Atomic Structure

1
Chapter 4

• Atomic Structure

2
Defining the atom

• This simple model goes back to the Greek
  philosopher Democritus (460 370 BC).
• Democritus reasoned that is a piece of matter,
  such as gold, were divided into smaller and
  smaller pieces, one would ultimately arrive at a
  tiny particle of gold that could not be divided
  further but that still retains the properties of
  gold.
• He used the word atom which literally means
  indivisible, to describe this undividable,
  ultimate particle of matter.

3
Defining the atom

• Plato (428 348 BC) and Aristotle (384 322 BC)
  argued against the existence of atoms, and their
  ideas prevailed for centuries.
• They believed that all matter was continuous.
• Most of those in the mainstream of enlightened
  thought rejected or remained ignorant of the
  atomic theory proposed by Democritus, though a
  few well-known scientists did refer to atoms.

4
Defining the atom

• Galileo Galilei (1564 1642) reasoned that the
  appearance of a new substance through chemical
  change involved a rearrangement of parts too
  small to be seen.
• Francis Bacon (1561 1626) speculated that heat
  might be a form of motion of small particles.
• Robert Boyle (1627 1691) was guided in his work
  on gases and other aspects of chemistry by what
  he called his corpuscular philosophy.
• Claims that all physical phenomena can be
  explained as the motions of different shapes,
  sizes, and extent of matter.
• None of these people, however, provided detailed,
  quantitative explanations of physical and
  chemical facts in terms of atomic theory.

5
Defining the atom

• Although Democrituss ideas agrees with
  scientific theory, they did not explain chemical
  behavior.
• They also lacked experimental support because
  Democrituss approach was not based on the
  scientific method.

6
Defining the atom

• The real nature of atoms and the connection
  between observable changes and events at the
  atomic level were not established for more than
  2000 years after Democritus.
• In 1803, John Dalton (1766 1844) used
  experimental methods to transform Democrituss
  ideas on atoms into a scientific theory.
• Dalton linked the existence of elements, which
  cannot be decomposed chemically, to the idea of
  atoms, which are indivisible.

7
Defining the atom

• Compounds, which can be broken down into two or
  more new substances, must contain two or more
  different kinds of atoms.
• Dalton went further to say that each kind of atom
  must have its own properties especially a
  characteristic mass.
• This idea allowed his theory to account
  quantitatively for the masses of different
  elements that combine chemically to form
  compounds.
• Thus, Daltons ideas could be used to interpret
  known chemical facts, and do so quantitatively.

8
Defining the atom

• The result was Daltons Atomic Theory
• All elements are composed of tiny indivisible
  particles called atoms.
• Atoms of the same element are identical. The
  atoms of any one element are different from those
  of any other element.
• Atoms of different elements can physically mix
  together or can chemically combine in simple
  whole-number ratios to form compounds.
• Chemical reactions occur when atoms are
  separated, joined, or rearranged. Atoms of one
  element, however, are never changed into atoms of
  another element as a result of a chemical
  reaction.

9
Who was the schoolmaster who studied chemistry
and proposed an atomic theory?

• John Dalton
• Jons Berzelius
• Robert Brown
• Dmitri Mendeleev

10
According to Daltons atomic theory, atoms

• Are destroyed in chemical reactions.
• Can be divided.
• Of each element are identical in size, mass, and
  other properties.
• Of different elements cannot combine.

11
Which of the following is not part of Daltons
atomic theory?

• Atoms cannot be divided, created, or destroyed.
• The number of protons in an atom is its atomic
  number.
• In chemical reactions, atoms are combined,
  separated, or rearranged.
• All matter is composed of extremely small
  particles called atoms.

12
According to Daltons atomic theory, atoms

• Of different elements combine in simple
  whole-number ratios to form compounds.
• Can be divided into protons, neutrons, and
  electrons.
• Of all elements are identical in size and mass.
• Can be destroyed in chemical reactions.

13
Which of the following statements is true?

• Atoms of the same element may have different
  masses.
• Atoms may be divided in ordinary chemical
  reactions.
• Atoms can never combine with any other atoms.
• Matter is composed of large particles called
  atoms.

14
Defining the atom

• John Daltons ideas were accepted by the
  scientific community because they helped early
  scientists understand two important laws
• The Law of the Conservation of Matter.
• The Law of Constant Composition. (Law of
  Definite Proportions)

15
Defining the atom

• The law of conservation of mass had been
  formulated a few years earlier by Antoine
  Lavoisier (1743 1794).
• He carefully determined the masses of reactants
  and products in a series of chemical reactions
  and found the sum of the masses of the reactants
  always equaled the sum of the masses of the
  products.
• This result can be understood from Daltons
  second and fourth parts of the atomic theory.

16
Defining the atom

• The French chemist Joseph Louis Proust (1754
  1826) formulated the law of constant composition
  (aka law of definite proportions) from his simple
  analysis of minerals.
• He determined that the ratio of masses of the
  elements in pure samples of a given compound did
  not vary, regardless of the samples origin.
• Daltons second and third parts rationalize this
  result, which must arise if the ratio of
  different atoms in a compound is fixed and if
  each kind of atom has a characteristic mass.

17
Defining the atom

• Example
• One of the compounds formed from carbon and
  oxygen always has a one-third greater mass of
  oxygen than carbon.
• This would be true if this compound contained one
  carbon atom for each oxygen atom giving the
  formula CO and if the mass of one oxygen atom
  was one-third greater than the mass of a carbon
  atom.
• Using this kind of reasoning, Dalton was able to
  determine ultimately which atoms were heavier
  than others and how much of one element could be
  expected to combine with another element in a
  compound.

18
Defining the atom

• Daltons theory was valuable because it explained
  existing facts, but he went further and proposed
  a new law, the law of multiple proportions.
• This states that if two elements form two
  different compounds, the mass ratio of the
  elements making up one compound is a whole number
  multiple of the mass ratio of the elements in the
  second compound.

19
Defining the atom

• Example
• Consider the two oxides of carbon Carbon
  monoxide (CO) and carbon dioxide (CO2).
• By careful laboratory measurements, it is
  possible to show that the ratio of the mass of
  oxygen to the mass of carbon in CO is 1.3333
  (mass of O/mass of C 1.3333).
• The ratio of the masses of O to C in CO2 is
  2.6666 (mass of O/mass of C 2.6666).
• The mass ratio for CO2 is exactly twice the value
  found for CO.

20
Defining the atom

• Daltons atomic theory was important because it
  suggested a new law and stimulated Dalton and his
  contemporaries to do a great deal more research,
  thereby contributing to scientific progress.
• A good theory not only accounts for existing
  knowledge but also stimulates the search for new
  knowledge.
• Although it was not until the 1860s that a
  consistent set of relative masses of the atoms
  was agreed on, Daltons idea that the masses of
  atoms are crucial to quantitative chemistry was
  accepted from the early 1800s on.

21
Defining the atom

• Despite their small size, individual atoms are
  observable with instruments such as scanning
  tunneling microscopes.
• The radii of most atoms fall within the range of
  5 10-11 m to 2 10-10 m.
• A scanning tunneling microscope (STM) holds a
  one-atom-sized tip over a sample that conducts
  electricity.
• Electric current flows between the tip and the
  sample.
• The current is stronger when the tip and the
  sample are close and weaker when they are farther
  apart.
• This change in voltage is interpreted to show
  surface topography.

22
Defining the atom

• With this new technology, scientists can move
  around atoms and arrange them into patterns.
• This can be used to create atomic-sized
  electronic devices, such as circuits and computer
  chips.
• This nanoscale technology could become
  essential to future applications in medicine,
  communications, solar energy, and space
  exploration.

23
The law of conservation of mass follows the
concept that

• Atoms are indivisible.
• Atoms of different elements have different
  properties.
• Matter is composed of atoms.
• Atoms can be destroyed in chemical reactions.

24
If each atom of element D has 3 mass units and
each atom of element E has 5 mass units, a
chemical molecule composed of one atom each of D
and E has

• 15 mass units.
• 2 mass units.
• 35 mass units.
• 8 mass units.

25
In oxides of nitrogen, such as N2O, NO, NO2, and
N2O3, atoms combine in small whole-number ratios.
This evidence supports the law of

• Conservation of mass.
• Multiple proportions.
• Definite composition.
• Mass action.

26
The law of multiple proportions can be partly
explained by the idea that

• Elements can combine in only one way to form
  compounds.
• Whole atoms of the same elements combine to form
  compounds.
• Elements in a compound always occur in a 11
  ratio.
• Only atoms of the same element can combine.

27
According to the law of definite proportions, any
two samples of KCl have

• The same mass.
• Slightly different molecular structures.
• The same melting point.
• The same ratio of elements.

28
Daltons atomic theory did not explain the law of

• Whole-number ratios.
• Definite proportions.
• Conservation of mass.
• Conservation of energy.

29
The law of definite proportions

• Contradicted Daltons atomic theory.
• Was explained by Daltons atomic theory.
• Replaced the law of conservation of mass.
• Assumes that atoms of all elements are identical.

30
Structure of the nuclear atom

• Electricity is involved in many of the
  experiments from which the theory of atomic
  structure was derived.
• The fact that objects can bear an electric charge
  was first observed by the ancient Egyptians, who
  noticed that amber, when rubbed with wool or
  silk, attracted small objects.
• You can observe the same thing when you comb your
  hair on a dry day your hair is attracted to the
  comb.
• A bolt of lightning or the shock you get when
  touching a doorknob results when an electric
  charge moves from one place to another.

31
Structure of the nuclear atom

• Two types of electric charge had been discovered
  by the time of Benjamin Franklin (1706 1790).
• He named them positive () and negative (-),
  because they appear as opposites and can
  neutralize each other.
• Experiments show that like charges repel each
  other and unlike charges attract each other.
• Franklin also concluded that charge is balanced
  if a negative charge appears somewhere, a
  positive charge of the same size must appear
  somewhere else.
• By the 19th century it was understood that
  positive and negative charges are somehow
  associated with matter perhaps in atoms.

32
Structure of the nuclear atom

• In 1896, the French physicist, Henri Becquerel
  (1852 1908), discovered that a uranium ore
  emitted rays that could darken a photographic
  plate, even though the plate was covered by black
  paper to protect it from being exposed by light
  rays.
• In 1898, Marie Curie (1867 1934) and coworkers
  isolated polonium and radium, which also emitted
  the same kind of rays, and in 1899 she suggested
  that atoms of certain substances disintegrate
  when they emit these unusual rays.
• She named this phenomenon radioactivity, and
  substances that display this property are said to
  be radioactive.

33
Structure of the nuclear atom

• Early experiments identified three kinds of
  radiation
• Alpha (?)
• Beta (?)
• Gamma (?)
• The behave differently when passed between
  electrically charged plates.
• Alpha and beta rays are deflected, but gamma rays
  pass straight through.
• This implies that alpha and beta rays are
  electrically charged particles, because charges
  are attracted or repelled by the charged plates.

34
Structure of the nuclear atom

• Even though an alpha particle was found to have
  an electric charge (2) twice as large as that of
  a beta particle (-1), alpha particles are
  deflected less, which implies that alpha
  particles must be heavier than beta particles.
• Gamma rays have no detectable charge or mass
  they behave like light rays.

35
Structure of the nuclear atom

• Marie Curies suggestion that atoms disintegrate
  contradicts Daltons idea that atoms are
  indivisible, and requires an extension of
  Daltons theory.
• If atoms can break apart, there must be something
  smaller than an atom that is, atomic structure
  must involve subatomic particles.

36
Structure of the nuclear atom

• Passing an electric current through a solution of
  a compound a technique called electrolysis
  can cause a chemical reaction, such as platting
  gold or silver onto another metal or the
  protection of chlorine from sodium chloride.
• In 1833, the British scientists Michael Faraday
  (1791 1867) showed that the same quantity of
  electric current caused different quantities of
  different metals to deposit, and postulated that
  those quantities were related to the relative
  masses of the atoms of those elements.
• Such experiments were interpreted to mean that,
  just as the atom is the fundamental particle of
  an element, a fundamental particle of electricity
  must exist.
• This particle of electricity was given the name
  electron.

37
Structure of the nuclear atom

• Further evidence that atoms are composed of
  smaller particles came from experiments with
  cathode ray tubes.
• These are glass tubes from which most of the air
  has been removed and that have a piece of metal
  called an electrode sealed into each end.
• When a sufficiently high voltage is applied to
  the electrodes, a cathode ray flows from the
  negative electrode (cathode) to the positive end
  (anode).
• Experiments showed that cathode rays travel in
  straight lines, cause gases to glow, can heat
  metal objects red hot, can be deflected by a
  magnetic field, and are attracted toward
  positively charged plates.
• When cathode rays strike a fluorescent screen,
  light is given off in a series of tiny flashes.
• Examples of Cathode Rays Televisions, computer
  monitors.

38
Structure of the nuclear atom
39
Structure of the nuclear atom

• This principle was first exploited by Sir Joseph
  John Thomson (1856 1940) to prove
  experimentally the existence of the electron.
• In his experiments he applied electric and
  magnetic field simultaneously to a beam of
  cathode rays.
• By balancing the effect of the electric field
  against that of the magnetic field and using
  basic laws of electricity and magnetism, Thomson
  calculated the ratio of charge to mass for the
  particles in the beam.
• He was not able to determine either charge or
  mass independently.

40
Structure of the nuclear atom

• Because cathode rays have a negative charge,
  Thomson suggested the particles were the same as
  the electrons associated with Faradays
  experiments.
• He obtained the same charge-to-mass ratio in
  experiments using 20 different metals as cathodes
  and several different gases.
• These results suggested that electrons are
  present in all kinds of matter and that they
  presumably exist in atoms of all elements.

41
Structure of the nuclear atom
42
Structure of the nuclear atom

• It remained for American physicist Robert Andrews
  Millikan (1868 1953) to measure the charge of
  an electron and thereby enable scientists to
  calculate its mass.
• In his apparatus, tiny drops of oil were sprayed
  into a chamber. As they settled slowly through
  the air, the droplets were exposed to x-rays,
  which caused air molecules to transfer electrons
  to them.
• Millikan used a small telescope to observe
  individual droplets.
• If the electric charge on the plates above and
  below the droplets is adjusted, the eletrostatic
  attractive force pulling a droplet upward is just
  balanced by the force of gravity pulling the
  droplet downward.
• From the equations describing these forces,
  Millikan calculated the charge of the droplet.

43
Structure of the nuclear atom
44
Structure of the nuclear atom

• Different droplets had different charges, but
  Millikan found that each was a whole-number
  multiple of the same smaller charge, -1.60 ?
  10-19 C (C represents the coulomb, the SI unit of
  electric charge).
• Millikan assumed this to be the fundamental unit
  of charge, the charge of an electron.
• Because the charge-to-mass ratio of the electron
  was known, the mass of an electron could be
  calculated.
• The currently accepted value for the electron
  mass is 9.109389 ? 10-28 grams, and the currently
  accepted value of the electric charge is
  -1.60217733 ? 10-19 C.
• When talking about the properties of fundamental
  particles, we always express the charge relative
  to the charge on the electron, which is given the
  value of -1.
• Additional experiments showed that cathode rays
  had the same properties as the beta particles
  emitted by radioactive elements, providing
  further evidence that the electron is a
  fundamental particle of matter.

45
In early experiments on electricity and matter,
an electric current was passed through a glass
tube containing

• Water.
• Gas under high pressure.
• Liquid oxygen.
• Gas under low pressure.

46
In a glass tube, electrical current passes from
the negative electrode called the _____, to the
other electrode.

• Cathode
• Anode
• Electron
• Millikan

47
The rays produced in a cathode tube in early
experiments were

• Unaffected by a magnetic field.
• Deflected away from a negative plate.
• Found to carry a positive charge.
• Striking the cathode.

48
After measuring the ratio of the charge of a
cathode ray particle to its mass, Thomson
concluded that the particles

• Had no mass.
• Had a very small mass.
• Had a very large mass.
• Carried a positive charge.

49
Millikans experiments

• Demonstrated that the electron carried no charge.
• Demonstrated that the electron carried the
  smallest possible positive charge.
• Measured the charge on the electron.
• Demonstrated that the electron was massless.

50
Because any element used in the cathode produced
eletrons, scientists concluded that

• All atoms contained electrons.
• Only metals contained electrons.
• Atoms were indivisible.
• Atoms carried a negative charge.

51
Structure of the nuclear atom

• The first evidence of a fundamental positive
  particle came from the study of canal rays,
  which were observed in a special cathode-ray tube
  with a perforated cathode (Eugen Goldstein
  (1850-1930)).
• When a high voltage is applied to the tube,
  cathode rays are observed.
• On the other side of the perforated cathode, a
  different kind of ray is seen.
• Because these rays are attracted toward a
  negatively charged plate, they must be composed
  of positively charged particles.

52
Structure of the nuclear atom

• Each gas used in the tubes gives a different
  charge-to-mass ratio for the positively charged
  particles.
• When hydrogen gas is used, the largest
  charge-to-mass ratio is obtained, suggesting that
  hydrogen provides positive particles with the
  smallest mass.
• These were considered to be the fundamental,
  positively charged particles of atomic structure
  and were later called protons (from a Greek word
  meaning the primary one) by Ernest Rutherford.
• The mass of a proton is known from experiment to
  be 1.672623 ? 10-24 grams.
• The relative charge on the proton, equal in size
  but opposite in sign to the charge on the
  electron , is 1.

53
Structure of the nuclear atom

• Because atoms normally have no charge, the number
  of protons must equal the number of electrons in
  an atom.
• Most atoms have masses greater than would be
  predicted on the basis of protons and electrons,
  indicating that uncharged particles must also be
  present.
• Because this third type of particle has no
  charge, the usual methods of detecting particles
  could not be used.

54
Structure of the nuclear atom

• In 1932, the British physicist James Chadwick
  (1891 1974) devised an experiment that produced
  these expected neutral particles and then
  detected them by having them knock hydrogen ions,
  a detectable species, out of parafin.
• This particles, now known as the neutron, has no
  electric charge and has a mass of 1.6749286 ?
  10-24 grams, slightly greater than the mass of a
  proton.

55
Structure of the nuclear atom

• J.J. Thomson had supposed that an atom was a
  uniform sphere of positively charged matter
  within which thousands of electrons circulated in
  coplanar rings.
• Thomson and his students thought the only
  question was the number of electrons circulating
  within this sphere.
• To test this model experimentally, they directed
  a beam of electrons at a very thin metal foil.
• They expected that as the beam traveled through
  the foil, the electrons would encounter the very
  large number of electrons within the atoms, and
  the negative charges would repel each other.
• A tiny deflection of the beam from its straight
  path should be observed with each encounter, the
  size of the total deflection related to the
  number of electrons in the atom.
• A deflection was observed, but it was much
  smaller than expected, forcing Thomson and his
  students to revise their estimate of the number
  of electrons, but not their model of the atom.

56
Structure of the nuclear atom

• About 1910, Ernest Rutherford (1871 1937)
  decided to test Thomsons model further.
• Rutherford had discovered earlier that alpha rays
  consisted of positively charged particles having
  the same mass as helium atoms.
• He reasoned that, if Thomsons atomic model were
  correct, a beam of such massive particles would
  be deflected very little as it passed through the
  atoms in a very thin sheet of gold foil.
• Rutherfords associate, Hans Geiger (1882
  1945), and a student, Ernst Marsden, set up the
  appartus and observed what happened when alpha
  particles hit the foil.

57
Structure of the nuclear atom

• Most of the particles passed straight through,
  but Geiger and Marsden were amazed to find that a
  few alpha particles were deflected at large
  angles, and some almost came straight back.
• Rutherford later said, It was about as credible
  as if you had fired a 15-inch (artillery) shell
  at a piece of paper and it came back and hit you.

58
Structure of the nuclear atom

• The only account for this was to discard
  Thomsons model and to conclude that all of the
  positive charge and most of the mass of the atom
  is concentrated in a very small volume.
• Rutherford called this tiny core of the atom the
  nucleus.
• The electrons occupy the rest of the space in the
  atom.
• From their results Rutherford, Geiger, and
  Marsden calculated that the nucleus of a gold
  atom had a positive charge in the range of 100 ?
  20 and a radius of about 10-12 cm.
• The currently accepted values are 79 for the
  charge and about 10-13 cm for the radius.

59
Structure of the nuclear atom
60
Structure of the nuclear atom

• These experiments of Thomson, Rutherford, and
  others led to the now-accepted model of the atom.
• Three primary particles protons, neutrons, and
  electrons make up all atoms.
• The model places protons and neutrons in a very
  small nucleus, which means that the nucleus
  contains all the positive charge and almost all
  the mass of an atom.
• Negatively charged electrons surround the nucleus
  and occupy most of the volume.
• Atoms have no net charge, so the number of
  electrons outside the nucleus equals the number
  of protons in the nucleus.

61
Structure of the nuclear atom

• Although experiments in the early part of the
  20th century established that electrons occupied
  the space outside the nucleus, the detailed
  arrangement of the electrons (electronic
  structure) was completely unknown to Rutherford
  and other scientists at that time.

62
Who discovered the nucleus by bombarding gold
foil with positively charged particles and noting
that some particles were widely deflected?

• Rutherford
• Dalton
• Chadwick
• Bohr

63
In Rutherfords experiments, most of the particles

• Bounced back.
• Passed through the foil.
• Were absorbed by the foil.
• Combined with the foil.

64
Because a few positively charged particles
bounced back from the foil, Rutherford concluded
that such particles were

• Striking electrons.
• Indivisible.
• Repelled by densely packed regions of positive
  charge.
• Magnetic.

65
Rutherford fired positively charged particles at
metal foil and concluded that most of the mass of
an atom was

• In the electrons.
• Concentrated in the nucleus.
• Evenly spread throughout the atom.
• In rings around the atom.

66
A nuclear particle that has about the same mass
as a proton, but with no electrical charge is
called a (an)

• Nuclide.
• Neutron.
• Electron.
• Isotope.

67
Which part of an atom has a mass approximately
equal to 1/2000 of the mass of a common hydrogen
atom?

• Nucleus.
• Electron.
• Proton.
• Electron cloud.

68
An atom is neutral because

• Neutrons balance the protons and electrons.
• Nuclear forces stabilize the charges.
• The numbers of protons and electrons are equal.
• The numbers of protons and neutrons are equal.

69
Most of the volume of an atom ic occupied by the

• Nucleus.
• Nuclides.
• Electron cloud.
• Protons.

70
Distinguishing among atoms

• What makes an atom of one element different from
  an atom of another element?
• All atoms of an element have the same number of
  protons in the nucleus (Atomic Number).
• Atomic Number (Z) the number of protons in the
  nucleus of each atom of that element.
• Currently known elements are listed in the
  periodic table inside the back cover of your
  book.
• The integer number at the top of the box for each
  element is its atomic number.
• The atomic number identifies an element.
• Because atoms are neutral, we know from the
  atomic number that their must be an equal number
  of electrons in the atom.

71
Distinguishing among atoms

• Atoms of a given element that differ in the
  number of neutrons, and consequently in mass, are
  called isotopes.
• Isotope Atom of the same element that have
  different masses.
• In Daltons atomic theory, he stated that atoms
  of the same element have the same mass, this was
  proven incorrect with the discovery of isotopes.
• Isotopes of a particular element all have the
  same number of protons and electrons, but a
  different number of neutrons.
• Although isotopes have different masses, they do
  not differ significantly in their chemical
  behavior.

72
Distinguishing among atoms

• Most atoms have at least two stable
  (nonradioactive) isotopes.
• A few have only one isotope (Al, F, P).
• Other elements have many isotopes (Tin 10
  stable isotopes).
• An atom of a specific isotope is called a
  nuclide.
• Hydrogen has three isotopes
• Protium (1 p, 1 e-, 0 n0) (most common)
• Deuterium (1 p, 1 e-, 1 n0)
• Tritium (1 p, 1 e-, 2 n0) (radioactive)

73
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74
Isotopes are atoms of the same element that have
different

• Principal chemical properties.
• Masses.
• Numbers of protons.
• Numbers of electrons.

75
Isotopes of an element contain different numbers
of

• Electrons.
• Protons.
• Neutrons.
• Nuclides.

76
The most common form of hydrogen has

• No neutrons.
• One neutron.
• Two neutrons.
• Three neutrons.

77
The only radioactive form of hydrogen is

• Protium.
• Deuterium.
• Tritium.
• Quadrium.

78
All isotopes of hydrogen contain

• One neutron.
• Two electrons.
• One proton.
• Two nuclei.

79
Distinguishing among atoms

• Generally, we refer to a particular isotope by
  giving its mass number.
• Mass Number (A) The total number of protons and
  neutrons in the nucleus of an isotope.
• Examples
• Oxygen-16 (8 p, 8 e-, 8 n0)
• Oxygen-17 (8 p, 8 e-, 9 n0)
• Oxygen-18 (8 p, 8 e-, 10 n0)

80
Distinguishing among atoms

• There are two ways for specifying isotopes
• Hyphen Notation
• The mass number is added with a hyphen to the end
  of the name for that specific element.
• Examples
• Boron-10
• Boron-11
• Nitrogen-14
• Nitrogen-15
• Magnesium-24
• Magnesium-25
• Magnesium-26
• Uranium-235
• Uranium-238

81
Distinguishing among atoms

• 2. Nuclear Symbols
• The atomic number is added to the symbol for the
  element as a subscript on the left and the mass
  number is added as a superscript on the left.

Mass Number
AZSy
Element Symbol
Atomic Number
82
Distinguishing among atoms

• Examples
• 63Li
• 73Li
• 3216S
• 3316S
• 3416S
• 3616S
• 3517Cl
• 3717Cl
• 6329Cu
• 6529Cu

83
An aluminum isotope consists of 13 protons, 13
electrons, and 14 neutrons. Its mass number is

• 13.
• 14.
• 27.
• 40.

84
Zn-66 (atomic number 30) has

• 30 neutrons.
• 33 neutrons.
• 36 neutrons.
• 96 neutrons.

85
Ag-109 has 62 neutrons. The neutral atom has

• 40 electrons.
• 47 electrons.
• 53 electrons.
• 62 electrons.

86
Chlorine has atomic number and mass number 35. It
has

• 17 protons, 17 electrons, and 18 neutrons.
• 35 protons, 35 electrons, and 17 neutrons.
• 17 protons, 17 electrons, and 52 neutrons.
• 18 protons, 18 electrons, and 17 neutrons.

87
Carbon-14 (atomic number 6), the radioactive
nuclide used in dating fossils, has

• 6 neutrons.
• 8 neutrons.
• 10 neutrons.
• 14 neutrons.

88
An electrically neutral atom of mercury (atomic
number 80) has

• 80 neutrons and 80 electrons.
• 40 protons and 40 electrons.
• 80 protons and 80 neutrons.
• 80 protons and 80 electrons.

89
Distinguishing among atoms

• Atoms have extremely small masses.
• The mass of the heaviest known atom is on the
  order of 4 ? 10-22 grams.
• Because it would be cumbersome to express such
  small masses in grams, we use the atomic mass
  unit (amu).
• One amu equals 1.66054 ? 10-24 grams.
• The masses of the proton and neutron are very
  nearly equal, and both are much greater than that
  of an electron a proton has a mass of 1.0073
  amu, a neutron 1.0087 amu, and an electron 5.486
  ? 10-4 amu.
• It would take 1836 electrons to equal the mass of
  one proton, so the nucleus contains most of the
  mass of an atom.
• The amu is presently defined by assigning a mass
  of exactly 12 amu to an atom of the 12C isotope
  of carbon.
• In these units the mass of the 1H nuclide is
  1.0073 amu and that of the 16O nuclide is 15.9949
  amu.

90
Distinguishing among atoms

• Most elements occur in nature as mixtures of
  isotopes.
• We can determine the average atomic mass of an
  element by using the masses of it various
  isotopes and their relative abundances (Percent
  Abundance).
• Percent Abundance The percentage of the atoms of
  a natural sample of the pure element represented
  by a particular isotope.

91
Distinguishing among atoms

• To find average atomic mass
• ( Abundance/100)(Mass of Isotope 1) (
  Abundance/100)(Mass of Isotope 2) ..
• The atomic mass of each stable element has been
  determined by experiment, and these masses appear
  in the periodic table in the back of the book.
• In the periodic table, each elements box
  contains the atomic number, the element symbol,
  and the atomic mass.
• For unstable (radioactive), synthetic elements,
  the mass number of the most stable isotope is
  given in parentheses.

92
Distinguishing among atoms

• Example
• Naturally occurring carbon is composed of 98.93
  12C and 1.07 13C. The masses of these nuclides
  are 12 amu (exactly) and 13.00335 amu. What is
  the average atomic mass of carbon?
• (0.9893)(12 amu) (0.0107)(13.00335 amu)
• 12.0107 amu

93
Distinguishing among atoms

• The average atomic mass of each element
  (expressed in amu) is also known as its atomic
  weight (used most commonly).
• Example
• Naturally occurring chlorine is 75.78 35Cl,
  which has an atomic mass of 34.969 amu, and
  24.22 37Cl, which has an atomic mass of 36.966
  amu. Calculate the average atomic mass of
  chlorine?
• (0.7578)(34.969 amu) (0.2422)(36.966 amu)
• 35.453 amu

94
Distinguishing among atoms

• Example
• Three isotopes of silicon occur in nature 28Si
  (92.23), which has a mass of 27.97693 amu, 29Si
  (4.68), which has a mass of 28.97649 amu and
  30Si (3.09), which has a mass of 29.97377 amu.
  Calculate the atomic weight of silicon.
• (0.9223)(27.97693 amu) (0.0468)(28.97649 amu)
  (0.0309)(29.97377 amu)
• 28.085 amu

95
In determining atomic mass units, the standard is
the

• C-12 atom.
• C-14 atom.
• H-1 atom.
• O-16 atom.

96
The abbreviation for atomic mass unit is

• amu.
• mu.
• a.
• µ.

97
The average atomic mass of an element is the
average of the atomic masses of its

• Naturally occurring isotopes.
• Two most abundant isotopes.
• Nonradioactive isotopes.
• Artificial isotopes.

98
Distinguishing among atoms

• Now that you can differentiate between atoms of
  different elements and also between isotopes of
  the same element, you need to understand how the
  elements are organized with respect to each
  other.
• A periodic table is an arrangement of elements in
  which the elements are separated into groups
  based on a set of repeating patterns.
• A periodic table allows you to easily compare the
  properties of one element (or group of elements)
  to another element (or group of elements).

99
Distinguishing among atoms
1 H Hydrogen 1.00794
Elements Symbol
Atomic Number
Average Atomic Mass
Elements Name
100
Distinguishing among atoms

• Each horizontal row of the periodic table is
  called a period.
• There are seven periods in the modern periodic
  table.
• The number of elements per period ranges from two
  to 32.
• Within a given period, the properties of the
  elements vary as you move across it from element
  to element.

101
Distinguishing among atoms

• Each vertical column of the periodic table is
  called a group or family.
• Elements within a group have similar chemical and
  physical properties.
• Each group is identified by a number and the
  letter A or B.
• Example Group 2A contains the elements
  beryllium, magnesium, calcium, strontium, barium,
  and radium.
• You will learn more about specific trends in the
  periodic table in Chapter 6.

102
Elements in a group or column in the periodic
table can be expected to have similar

• Atomic Masses.
• Atomic Numbers.
• Numbers of Neutrons.
• Properties.

103
A horizontal row of blocks in the periodic table
is called a (an)

• Group.
• Period.
• Family.
• Octet.