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The Periodic Table

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Title: The Periodic Table


1
The Periodic Table
  • Chapter 6
  • Chemistry 112

2
6.1 Development of the Modern Periodic Table
  • The properties of the elements in the table
    repeat in a periodic way which is why the term
    periodic is used to describe the table.
  • Example periods of the moon

3
  • Antoine Lavoisier was credited with
  • Compiling a list of 23 elements including silver,
    gold, carbon and oxygen

4
  • In 1864, Newlands noticed if he arranged the
    elements by atomic mass, their properties
    repeated every eight element. He named this
    relationship the Law of Octaves, after the
    musical notes that repeat every eight tone. The
    acceptance of this law was hampered because the
    law did not work for all elements and octave
    was criticized because many thought the musical
    analogy was unscientific.

5
  • Lothar Meyer and Dmitri Mendeleev were the two
    scientists that demonstrated a relationship
    between atomic mass and elemental properties in
    1869. Mendeleev received the most credit
    because he published first.

6
  • Mendeleev is known as the father of the periodic
    table

7
  • Mendeleevs table was widely accepted because he
    was able to predict the existence and properties
    of undiscovered elements

8
  • Henry Mosley is credited with arranging the
    Periodic Table by atomic number, instead of by
    atomic mass.

9
  • Periodic Law is a periodic repetition of chemical
    and physical properties of the elements when they
    are arranged by increasing atomic number.

10
  • Group elements in the same column (also known
    as family)
  • Periods elements in the same horizontal row

11
  • A Representative elements
  • B transition elements

12
  • The three main classifications of elements are
    Metals, Non-metals, and metalloids.

13
  • Some common properties of metals, nonmetals, and
    metalloids are
  • Metals hard, shiny, conduct electricity and
    heat, malleable, ductile.
  • Nonmetals brittle, dull, poor conductors
  • Metalloids properties of metals and nonmetals

14
  • Metals are to the left of the stairstep (not
    including the metalloids), except Hydrogen.
  • Non-metals are to the right.
  • All elements touching the line are metalloids
    EXCEPT for Aluminum

15
  • These definitions will be discussed more
    completely later in your notes
  • Atomic radii - Atomic radius measures size of
    atoms and is defined by how closely an atom lies
    to a neighboring atom.
  • Ionization energy - Energy required to remove an
    electron from an atom which is used to overcome
    the attraction between the positive nucleus and
    the negative electron.
  • Electronegativity - The relative ability of an
    atom to attract electrons in a chemical bond.

16
Advancement of Chemistry
  • There are three factors that led to the
    advancement of chemistry in the 1800s
  • The advent of electricity
  • Development of a spectrometer
  • The industrial revolution

17
Oxidation Numbers
  • The positive or negative charge of an ion
  • Equal to the number of electrons transferred from
    an atom to form an ion
  • Sodium loses one electron 1
  • Chlorine gains one electron -1
  • Why do they act like this????? Think about how
    many valence electrons all atoms want.

18
Oxidation Numbers of Groups
  • Alkali metals have 1 valence electron that they
    want to LOSE. They all form a 1 charge.
  • Alkaline earth metals have 2 valence electrons
    that they want to LOSE. They all form a 2
    charge.
  • Transition metals for ions in more than one way.
    We will skip the d-block for this chapter and
    revisit it next unit.

19
  • Boron Group has three valence electrons they want
    to LOSE. They all form a 3 charge.
  • Carbon Group has four valence electrons. It is
    just as easy to lose four as it is to gain four
    electrons. Carbon group forms a 4 or -4 charge.
    (there are some exceptions that we will talk
    about next unit).

20
  • Nitrogen group has five valence electrons. They
    want to GAIN three more to make a total of eight.
    They all form a -3 charge.
  • Oxygen group has six valence electrons and they
    want to GAIN two more. They all form a -2
    charge.
  • Halogens all have seven valence electrons and
    want to GAIN one more. They all form a -1
    charge.
  • Nobel gases have a full valence shell already and
    do not form ions.

21
Practice Problems What are the oxidation
numbers of the following ions?
  • K
  • P
  • I
  • Ar
  • Al
  • S
  • 1
  • 3-
  • 1-
  • 0
  • 3
  • 2-

22
Write the ion with the proper oxidation number or
charge for each of the following
  • Mg
  • Cs
  • Al
  • N
  • O
  • Br
  • Kr

23
Trends on the Periodic Table
  • You will be responsible for five trends.
  • 1. Atomic radius
  • 2. Metal reactivity
  • 3. Non-metal reactivity
  • 4. Ionization Energy
  • 5. Electronegativity

24
Atomic Radius
  • All trends on the periodic table relate to atomic
    radius! They all have to do with how easy it is
    to lose or gain electrons.
  • Atomic radius measures size of atoms
  • Defined by how closely an atom lies to a
    neighboring atom.
  • Trend for within a period and within a
    group.graph to find out!!

25
Trends for Atomic Radius
  • Atomic Radius INCREASES as you go down the P.T.
    because there are increasing number of energy
    levels (think of it as layers of fatgetting
    bigger).
  • Atomic Radius DECREASES across the P.T. because
    there is more of a positive pull on the
    electrons. More of an attraction pulls the
    electrons a bit closer. (think about having four
    people pull on a rope versus five. Five people
    are able to pull harder and will therefore bring
    the rope closer)

26
Hands on a clock
  • A way to remember the trends is to relate them to
    hands on a clock.
  • You point an arrow in the direction where the
    trend INCREASES.
  • For atomic Radius, you would have a hand pointing
    down and a hand pointing left.
  • This looks like 930 on a clock.

27
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28
Practice Problems Which atomic radius is
bigger? (Hint, write he symbols how they appear
on the periodic tablewhich ever one the arrow
points to is bigger)
  • Na or S
  • Ba or Ca
  • K or Ra
  • Cl or Au
  • As or P
  • Fr or Cs
  • Na
  • Ba
  • Ra (period matters more than group)
  • Au
  • As
  • Fr

29
Ion Radius
  • Atoms that lose electrons to become cations,
    always become smaller. (when you lose weight, you
    become smaller!) Since there are less electrons
    that are repelling each other, they can settle in
    a little closer to each other.
  • Atoms that gain electrons to become anions,
    always become larger. (when you gain weight, you
    become bigger). More electrons mean more
    repulsion. They have to spread out even further!

30
Practice Problems Which is SMALLER! (pay
attention to the questions)
  • Ca atom or Ca ion (hintwhat oxidation number
    does Ca form?)
  • F atom or F ion
  • Fe2 or Fe3
  • Ca2 ion is smaller because it loses 2 e-
  • F atom is smaller because the ion gains 1e-
  • Fe3 is smaller because it loses 3 e- compared to
    Fe2 that loses 2 e-

31
Metal Reactivity
  • When Metals react, they lose electrons due to the
    fact they have a low amount of valence electrons.
  • Would it be easier for an electron to be lost if
    it was close to the nucleus or far away?
  • Electrons are lost easier when they are farther
    away from the nucleus since there is less of a
    positive pull on them.
  • Therefore, the greater the Atomic Radius, the
    more reactive the metal is. (easier to lose
    more reactive)

32
Metal Reactivity Trendhttp//video.google.com/vid
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Blocked by Cobb County but you can watch this at
home!
  • INCREASES DOWN a Group due to the fact that the
    electrons are further from the positive pull of
    the nucleus and they are easier to remove.
  • DECREASES ACROSS a period due to the fact that
    there is more of a positive pull holding onto the
    electrons and they are harder to remove.

33
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34
Which metal is more reactive?
  • Na or Cs
  • Ni or Fe
  • Au or Cu
  • Ag or K
  • Cs
  • Fe
  • Au
  • K (group always matters morealkali metals are
    more reactive than transition metals! Good thing
    or we wouldnt be able to wear jewerly!!)

35
Non Metal Reactivity
  • When Non-metals react, they gain electrons due to
    the fact that they have a high number of valence
    electrons.
  • Would it be easier to gain electrons if they were
    closer to the nucleus or farther away?
  • Electrons are gained easier if they are closer to
    the nucleus due to the positive attraction.
  • The trend is exactly opposite of Metal Reactivity
    and Atomic Radius.
  • Noble Gases are excludedthey hardly ever react!

36
Non-Metal Reactivity Trend
  • DECREASES DOWN a group due to the fact that the
    atomic radius is getting larger and it is harder
    for the positive nucleus to attract electrons.
  • INCREASES ACROSS a period due to the fact that
    there is more of a positive pull and is easier to
    attract the electrons.
  • What time would this be on a clock?

37
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38
Which non-metals is more reactive?
  • F or Br
  • S or Cl
  • N or He
  • F
  • Cl
  • N (remember Nobel gases dont react!)

39
Ionization Energy
  • Energy required to remove an electron from an
    atom.
  • Energy is needed to overcome the attraction
    between the positive charge in the nucleus and
    the negative charge of the electron.
  • What type of atoms want to lose electrons?
  • Metalstherefore, they have lower I.E.

40
  • 1st I.E. Energy to remove the 1st electron
  • 2nd I.E. Energy to remove the 2nd electron
  • 3rd I.E. etc.

41
Trends for I.E.
  • I.E. DECREASES down a group. It becomes easier
    to remove an electron because there are more
    energy levels blocking the positive pull.
  • I.E. INCREASES across a period. There is more of
    a positive pull holding on to the electrons,
    therefore, it is harder to remove them. You are
    also approaching the Non-metals, which gain
    electrons!

42
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43
When will you see a drastic increase in I.E.???
  • After the element has lost its valence electrons.
  • Example Potassiums 1st I.E. is low. Its 2nd
    is very high.
  • Calciums 1st and 2nd I.E. is low. Its 3rd is
    high.
  • See page 168 Table 6-2

44
Practice problems Which atom has a lower 1st
I.E.?
  • Ca or Br
  • Cl or I
  • Au or Cu
  • Ag or Rb
  • Mg or I
  • Ca
  • I
  • Au
  • Rb
  • Mg (metals always have a lower I.E. than
    nonmetals.metals want to lose, nonmetals want to
    gain)

45
Electronegativity
  • The relative ability of its atoms to attract
    electrons in a chemical bond.
  • What type of elements want to ATTRACT electrons??
  • Nonmetals! Therefore, nonmetals have a high E.N.
    value.
  • Numerical value from 0.0 (least) to 4.0 (most)
  • Unit is a Pauling
  • Noble Gases are excluded!!

46
Trends for E.N.
  • Down a group, E.N. DECREASES. More energy levels
    makes it harder to attract electrons
  • Across a period, E.N. INCREASES. More and more
    valence electrons, and less for the atoms to have
    a full octet. More positive pull also. Also,
    pointing towards the non-metals.

47
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48
Which element is more electronegative?
  • F or Be
  • S or Po
  • Sr or Br
  • Kr or As
  • F (highest E.N. value on P.T.)
  • S
  • Br (Nonmetals always higher than metals)
  • As (Nobel gases are are excluded)

49
Using E.N. values to calculate bond types
  • Ionic Bond electrostatic force holds oppositely
    charged particles together in an ionic compound.
    Formed by losing or gaining electrons.
  • Polar Covalent Bond formed when electrons are
    not shared equally.
  • Non-polar Covalent bond formed from the equal
    sharing of valence electrons.

50
Like Dissolves Like
  • Water is a polar covalent solvent (meaning one
    end of the molecule has a slight charge).
  • Only Ionic compounds and Polar compounds will
    dissolve in water because they also have a
    charge.
  • Non-polar covalent compounds will not dissolve in
    water because it does not have a charge.

51
E.N. differences
  • Subtract the lowest E.N. element in a compound
    from the highest E.N. element (the subscripts do
    not matter)
  • Difference greater than 1.70 IONIC
  • Difference 0.5 to 1.70 Polar
  • Difference less than .5 Non-polar

52
Use your book pg. 403 to find the E.N. values of
each element. What kind of bond will the
compounds likely have?
  • SO2
  • BaCl2
  • H2S
  • Ca2N3
  • As2O3
  • .94 polar covalent
  • 2.27 ionic
  • 0.3 non-polar
  • 2.04 ionic
  • 1.26 polar covalent
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