Chemistry Chapter 4

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Chemistry Chapter 4
The Structure of the Atoms
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Top Ten
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Table 3.1
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Table 3.3
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History of Chemistry

• 400 B.C. Greeks proposed 4 elements
• Earth
• Fire
• Water
• Air
• Next 2000 yearsalchemy
• During this period discoveries were made
• Hg, S, Sb
• prepared acids

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Controversial Greek Thinking!
Democritus (460-370 B.C)
-Matter is composed of atomos (now
atoms) -Atoms were homogeneous
indivisible -Could not answer what holds atoms
together
Aristotle (384 B.C.-322B.C.)
-Matter was continuous and indefinitely
divisible (did not believe in atoms) -Matter
made of earth, fire, air, water -Idea was
accepted for nearly 2000 years!
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Indivisible or Divisible?Democritis vs. Aristotle
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Late 1700s

• Most chemists accepted element definition
• Understood elements combined to form compounds
  with various properties
• Disagreed whether compounds are always in the
  same ratio

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What happened in 1790?

• Study of matter was revolutionized by new
  emphasis on Quantitative Analysis
• Aided by improved balances
• Measurements were actually ACCURATE!!!

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Robert Boyle

• Founder of Modern Chemistry (1627-1691)
• Took the Al out of Alchemy (although he started
  as one)
• First scientist to understand the importance of
  careful measurement
• Insisted science be based on experiments
• Famous for P1/V

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Antoine Lavoisier

• Father of Modern Chemistry (1743-1794)
• Recognized and named hydrogen and oxygen
• Introduced the metric system
• Wrote first list of elements and revised
  nomenclature
• Because of prominence in pre-revolutionary
  government, was beheaded at the height of the
  French revolution

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John DaltonBeginning of Modern Atomic Theory

• Englishman from a Quaker family (1766-1844)
• Revolutionized chemistry by emphasizing that
  atoms can have weights and weights can be
  measured (quantitative)
• Opened a school at age 12
• Color blind/researched
• Interested in botany
• Theory not accepted until 1905 Albert Einstein
  paper

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Daltons Atomic Theory (1808)

• Matter is composed of extremely small particles
  called atoms
• Atoms are indivisible and indestructible.

• Atoms of a given element are identical in size,
  mass, and chemical properties.
• Atoms of specific element are different from
  those of another element.
• Different atoms combine in simple whole-number
  ratios to form compounds.
• In chemical reactions, atoms are separated,
  combined, or rearranged

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Dalton vs. Today
Matter is composed of extremely small particles called atoms. True
Atoms are indivisible and indestructible. Made up of smaller particles (protons, neutrons, electrons) except in nuclear chemistry.
Atoms of a given elemet are identical in size, mass, and chemical properties. Atoms of a given element have same p and e-, but may differ in of neutrons
Atoms of a specific element are different from those of another element True, how we identify them
Different atoms combine in simple whole-number ratios to form compounds True, Law of Multiple Proportions
In chemical reactions, atoms are separated, combined, or rearranged True
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Law of Conservation of Mass
Mass is neither created nor destroyed during
chemical or physical reactions.
Total mass of reactants Total mass of products
Antoine Lavoisier
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Law of Multiple Proportions

• If two or more different compounds are composed
  of the same two elements, then the ratio is
  always small whole numbers. (CO, CO2)

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What does this mean? (Law of Definite Composition)

• 50.0 g sample of pure H2O decomposed into its
  elements
• would find 5.6 g H and 44.4 g oxygen
• mass would be
• mass H 5.60 g x 100 11.2 H
• total mass 50.0 g
• mass 0 44.4 g x 100 88.8 O
• total mass 50.0 g

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Law of Definite (or Constant) Composition

• The fact that a chemical compound contains the
  same elements in exactly the same proportions by
  mass regardless of the size of the sample or the
  source of the compound.

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Figure 3.2 Representation of NO, NO2, and N2O.
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• What does an atom look like?
• (Sketch it on your paper!)

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This is The Modern Atomic Model

• Atom The smallest particle of an element that
  retains the properties of the element
• Only seen by STM (Scanning Tunneling microscope)

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Subatomic Particles
Particle Charge Mass (amu) Location
Electron (e-) J.J.Thomson 1897 Nobel Prize 1906 Robert Milllikan (1910s) -1 5.486x10-4 amu 9.1 x 10-28 g 1/1840 of H Electron cloud
Proton (p) Thomson/Goldstein-1907 Rutherford 1920 1 1.007 amu 1.673 x 10-24 g Nucleus
Neutron (no) Chadwick 1932 Nobel Prize 1935 0 1.009 1.675 x 10-24 g Nucleus
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Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to
deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
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Figure 3.7 Schematic of a cathode ray tube.
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Some ModernCathode Ray Tubes
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Mass of the Electron
1909 Robert Millikan determines the mass of the
electron.
The oil drop apparatus
Mass of the electron is 9.1 x 10-28 g
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Conclusions from the Study of the Electron

• Cathode rays have identical properties
  regardless of the element used to produce them.
  All elements must contain identically charged
  electrons.
• Atoms are neutral, so there must be positive
  particles in the atom to balance the negative
  charge of the electrons
• Electrons have so little mass that atoms must
  contain other particles that account for most of
  the mass

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Thomsons Atomic Model
Thomson believed that the electrons were like
plums embedded in a positively charged pudding,
thus it was called the plum pudding
model. Based on the following facts (1) atoms
contain small, negatively charged particles
called electrons and (2) the atoms of the element
behave as if they have no charge at all
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Ernest Rutherford

• 1871-1937
• Learned physics in JJ Thomsons lab
• Did much work with alpha particles ( charged
  part with mass)
• Most famous for his GOLD FOIL EXPERIMENT

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Figure 3.5 Rutherfords experiment.
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Try it Yourself!
In the following pictures, there is a target
hidden by a cloud. To figure out the shape of the
target, we shot some beams into the cloud and
recorded where the beams came out. Can you figure
out the shape of the target?
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The Answers
Target 1
Target 2
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Figure 3.3 Plum Pudding model of an atom.
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Figure 3.6 Results of foil experiment if Plum
Pudding model had been correct.
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Figure 3.6 Actual results.
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Rutherfords Findings

• Most of the particles passed right through
• A few particles were deflected
• VERY FEW were greatly deflected

Like howitzer shells bouncing off of tissue
paper!
Conclusions

• The nucleus is small
• The nucleus is dense
• The nucleus is positively charged

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Disbelievers.

• Albert Einstein when to his grave not totally
  believing it
• According to classical physics, the electron
  would have collapsed into the nucleus
• 1910-1930 began the Quantum Physics Revolution
  (the physics of atomic and subatomic particles)

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The Atomic Scale

• Most of the mass of the atom is in the nucleus
  (protons and neutrons)
• Electrons are found outside of the nucleus (the
  electron cloud)
• Most of the volume of the atom is empty space

q is a particle called a quark
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The Quark
Oopswrong Quark!
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About Quarks
Protons and neutrons are NOT fundamental
particles.
Protons are made of two up quarks and one
down quark.
Neutrons are made of one up quark and two
down quarks.
Quarks are held together by gluons
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Figure 3.9 A nuclear atom viewed in cross
section.
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Atomic Number
Atomic number (Z) of an element is the number of
protons in the nucleus of each atom of that
element.
Element of protons Atomic (Z)
Carbon 6 6
Phosphorus 15 15
Gold 79 79
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Isotopes
Elements occur in nature as mixtures of isotopes.
Isotopes are atoms of the same element that
differ in the number of neutrons
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Figure 3.10 Two isotopes of sodium.
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Mass Number
Mass number is the number of protons and neutrons
in the nucleus of an isotope.
Mass p n0
Nuclide p n0 e- Mass
Oxygen - 10
- 33 42
- 31 15
8
8
18
18
Arsenic
75
33
75
Phosphorus
15
31
16
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Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon 12.011
Isotope Symbol Composition of the nucleus in nature
Carbon-12 12C 6 protons 6 neutrons 98.89
Carbon-13 13C 6 protons 7 neutrons 1.11
Carbon-14 14C 6 protons 8 neutrons lt0.01
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IsotopesAgain (must be on the test)
Isotopes are atoms of the same element having
different masses due to varying numbers of
neutrons.
Isotope Protons Electrons Neutrons Nucleus
Hydrogen1 (protium) 1 1 0
Hydrogen-2 (deuterium) 1 1 1
Hydrogen-3 (tritium) 1 1 2
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Chlorine Practice Problem

• Chlorine exists as 2 isotopes in nature. Cl-35
  (atomic mass 34.969 amu) has a 75.77 relative
  abundance. Cl-37 has an atomic mass 36.966 amu.
• What is the abundance of the Cl-37 isotope?

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Chlorine Practice Problem

• Chlorine exists as 2 isotopes in nature. Cl-35
  (atomic mass 34.969 amu) has a 75.77 relative
  abundance. Cl-37 has an atomic mass 36.966 amu
• Calculate the atomic mass of Chlorine.

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Chlorine Practice Problem

• Chlorine exists as 2 isotopes in nature. Cl-35
  (atomic mass 34.969 amu) has a 75.77 relative
  abundance. Cl-37 has an atomic mass 36.966 amu.
• How many times more massive is Cl-37 than Cl-35?

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Atomic and Atomic Mass

• The Periodic Law

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Chinese Periodic Table
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Triangular Periodic Table
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Mayan Periodic Table
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Stowe Periodic Table
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A Spiral Periodic Table
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The Year 1860.

• First International Congress of Chemists
• 60 to 70 of 113 elements had been discovered
• Italian chemist Cannizzaro presented method for
  measurement of atomic mass that all could agree
  on.

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John Newlands (1837-1898)

• English chemist that in 1864 proposed an
  organizational scheme for the elements
• Noticed properties repeated periodically every 8
  elements
• Called Law of Octaves
• Did not work for all elements
• Criticized analogy for being non-scientific

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Dimitri Mendeleev

• Writing a book about the same time.
• Wanted to include new information of atomic
  masses
• Wanted to find an arrangement for all of the
  information on the 60-70 elements
• Lothar Meyer (German chemist) did the same but
  Mendeelev published first

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Modern Russian Table
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Law of Mendeleev

• Properties of the elements recur in regular
  cycles (periodically) when the elements are
  arranged in order of increasing atomic mass.

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Missing?
14 Si 28.09
??
50 Sn 118.71
Named missing element Ekasilicon
From base word eka meanging next in order
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Ekasilicon
Predicted Properties Observed Properties
Atomic Mass 72 amu
Density 5.5 g/cm3
Melting Point 825 C
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Ekasilicon
Predicted Properties Observed Properties
Atomic Mass 72 amu 72.61 amu
Density 5.5 g/cm3 5.32 g/cm3
Melting Point 825 C 938 C
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Ekasilicon
Predicted Properties Observed Properties
Oxide Formula XO2 GeO2

Chloride Formula XCl4 GeCl4
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Mendeleevs Periodic Table
Dmitri Mendeleev
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Henry Moseley (1887-1915)

• After more elements were discovered, order of
  atomic mass wasnt working
• 1913 Discovered atoms have unique number or
  protons in nucleus and rearranged table in order
  of atomic number
• Periodic Law The statement that there is a
  periodic repetition of chemical and physical
  properties of elements when they are arranged by
  increasing atomic number.

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The Periodic Table
Period
Group or Family
Group or family
Period
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Figure 3.12 Elements classified as metals and
nonmetals.
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Properties of Metals

• Metals have luster (shiny) when smooth and clean
• Metals are solid _at_ room T
• Metals are good conductors of heat and
  electricity
• Metals are malleable
• Metals are ductile
• Metals have high tensile strength
• Located left of stairstep

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Figure 3.17 Spherical atoms packed closely
together.
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Metalloids

• Mostly brittle solids
• Properties between metal and non-metal
  (semi-conductors)
• Found in nature only as compounds
• Once obtained as free metals, are stable in the
  presence of air

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Properties of Nonmetals
Carbon, the graphite in pencil lead is a great
example of a nonmetallic element.

• Nonmetals are poor conductors of heat and
• electricity
• Nonmetals tend to be brittle and dull looking if
  solid
• Many nonmetals are gases at room T

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Examples of Nonmetals
Microspheres of phosphorus, P, a reactive
nonmetal
Sulfur, S, was once known as brimstone
Graphite is not the only pure form of carbon, C.
Diamond is also carbon the color comes from
impurities caught within the crystal structure
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Figure 3.14 Nitrogen gas contains N2 molecules.
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Figure 3.13 A collection of argon atoms.
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Figure 3.14 Oxygen gas contains O2 molecules.
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Table 3.5
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Valence e-

• The e- in an elements outermost orbital they
  determine the chemical properties of the element.
• Group 1 1 valence e-
• Group 2 2 valence e-
• Group 13 3 valence e-
• Group 14 4 valence e-
• Group 15 5 valence e-
• Group 16 6 valence e-
• Group 17 7 valence e-
• Group 18 8 valence e-

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Figure 3.19 The ions formed by selected members
of groups 1, 2, 3, 6, and 7.
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Atomic Radius

• Metal one half the distance between two
  adjacent atoms in a crystal
• Non-metal half the idstance between nuclei of
  identical atoms that are chemically bonded
  together.
• Trends
• Atomic radii generally decreases left to right
  across a period (more e-, same E level)
• Atomic radii generally increases down a group
  (added e- and added E level)

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Ionic Radius

• Ion An atom or a group of atoms that has a
  charge (due to loss or gain of e-)
• Cations ( charged ions)
• Radii become smaller due to loss of e- and
  decreased repulsion/often lose E level
• Anions (- charged ions)
• Radii become larger due to increase in of e-
  and repulsion among them (w/o change in proton )

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Table of Ion Sizes
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Octet Rule

• Atoms tend to gain, lose, or share e- in order to
  acquire a full set of 8 valence e-.

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Electronegativity

• The relative ability of an elements atoms to
  attract e- in a chemical bond.
• Defined in Paulings units
• F most electronegative element
• Fr least electronegative element
• Noble Gases none
• Trends
• E.N. increases across a period
• E.N. decreases down a group

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Ionization E

• The E required to remove an e- from a gasseous
  atom
• The lower the I.E, the more easily an atom given
  away its e-.
• Trends
• I.E. increases from left to right
• I.E. decreases from top to bottom

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Figure 3.11 The periodic table
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Figure 3.11 The periodic table
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Figure 3.20 Pure water does not conduct a
current.
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Figure 3.20 Water containing dissolved salt
conducts a current.
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Radioactivity

• Radioactivity a process in which some substance
  spontaneously emit radiation
• Radiation The rays and particles emitted by the
  radioactive materials.
• Nuclear Reaction A change in an atoms nucleus.
• Radioactive Decay emitted radiation in a
  spontaneous process
• Major Breakthrough 1890s.
• Only happens in radioactive atoms with unstable
  nuclei

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Types of Radiation

• Alpha Radiation
• Made up of alpha particles
• Deflected toward a negatively charged plate
• Beta Radiation
• Made up of beta particles
• Deflected toward a positively charged plate
• Gamma Radiation
• High E radiation that has no charge and mass
• Not deflected by electronic or magnetic fields

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Allotropes

• Different forms of a given element
• Different properties b/c different arrangement of
  atoms
• EX diamond, graphite, Buckminsterfullerene (All
  Carbon!)
• Look in your bookp. 70

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Periodic Table with Group Names
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Periodic Table with Group Names