Chapter 4: Elements, Atoms and the Periodic Table

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• Chapter 4 Elements, Atoms and the Periodic Table

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Elements

• Periodic Table lists 115 known elements
• Elements are being added every year
• Element- Cannot be broken into simpler
  substances.
• Compound- Combination of two or more elements.
• There are 92 naturally occurring elements.
• 23 elements have been synthesized.
• Each element has its own unique symbol.
• A symbol is a combination of 1 or 2 letters.

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Periodic Table of the Elements
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Elemental Abundances

• Hydrogen, H, is the most abundant element in the
  universe.
• Oxygen, O, is the most abundant element in the
  earths crust.
• Silicon, Si, is also abundant in the earths
  crust.

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Abundance of Elements in the Universe
Insert figure 4.3
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Abundance of Elements in Earths Crust
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Classification of Elements

• Metals
• Located to the left and below the diagonal line
  in the periodic table.
• High Luster, combine with non-metals, conductive
• Non-metals
• Located to the right and above the diagonal line
  in the periodic table.
• Combine with one another and metals.
• Non-conductive

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Metalloids

• Lie along the diagonal line in the periodic
  table.
• Properties intermediate between metals and
  non-metals. (semiconductors)

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Classification of Elements
Insert figure 4.6
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Classification of Elements

• Periods Horizontal rows of elements
• Groups Vertical columns of elements
• Elements in the same group have similar
  properties
• Group IA. Alkali metals have similar properties
• Group VIIA Halogens diatomic (2 atoms combined by
  a chemical bond in an element.
• Hydrogen, H2, oxygen, O2, nitrogen, N2 also form
  diatomic molecules.

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Physical State at Room Temperature (20o C)

• Depends on melting point and boiling point
• Gases
• Noble Gases He, Ne, Ar, Kr, Xe, Rn
• Other Gases H2, N2, O2, F2, Cl2
• Liquids
• Br2, Hg
• Solids- All other Elements
• Metals like Na, Ag, Fe, Au, Pb, K, Ca, Cu, Li,
  Sn, Cr
• Non-metals like B, C, P, S, I2,
• Metalloids like Si, As, Sb, Te

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Physical Properties

• Conductivity
• Ease of transmitting heat or electricity
• Metals are good conductors
• Non-metals are poor conductors
• Metalloids are semiconductors used in electronic
  circuits.
• Luster-ability to reflect light.
• Malleable-can be rolled and hammered into shape.
• Ductile-can be drawn into wires.
• Metals have luster, many are malleable and
  ductile.
• Non-metals are dull and brittle.

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Diatomic Elements
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Greek Philosophers

• Aristotle relied on logic--continuous matter
• Leucippus Democritus atomos-individual
  particles of matter which could not be
  subdivided. Substance were mixtures of different
  types of atoms.
• Lucretius Roman poet who wrote about atoms.

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Atoms

• The Greeks based their models on logic and
  speculation, not experiment.
• Several observations led up to the formulation of
  atomic theory by Dalton. Daltons model of the
  atom was based on observation and experiment, not
  speculation.
• Based on Laws of Nature as discovered by several
  people.

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Four pieces of evidence upon which the idea of
the atom is based.

• Law of Conservation of Mass
• Law of Definite Proportions
• Law of Multiple Proportions
• Law of Combining Volumes

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Law of Conservation of Mass

• Lavoisier
• Matter can neither be created nor destroyed
  during a chemical change.

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Law of Definite Proportions

• A compound always contains elements in certain
  definite proportions, never in any other
  combination also called the law of constant
  composition.
• Proust- by decomposition
• Berzelius- by synthesis

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Proust by Decomposition

• Proust found that when he decomposed copper (II)
  carbonate into its constituent elements the ratio
  of the masses of the elements was always the
  same, regardless of where the copper (II)
  carbonate came from
• If any 123.5 g sample of copper (II) carbonate
  was decomposed into its elements, there would be
  63.5 g Cu, 12 g C, and 48 g O.

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Composition of Copper (II) Carbonate

• Substance mass mass
• (g) fraction mass
• CuCO3 123.5
• Cu 63.5 63.5/123.5
• 0.514 51.4
• C 12 12/123.5
• 0.097 9.7
• O 48 48/123.5
• 0.389 38.9
• 100

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Using Mass Fraction Data to Determine Masses of
Elements From Other Quantities of CuCO3

• Substance mass fraction mass
• CuCO3 1 103
• Cu 0.514 x103 53
• C 0.097 x103 10
• O 0.389 x103 40
• 
  103

5.3 parts C 4 parts C 1 part C
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Berzelius by Synthesis

• Berzelius found that when he combined 828.8 g of
  Pb with 128.4g of S, he formed 957.2 g PbS.
• But if he tried to combine 828.8 g of Pb with
  192.6 g of S he formed 957.2 g PbS with 64.2g of
  S left over.
• And if he tried to combine 1243.2 g Pb with 128.4
  g S, he formed 957.2 g PbS with 414.4 g Pb left
  over.

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Law of Definite Proportions by Synthesis
(Berzelius)
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Law of Multiple Proportions

• Dalton
• Elements may combine in more than one proportion
  to form more than one compound (CO and CO2).
• For a fixed mass of one element in both
  compounds, the ratio of the masses of the other
  element in the two compounds is as small whole
  numbers.

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Compare the Masses in CO and CO2

• Compound C O
• CO 42.8 57.1
• CO2 27.3 72.7
• 28 g of CO contains 28 x 0.428 12 g C
• 28 x 0.571 16 g O
• 44 g of CO2 contains 44 x 0.273 12 g C
• 44 x 0.727 32 g O
• For 12 g C, the ratio of O is 1632 12

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Law of Combining Volumes

• The volumes of gaseous reactants and products are
  in a small whole number ratio when all
  measurements are made at the same temperature and
  pressure.
• Cavendish-burned hydrogen and oxygen to form
  water
• Volta-battery
• Nicholson and Carlisle-used electrolysis to
  decompose water into hydrogen oxygen.

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Daltons atomic theory

• All matter is composed of extremely small
  particles called atoms.
• All atoms of a given element are alike, but atoms
  of one element differ from the atoms of any other
  element.
• Compounds are formed when atoms of different
  elements combine in fixed proportion.

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Hypotheses (cont.)

• A chemical reaction involves a rearrangement of
  atoms. No atoms are created or destroyed or
  broken apart in a chemical reaction.

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Daltons Atomic Theory Explains the Experimental
Laws

• The Laws of nature can be described graphically
• Law of Conservation of Mass
• Law of Definite Proportions
• Law of Multiple Proportions

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Atomic Representation of the Law of Definite
Proportions

• By Decomposition (Proust)
• O
• Cu C O
• O

• By Decomposition (Proust)
• O
• Cu C O
• O

O
16
16
Cu
C
12
16
63.5
63.5
12
O
63.5 12 16 16 16
63.5 12 16 16 16
16
16
O
16
Total mass 123.5
Total mass 123.5
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Atomic Representation of the Law of Definite
Proportions

• By Synthesis (Berzelius)

s
207.232.1239.3
Pb
s
Pb
239.3
s
Pb
s
Pb
239.3
s
s
Pb
Pb
4x32.1128.4
4x207.2 828.8
s
s
Pb
239.3
Pb
Total 828.8128.4957.2
957.2
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Atomic Representation of the Law of Definite
Proportions

• By Synthesis (Berzelius) (Continued)

s
207.232.1239.3
Pb
s
s
Pb
239.3
s
s
Pb
s
Pb
239.3
s
s
s
s
Pb
Pb
32.1x2
4x207.2828.8
239.3
s
s
957.2
64.2
Pb
4x32.1128.4
Pb
(left over) (PbS)
64.2192.6
1021.2
Total 828.8128.4957.264.21021.4
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Law of Definite Proportions Illustrated with
Atoms
Insert figure 4.18
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Atomic Representation of the Law of Multiple
Proportions
Insert figure 4.19
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Subatomic Particles

• Electrolysis (Faraday, Nicholson and Carlisle)
  showed that matter was electrical.
• Atoms contain small, sub-atomic particles.
• Electron, proton, neutron
• Two types of charges , -
• Opposite charges attract, Like charges repel
• Protons electrons - neutrons no charge

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Masses of Atoms and Particles

• Dalton compared the masses of elements, setting
  hydrogen equal to 1.
• An atom having a mass 12 times H would have a
  mass of 12.
• Now we compare masses in terms of atomic mass
  units (amu)
• 1 amu 1.6606 x 10-24 g

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Subatomic Particles
Particle Symbol Charge Mass (amu) (approxima
te) Electron e- -1 0.000544 Proton p 1 1 Neu
tron n 0 1
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Atoms

• Atoms are extremely small. They have been
  photographed with electron microscopes, scanning
  tunneling microscopes, and atomic force
  microscopes.
• The atoms have a diameters of 0.1-1.0 nm
• The nucleus (where the protons and neutrons are)
  has a diameter
• of about 10-15 m.

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Insert figure 4.23
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Atomic NumberMass Number

• Atomic Number identifies the atom of an element.
  Equals the number of protons in the nucleus.
• Mass Number equals the number of protons plus
  neutrons in the nucleus. Approximately equal to
  the mass of the atom in amu.

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Define the term isotope

• Elements which have atoms which have the same
  number of protons, but different numbers of
  neutrons.
• Dalton was wrong!
• Different isotopes of the same element have the
  same number of protons, but different numbers of
  neutrons (and therefore different masses).

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Three isotopes of hydrogen.

• Protium (the most abundant form) contains 1
  proton and 1 electron. (mass 1 u)
• Deuterium contains 1 proton, 1 neutron, and 1
  electron. (mass 2 u)
• Tritium (the radioactive isotope) contains 1
  proton, 2 neutrons, and 1 electron. (mass
  3 u)

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Insert figure 4.24
Isotopes of Hydrogen
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A and Z

• Z is the atomic number, the number of protonsthe
  number of electrons in a neutral atom. The
  charge on the nucleus.
• A is the mass number or nucleon number, the
  number of protons neutrons in the atom. The
  approximate mass of the atom.

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Parts of the nuclear symbol related to number of
protons and neutrons, atomic mass, atomic number.

• A nuclear symbol has the following form
• AZX
• Where
• A is the mass number (N P)
• Z is the atomic number (P)

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Isotope Symbols

• In the symbol 146C
• The atomic number, Z is 6
• The mass number, A is 14
• The number of protons is 6.
• The number of neutrons can be calculated from the
  formula A N P N A - P
• N 14 - 6 8 neutrons.

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Isotope Symbols

• Give the isotope symbol for the isotope which
  contains 22 protons and 28 neutrons.
• Z P 22
• A N P 22 28 50
• Element 22, from the periodic table is Titanium,
  Ti
• 5022Ti

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Atomic Mass

• Determine atomic mass, given isotope masses and
  abundances.
• To determine atomic mass, multiply each isotope
  mass as determined with a mass spectrometer by
  the isotopic abundance.
• Then add these products together.
• The sum is the atomic mass of the element.

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Isotopic Masses
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Example Problem

• Determine the atomic mass of chlorine, which
  consists of two naturally occurring isotopes
• Isotope mass (amu) abundance ()
• 3517Cl 34.96885 75.53
• 3717Cl 36.97790 24.47
• 34.96885x0.7554 26.415
• 36.97790x0.2447 9.048
• 35.46 atomic mass

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Atomic Mass

• Define atomic mass and relate it to carbon-12.
• Atomic mass of an isotope is measured with a mass
  spectrometer in which an atom is ionized and then
  accelerated toward a negative electrode.
• A magnet bends the beam of ions. The greater the
  bending, the less mass. The mass is compared
  with 126C being assigned a mass of exactly
  12.000000000000...

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The Atomic Mass of an Element

• The atomic mass of an element is the average of
  the masses of the naturally occurring isotopes
  weighted according to their abundance.
• Atomic mass is measured in atomic mass units
  (amu).
• 1amu 1.66054x10-24g
• One 126C atom has a mass of exactly
• 12.0 amu or 12x1.66054x10-24g/amu
• 1.99265x10-23g

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Avogadros Number

• Relate molar mass to atomic mass, Avogadros
  Number, and the mass of one atom.
• The molar mass is the mass of one mole of an
  element. One mole has 6.022 x 1023 atoms.
• The molar mass is also the atomic mass measured
  in grams. Thus, the atomic mass of Na is 22.99
  the molar mass is 22.99 g.

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Example Problems

• Calculate the number of moles present in 9.9x1022
  atoms of Na.
• Mole atoms/6.022x1023 atoms/mole
• 9.9x1022 atoms/6.022x1023 atoms/mol

1.64x10-1 0.16 mol
Calculate the number of atoms of Na
present in 4.00 mol.
Atomsmolx6.022x10234.00molx6.022x1023atom/mol
2.41x1024atoms
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Molar Mass of Monatomic Elements

• One atom of Na has a mass of 22.99 amu.
• One mole of Na has a mass of 22.99 g,
• or (22.99g/mol)/(6.022x1023atoms/mole
• 3.818x10-23g/atommass of one Na atom

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Number of particles and moles

• Relate the actual number of atoms or molecules to
  the number of moles.
• One mole is 6.022x1023 entities (atoms,
  molecules, ions, etc.)
• entities6.022x1023 x moles
• moles(entities)/6.022x1023

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Definitions

• Relate the terms molecular mass, formula mass,
  formula weight, molecular weight, molar mass.
• All of these terms mean about the same thing.
  Numerically, they are all equivalent
• When measured in grams or amu, the more precise
  term is mass. Technically molecular refers only
  to molecules, formula can apply to ionic
  compounds, ions, atoms.
• Molar mass is the atomic mass, molecular mass, or
  formula mass measured in grams

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Determining Formula Mass

• (Also called molecular weight. When expressed in
  g, this is the molar mass, M)
• The formula mass is the sum of the atomic masses
  of the atoms in the element times their
  subscripts.

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For the Compound with FormulaAaBb

• M (AaBb) aA bB, where A and B are the atomic
  masses of A and B, respectively.
• Calculate the formula mass and molar mass of
  C6H12O6.
• M 6(12.0) 12(1.0) 6(16.0) 180
• 180 g (molar mass)

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Molar Mass, Mass, Moles

• Given two of the following, determine the third
  molar mass, mass, number of moles.
• Formulas
• Mol g/M
• g (mol)(M)
• M g/mol

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Example Problems

• Calculate the number of moles present in 50.0g of
  glucose (C6H12O6 , M 180 g/mol)
• mol g/M 50.0g/180g/mol

.278 mol
Calculate the mass of 0.233 mol of glucose.
g mol(M) 0.233 molx180g/mol
41.9 g
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Calculate Formula Mass

• Calculate the formula mass of an unknown compound
  where 2.78 moles of the substance has a mass of
  500g.
• M g/mol 500g/2.78mol
• 180 gmol-1
• Formula mass 180