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Title: Chemistry Chapter 4


1
Chemistry Chapter 4
  • The Structure of the Atom

2
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

3
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

4
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

5
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)

6
  • 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

7
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

8
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

9
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

10
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-

11
  • 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

12
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

13
  • 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

14
  • 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

15
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

16
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

17
  • 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)

18
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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      
20
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)

21
  • 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

22
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)

23
  • 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

24
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

25
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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

27
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

28
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

29
  • 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

30
  • 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

31
  • 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

32
  • 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

33
  • 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

34
Alpha, beta gamma equations
  • Examples of alpha
  • Examples of beta
  • Gamma will emit either alpha or beta particle
    gamma rays
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