Chemistry Chapter 4

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Chemistry Chapter 4

• The Structure of the Atom

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4.1 Early Ideas About Matter

• Objectives
• 1. Compare and contrast the models of Democritus,
  Aristotle and Dalton
• 2. Understand how Daltons theory explains
  conservation of mass

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Democritus

• Definition An atom is the smallest particle of
  an element that retains its identity in a
  chemical reaction
• Democritus believed that atoms were indivisible
  and indestructible
• His approach was not based on scientific method

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Aristotle

• Aristotle did not believe empty space could exist
• He discredited the ideas of Democritus because
  Democritus could not explain what held atoms
  together
• Aristotle said that all matter was composed of
  air, water, earth fire
• These ideas went unchallenged for 2000 years

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Dalton

• Dalton used scientific method and transformed
  Democrituss idea into scientific theory
• There are 6 points to his theory (check 4.1 in
  your book)
• Here is some of Daltons theory
• 1. all elements (matter) are composed of atoms
• 2. atoms of the same element are identical (been
  disproven)

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• 3. Atoms are indivisible indestructible (been
  disproven)
• 4. Atoms of any element are unique from atoms of
  any other element
• 5.Elements can combine in whole number ratios to
  form compounds
• 6. Chemical reactions occur when atoms are
  separated, joined, or rearranged
• Also, Atoms of one element are never changed into
  atoms of another element in a chemical reaction

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Dalton and Conservation of Mass

• Recall that mass is always conserved in chemical
  reactions
• Using Daltons theory the number of atoms of each
  type is the same before and after reacting
• Daltons experiments provided evidence and
  explanation of the composition of chemical
  compounds
• Conservation of mass led to acceptance of
  Daltons theory

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Instruments for observing atoms

• An atom is the smallest particle of an element
  that still retains all the properties of that
  element
• Atoms are very small a copper penny would have
  roughly 2.4 x 1022 atoms (unfathomably small !!!)
• The radius of atom ranges from 5 x 10-10 m to 2 x
  10-10 m
• Despite their small size atoms are observable
  with instruments such as a scanning tunneling
  microscope

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4.2 Defining the Atom

• Objectives
• 1. Define atom
• 2. Distinguish between subatomic particles in
  terms of relative charge and mass
• 3. Describe the structure of the atom, including
  the locations of the subatomic particles

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3 Types of Subatomic Particles

• Definition an atom is the smallest particle of
  an element that retains the properties of the
  element
• Atoms formed from 3 types of subatomic particles
  electrons, protons, and neutrons
• Electrons have a negative charge, weigh 1/1840th
  of protons and neutrons
• Symbol is e-

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• Protons have a positive charge and weigh 1840
  times as much as electrons
• Proton symbol is p
• Neutrons have no charge, weigh nearly the same as
  protons
• Neutron symbol is n0

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Thomson, Rutherford atomic structure

• J.J. Thomson- did cathode ray experiments that
  showed charged particles exist in atoms
• Thomson invented the plum pudding model
• The model consisted of a spherically shaped atom
  composed of a uniformly distributed positive
  charge
• The negatively charged electrons existed in the
  middle of the of the positive charges

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• Rutherford Thomson were contemporaries
• Rutherford set out to disprove Thomsons model
• Rutherford-alpha particle experiments
• 1. Atoms contain a nucleus with protons and
  neutrons, electrons orbit around nucleus
• Definition nucleus- tiny, central core of atom
  composed of p and n0

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• 2. Atoms composed mainly of empty space
• In the nuclear atom, the protons and neutrons are
  located in the nucleus
• The electrons are distributed around the nucleus
  and the nucleus occupies almost all the volume of
  an atom

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4.3 How atoms differ

• Objectives
• 1. Explain the role of atomic number in
  determining identity of an atom
• 2. Calculate the number of electrons, protons
  neutrons in an atom, given its mass number and
  atomic number
• 3. Define isotope
• 4. Explain why atomic masses are not whole
  numbers

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Atomic Number

• Elements differ because they have a different
  number of protons in the nucleus
• Protons determine atom identity
• How many protons is referred to as atomic number
• Given as the top number in the periodic table
• In an atom, the number of electrons is equal to
  the number of protons

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• Definition Mass Number is the total weight of an
  atom
• Mass number equals the number of protons plus the
  number of neutrons mass p n0
• Weight of electrons ignored (1/1840th)
• Given as the bottom number in the periodic table
• Number of neutrons is the difference between mass
  number and atomic number
• n0 mass - atomic number (p)

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• An atoms composition can be found if the atomic
  number or mass number is known

Element Atomic Mass Pro-tons Neu- trons Elec- trons
K 19       19
          5
  16        
    23      
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Why Use the Periodic Table

• Definition a periodic table is an arrangement of
  elements in which the elements are separated into
  groups based on a set of repeating properties
• A periodic table allows you to easily compare the
  properties of one element (or group of elements)
  to another element (or group of elements)

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• Definition a period is each horizontal row of
  the periodic table
• There are seven rows or periods in the modern
  periodic table
• Definition a group or family is each vertical
  column of the periodic table
• Each group is identified by a number and letter A
  or B
• There are a total of 18 groups, 8 As and 10 Bs

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Isotopes

• Definition Atoms of the same element with a
  different number of neutrons are isotopes
• Because isotopes of an element have different
  numbers of neutrons, they also have different
  mass numbers
• Mass number is different, same element
• Written with number of neutrons behind,
  carbon-12, carbon-13
• Or with superscripts for neutrons (13C) and
  subscripts for protons (6C) put together they are
  written as (613 C)

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• Actual masses of protons, neutrons and electrons
  is really small
• More useful to talk about relative masses
• A reference was needed
• Carbon-12 chosen as standard
• Carbon-12 was assigned 12 Atomic Mass Units
• Definition atomic mass unit (amu) is one twelfth
  the mass of a carbon-12 atom
• One mass unit is given for each proton and
  neutron

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Calculating Atomic Mass Why Atomic Masses are
not Whole Numbers

• Most elements in nature are found as mixtures of
  isotopes
• Definition relative abundance is how often you
  encounter isotopes in nature and is expressed as
  a percent
• Relative abundance of isotopes determines an
  averaged mass number
• Definition atomic mass of an element is the
  weighted average mass of the atoms in a naturally
  occurring sample of the element
• Atomic mass accounts for the isotopes of an
  element as it occurs in nature

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• To calculate the atomic mass of an element,
  multiply the mass of each isotope by its natural
  abundance, expressed as a decimal, then add the
  products together
• Problem Calculate the atomic mass of bromine.
  The 2 isotopes of bromine have atomic masses and
  relative abundances of 78.92 amu (50.69) and
  80.92 amu (49.31)
• A 79.91 amu

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4.4 Unstable Nuclei and Radioactive Decay

• Objectives
• 1. Explain the relationship between unstable
  nuclei and radioactive decay
• 2. Characterize alpha, beta and gamma radiation
  in terms of mass and charge

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The Relationship

• Definition radioactivity is spontaneously
  emitted radiation produced by some substances
• Definition radiation is the rays and particles
  emitted by those substances
• Definition nuclear radiation is a reaction that
  involves a change in an UNSTABLE atomic nucleus

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• Most radioactive atoms tend to be heavier
  elements with a large nucleus making it emit
  radiation because their nucleus is unstable
• Definition radioactive decay is a spontaneous
  process in which unstable nuclei lose energy by
  emitting radiation
• 200 years ago, scientists began investigating
  radiation by directing it between electrically
  charged plates

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• Definition alpha radiation is the radiation that
  was deflected toward the negatively charged plate
• It consists of alpha particles
• Alpha particles are composed of 2 neutrons and 2
  protons, no electrons
• Therefore an alpha particle carries a charge of
  2 is the nucleus of a helium (He) atom

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• Definition beta radiation is the radiation that
  was deflected toward the positively charged plate
• Each beta particle is an electron with a 1-
  charge is an electron which is emitted
• Definition gamma rays or gamma radiation is high
  energy radiation with no mass
• It is denoted by the Greek letter gamma which
  looks like this

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• The main factor in determining nuclear stability
  is the ratio of neutrons to protons
• Too many or too few neutrons are unstable and
  cause the nucleus to lose energy through
  radioactive decay until a stable nucleus is
  formed
• Alpha and beta particles are emitted until a
  stable nonradioactive nucleus is formed

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• Alpha radiation is the least harmful, can burn
  skin is stopped by paper or clothing
• is not harmful unless ingested or breathed in
• The electron of a beta particle is stopped by
  aluminum is more penetrating than alpha
• Gamma radiation is pure energy which is very
  penetrating harmful

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Alpha, beta gamma equations

• Examples of alpha
• Examples of beta
• Gamma will emit either alpha or beta particle
  gamma rays