oxyanions the anions are composed of oxygen and
  one other element
• ExSO42- (sulfate), NO2- (nitrite) , MnO4-
  (permanganate)
• two oxyanions of the same element
• The anion with the smaller number of oxygens uses
  the roots of the element plus ite
• The higher number use the root plus ate
• Ex SO32- sulfite, NO2- nitrite,
  PO3-3 phosphite
• SO42- sulfate, NO3- nitrate,
  PO4-3 phosphate

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Four oxyanions

• There are four oxyanions containing Cl
• The middle two are named as two oxyanions
• The one with one less oxygen than the chlorite
  has a prefix of hypo
• The one with one more oxygen than the chlorate
  has a prefix of per
• Ex ClO- hypochlorite
• ClO2- chlorite
• ClO3- chlorate
• ClO4- perchlorate

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Naming binary compounds

• Use name of metal with no changes
• Change the name of the anion by taking the stem
  and add the suffix ide
• ex CI2 chlorine
• Cl- chlorine ( ine ) ide chloride
• Examples
• NaCl - sodium chloride
• MgCl2 - magnesium chloride

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Naming Binary Ionic Compounds

• Name the following binary ionic compounds
• Metal nonmetal compound
  name
• KI potassium iodine potassium
  iodide
• Li2S lithium sulfur lithium
  sulfide
• Mg3N2 magnesium nitrogen magnesium nitride

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Metals with multiple oxidation states

• Two methods Stock and classical system
• Stock system (used at CHS)
• metal name and the oxidation state in Roman
  numbers in parenthesis
• Ex Fe2 iron (II)
• Form compound by balancing charge of metal with
  correct number of nonmetals
• Ex CoCl3 cobalt(III) chloride
• Charge of metal charge of anion x subscript
• subscript of
  cation

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Example 2.4 Write names or formulas for
Sample Problem

• rubidium bromide AlCl3
• barium nitride Ca3P2
• cobalt (II) bromide NaI
• Strontium oxide PbS2

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Solution

• RbBr Aluminum Chloride
• Ba3N2 Calcium Phosphide
• CoBr2 Sodium Iodide
• SrO Lead (IV) sulfide

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Example 2.5 Write names or formulas for
Sample Problem

• (NH4)2S strontium hydroxide
• K2Cr2O7 cobalt (II) sulfate
• Al(NO2)3 calcium phosphate
• Fe(CN)2 tin (IV) carbonate

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Solution

• ammonium sulfide Sr(OH)2
• potassium dichromate CoSO4
• aluminum nitrite Ca3(PO4)2
• Iron (II) cyanide Sn(CO3)2

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Binary Molecular Compounds

• (Two nonmetals bonded together may also include
  a metalloid in formula)
• e.g., CO, NO, HF, SiO2
• a. First symbol is usually element to furthest
  left in p.t.
• b. Numbers of atoms indicated by subscripts are
  written as prefixes.

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Binary Molecular Compounds

• The name consists of two words.
• Directions
• 1. Write name of first element preceded by
  prefix EXCEPT do not write mono- if only have ONE
  of first element.
• 2. Name of second element ends with -ide is
  also preceded by prefix.

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Names of Binary Compounds
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Example

• Consider the compounds CO and CO2

Name the element that appears first in the
formula CARBON The second element has an
altered name retain the stem of the element name
and replace the ending by -ide OXYGEN
? OXIDE
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Sample Problem
Example 2.4 Write names or formulas for the
following B2O3 tetraphosphorus
pentachloride AsO5 dihydrogen monoxide
As2O7
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Solution

• Diboron trioxide P4Cl5
• Arsenic pentoxide H2O
• Diarsenic heptoxide

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Acknowledgements

• Thomson/Brooks Cole (Textbook Publishers)
• Mark P. Heitz, State University of New York at
  Brockport (Prentice Hall, Book Publishers)