Chem 400 Chapter 2

1
Chem 400 Chapter 2
2
Going Further The Structure of Atoms

• Dalton thought that atoms were the smallest
  particle of matter, but through a series of
  experiments starting in the late 1800s, this was
  proved to be incorrect.
• In a series of experiments by various scientists,
  the existence of electrons, protons, and neutrons
  were deduced and verified.

3
Going Further The Structure of Atoms

• Electrons were discovered in 1897 by JJ Thomson,
  and they were found to have a negative electric
  charge.
• Protons were hypothesized in 1911 by Ernest
  Rutherford, and were verified in 1919 by
  Rutherford. They have a positive electric charge.
• Neutrons were discovered in 1932 by James
  Chadwick. They have no electric charge.

4
Going Further The Structure of Atoms

• There are 2 important experiments that you should
  be aware of
• Robert Millikans Oil-Drop Experiment of 1909,
  which enabled him to measure the magnitude of the
  electric charge on electrons, and calculate their
  mass in grams.
• Ernest Rutherfords 1911 Gold-Foil Experiment,
  also called the Alpha-Scattering Experiment. This
  experiment enabled him to hypothesize the
  existence of protons in the nucleus of atoms.

5
Oil-Drop Experiment
6
(No Transcript)
7
The Structure of Atoms

• From this experiment, Millikan obtained the
  actual charge on an electron, -1.60x10-19 C.
• And from this charge and Thomsons charge/mass
  ratio, the exact mass of an electron was
  calculated to be 9.10x10-28 g.
• So from these experiments, scientists deduced
  that atoms were made up of even smaller subatomic
  particles, one of which was the electron.
• Since electrons have a negative charge, while
  atoms are neutral, scientists also realized that
  there had to be at least 1 more subatomic
  particle with a positive charge.

8
The Structure of Atoms

• Where was this atomic particle what did it look
  like?

9
Gold-Foil Experiment
10
The Structure of Atoms

• How could something in the atom cause such huge
  deflections in a massive positively charged
  particle like an alpha particle?
• So Rutherford proposed that atoms were composed
  of mostly empty space (where the electrons moved
  in circular orbits) with a very small, very
  massive, very dense center called the nucleus.
• The nucleus has a positive charge. This was
  Rutherfords Atomic Model.

11
The Structure of Atoms
12
The Structure of Atoms

• Rutherford then proved the existence of protons
  in 1919, and neutrons were discovered by James
  Chadwick in 1932.
• So whats the overall picture of an atom, and
  what are the sizes, masses, charges, and
  densities of the particles and regions?

13
The Structure of Atoms
Particle Mass (g x10-24) Mass (amu) Relative Mass Charge (C x10-19) Relative Charge Location in Atom
proton, p 1.673 1.0073 1 1.602 1 nucleus
electron, e or e- 0.000911 0.000549 0 -1.602 -1 electron cloud
neutron, n 1.675 1.0087 1 0 0 nucleus
14
The Structure of Atoms

• The diameter of a typical atom is around 1x10-10
  m or 1 Å.
• The diameter of a typical nucleus is only 0.0001
  Å.
• You can see that most of the mass of the atom is
  contained in a very small volume, so the nucleus
  is incredibly dense.
• The density of a typical nucleus is 1x1013 to
  1x1014 g/cm3, beyond our comprehension! If a
  matchbox had this density, it would weigh 2.5
  billion tons!

15
The Structure of Atoms
16
Atomic Number, Mass Number and Isotopes

• Dalton thought that atoms of different elements
  differed mainly by mass, but we now know that
  atoms of different elements differ by the number
  of protons which they contain.
• The number of protons which an element contains
  is called the Atomic Number, Z.
• The Atomic Number is found on the Periodic Table
  above the elemental symbol.

17
Atomic Number, Mass Number and Isotopes
18
Atomic Number, Mass Number and Isotopes

• It is true that every atom of the same element
  contains the same number of protons.
• So every H atom has 1 proton, and every C atom
  has 6 protons.
• So the number of protons defines the element.
• But it is NOT true that all atoms of the same
  element are identical.
• Whats different? Well, what else is there?

19
Atomic Number, Mass Number and Isotopes

• Although all atoms of the same element have the
  same number of protons, they DO NOT have the same
  number of neutrons!
• and if it is a neutral atom, they all have the
  same number of electrons
• Atoms of the same element which have different
  numbers of neutrons are called isotopes.

20
Atomic Number, Mass Number and Isotopes

• So isotopes of the same element differ by the
  number of neutrons.
• And since neutrons are the same relative mass as
  protons, isotopes also differ by mass.
• Although this is not shown on the Periodic Table,
  every element has at least 2 isotopes (except
  some of the newly synthesized elements like Mt).

21
Atomic Number, Mass Number and Isotopes

• To show different isotopes, we have several
  different isotopic notations or isotopic symbols.
• They all use the Mass Number, A, which is the sum
  of the protons and neutrons in the nucleus.
• For example, H has 3 common isotopes, H-1, H-2,
  and H-3.
• Carbon also has 3 common isotopes, C-12, C-13,
  and C-14.
• The number after the symbol or the superscript
  left number is the Mass Number.

22
Atomic Number, Mass Number and Isotopes
Mass Number
Mass Number
Atomic Number

• Practice with Isotopic Notation How many
  electrons, protons, and neutrons do the following
  isotopes have?

23
Ions and Ionic Isotopic Notation

• A neutral atom has equal numbers of protons and
  electrons. Why?
• During chemical reactions, atoms may gain or lose
  (or share) electrons, e-.
• If an atom gains or loses 1 or more e-, there is
  an imbalance between protons and e-, and the
  result is a charged particle called an ion.
• So an ion is formed when an atom gains or loses
  e-.

24
Ions and Ionic Isotopic Notation

• If an atom loses 1 or more e-, then it has more
  protons than e-, so the ion has a charge. It is
  called a cation.
• If an atom gains 1 or more e-, then it has less
  protons than e-, so the ion has a - charge. It is
  called an anion.
• Note that it is difficult to gain more than 3 e-
  or lose more than 4 e-.
• How do we show ions?

25
Ions and Ionic Isotopic Notation

• If we have an isotope which is an ion, we can
  show a complete isotopic notation for the ion.
• How many protons, e-, and neutrons do the
  following ionic isotopes have?

26
Elements and the Periodic Table

• Elements are fundamental substances.
• They cant be broken down into smaller substances
  by chemical reactions.
• The Periodic Table arranges the known elements
  (114 of them).
• 90 of these are naturally occurring, while the
  rest have been synthesized in nuclear reactions.

27
Elements and the Periodic Table
28
Elements and the Periodic Table

• Notice that the elements names have been given
  shorthand notations (called symbols) of 1 or 2
  letters.
• Unnamed elements actually have a 3 letter
  designation until they are named.
• The first letter is ALWAYS capitalized, while the
  second letter is ALWAYS lowercase.
• What elements do you have to memorize (names and
  symbols)? 1-40 42 46-57 76-90 92 and 94.

29
Elements and the Periodic Table

• Although most of the symbols are obviously
  related to the name, like N for nitrogen, others
  seem to make no sense, like Pb for lead!
• This is because some of the symbols come from
  old Latin names or other languages.
• Plumbum was an old Latin name for lead.
• W for tungsten comes from the German name wolfram.

30
Elements and the Periodic Table

• Chemistry in some fashion has been around for
  centuries.
• Some elements were known thousands of years ago.
• But most elements were discovered and identified
  in the last 250 years.

31
Elements and the Periodic Table

• In the early to mid 1800s, chemists were trying
  to organize the 60-some known elements into some
  sort of pattern.
• Mendeleev designed a Periodic Table in 1869 which
  was based on the masses of the known elements
  (atomic weights) and the compounds they formed
  with hydrogen (hydrides) or oxygen (oxides).
• Todays Table is similar, but the elements are
  arranged by atomic numbers (number of protons)
  instead of by atomic weights.

32
Elements and the Periodic Table

• If you look at a Periodic Table, there are 18
  columns called Groups or Families.
• They are called families as they share common
  chemical properties or characteristics.
• The 7 rows are called Periods.
• The groups are numbered 2 ways on US Tables.
• The old US system uses numbers with A or B
  sections, while the internationally approved
  system simply numbers the groups from 1 to 18
  going across from left to right.

33
Elements and the Periodic Table

• There are several basic regions on the Table
• Metals
• Nonmetals
• Semimetals (metalloids or semiconductors)
• Main Block or Representative Elements
• Transition Metals
• Inner Transition Metals
• Lanthanides
• Actinides

34
Elements and the Periodic Table

• Several important Groups also have names
• Group 1, except hydrogen, are the Alkali Metals
• Group 2 is the Alkaline Earth Metals
• Group 17 are the Halogens
• Group 18 are the Noble Gases

35
Elements and the Periodic Table

• Metals lustrous, silvery, malleable, ductile,
  generally hard, solids except Hg, conductors,
  lose electrons to become cations, react with
  nonmetals to form ionic salts.
• Nonmetals nonconductors, react with metals, gain
  electrons to form anions, brittle, most gases (1
  l, 5 s)
• Metalloids B, Si, As, Te, At, Ge, Sb in between
  metals nonmetals, semiconductors, solids

36
Elements and the Periodic Table

• Here are some shared characteristics in the
  regions and groups
• Alkali Metals Very reactive metals, soft, not
  found in nature as pure element
• Alkaline Earth Metals same as Alkali metals but
  less reactive
• Halogens most reactive nonmetals, corrosive, not
  found in nature as pure element

37
Elements and the Periodic Table

• Noble Gases also called Inert gases as very
  nonreactive, dont form compounds except Xe
• Lanthanides f-fillers, rare earth metals, inner
  transition metals, reactive, silvery-grey
• Actinides f-fillers, rare earth metals, inner
  transition metals, reactive, silvery-grey,
  radioactive, synthetic above 92
• Why is hydrogen placed in Group 1 if it is NOT an
  Alkali Metal and is actually a nonmetal?

38
Elements and the Periodic Table
39
Elements and the Periodic Table
40
Atomic Mass and Weighted Averages of Elements

• As atoms have a very tiny mass in grams,
  scientists use another scale to state the masses
  of atoms, the atomic mass unit, amu. You see
  this in the table with the masses of p, e, and n
  given earlier.
• The conversion factor between mass in g and mass
  in amu is
• 1 amu 1.66054x10-24g OR
• 1 g 6.02214x1023amu

41
Atomic Mass and Weighted Averages of Elements

• The average atomic mass of the elements is shown
  beneath the elemental symbol on the Periodic
  Table.

42
Atomic Mass and Weighted Averages of Elements

• But every element has different isotopes with
  different masses!
• Thats why the atomic masses on the Table are
  average masses it is really the mass in amu of a
  single average atom of an element.
• But what does an average atom of an element
  look like?
• What does the average student look like? Does
  it exist?

43
Atomic Mass and Weighted Averages of Elements

• For H, 99.985 of all H atoms are H-1, while
  0.015 are H-2 (there are basically 0 H-3).
  This is called the natural abundance or
  -abundance of an isotope.
• So shouldnt the average H atom look a lot like
  H-1, and shouldnt the average atomic mass of H
  be very close to the mass of the H-1 isotope?
• Because the different isotopes do not have equal
  natural abundances, we calculate atomic masses of
  elements using a weighted average of all the
  isotopes

44
Atomic Mass and Weighted Averages of Elements

• So if lead has 4 common isotopes with the
  following masses and -abundance, what is the
  atomic mass of lead?
• Pb-204 203.973020 amu 1.40
• Pb-206 205.974440 amu 24.1
• Pb-207 206.975872 amu 22.1
• Pb-208 207.976627 amu 52.4

45
Molar Mass of Atoms Avogadros Number

• The atomic mass is the mass in amu of a single
  average atom.
• Is this useful in the lab? Can we pick out and
  weigh an individual atom?
• Chemists weigh in g, which is a HUGE number of
  atoms.
• So we need a unit to express large numbers of
  atoms or molecules without using scientific
  notation.
• Chemists defined a counting unit to do this.

46
Molar Mass of Atoms Avogadros Number

• They chose a unit called mole so that the atomic
  mass on the Periodic Table is also used for
  measuring grams.
• 1 mole 1 mol 6.02214x1023 things
• So if you have 1 mol of pennies, how many pennies
  do you have? How many dollars is this?
• The number 6.02214x1023 is Avogadros number.

47
Molar Mass of Atoms Avogadros Number

• Avogadros number is very important as it is a
  conversion factor between number of things and
  moles of things.
• EX if you have 2.5 mol of aluminum, how many
  atoms of aluminum do you have?
• Avogadros number is also very special as if we
  have 1 mol of an element, we have the atomic mass
  in g of the element.

48
Molar Mass of Atoms Avogadros Number

• If we have 1 mol of Na atoms, prove that this is
  22.99 g of Na

49
Molar Mass of Atoms Avogadros Number

• The molar mass is the mass in g of exactly 1 mol
  of an element.
• The atomic mass and the molar mass are the same
  number, they differ only by units!

50
Molar Mass of Atoms Avogadros Number

• The molar mass is also a conversion factor it
  converts between mol of an element and g of an
  element.
• Ex If you have 25.7840 g of gold, how many mol
  of gold is this?
• There are 6 types of simple calculations that we
  can do using molar masses and Avogadros number.