Topic 3: Periodicity

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Topic 3 Periodicity

• 3.1 The periodic table
•  
• 3.1.1      Describe the arrangement of elements
  in the periodic table in order of increasing
  atomic number
• 3.1.2      Distinguish between the terms group
  and period
• 3.1.3      Apply the relationship between the
  electron arrangement of elements and their
  position in the periodic table up to z20.
• 3.1.4      Apply the relationship between the
  highest occupied energy level for an element and
  its position in the periodic table.
•  

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• Groups vertical columns (18)
• Have similar properties because have same number
  of electrons in outer shell
• Periods horizontal row (7)
• Family Names
• Group 1 alkali metals
• Group 2 alkaline earth metals
• Group 17 halogens
• Group 18 noble gases
• Group 3-12 Transition metals
• Groups 1,2, 13-18 representative elements

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3.2 Physical properties  3.2.1      Define the
terms first ionization energy and
electronegativity3.2.2      Describe and explain
the trends in atomic radii, ionic radii, first
ionization energy, electronegativities and
melting points for alkali metals (Li ? Cs) and
the halogens (F ? I).3.2.3      Describe and
explain the trends in atomic radii, ionic radii,
first ionization energy, and electronegativities
for elements across period3.2.4      Compare the
relative electronegative values of two or more
elements based on their position on the periodic
table.
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Atomic Size

• The electron cloud doesnt have a definite edge.
• They get around this by measuring more than 1
  atom at a time.
• Summary it is the volume that an atom takes up
• http//www.mhhe.com/physsci/chemistry/essentialche
  mistry/flash/atomic4.swf

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Group trends
H

• As we go down a group (each atom has another
  energy level) the atoms get bigger, because more
  protons and neutrons in the nucleus

Li
Na
K
Rb
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Trends Within Groups (Families)

• Increase as you move down a group.
• Even though nuclear charge increases as you go
  down a group, the orbital sizes increase so much
  that the atom becomes larger.
• The outer electrons are farther from the nucleus
  and are shielded from the positive charge of the
  nucleus by the other electrons.

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Periodic Trends

• atomic radius decreases as you go from left to
  right across a period.
• Why? Stronger attractive forces in atoms (as you
  go from left to right) between the opposite
  charges in the nucleus and electron cloud cause
  the atom to be 'sucked' together a little
  tighter. Remember filling up same energy level,
  little shielding occurring.

Na
Mg
Al
Si
P
S
Cl
Ar
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Ionization Energy

• Ionization Energy The energy needed to overcome
  the attraction between the positive charge in the
  nucleus and the negative charge of the electron.
• How much energy is needed to remove an electron
  from an atom.
• Energy is measured in Joules.

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Ionization Energy

• High ionization energy values indicate a strong
  hold on electrons.
• Unlikely to become a positive ion.
• Low ionization energy values indicate a weak hold
  on electrons.
• Likely to become positive ions.

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Ionic Size

• Cations form by losing electrons.
• Cations are smaller than the atom they come from.
• Metals form cations.
• Cations of representative elements have noble gas
  configuration.

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Ionic Size

• Anions form by gaining electrons.
• Anions are bigger than the atom they come from.
• Nonmetals form anions.
• Anions of representative elements have noble gas
  configuration.

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Periodic Trends

• Metals losing from outer energy level, more
  protons than electrons so more pull, causing it
  to be a smaller species.
• Non metals gaining electrons in its outer energy
  level, but there are less protons than electrons
  in the nucleus, so there is less pull on the
  protons, so found further out making it larger.

N-3
B3
O-2
F-1
Li1
C4
Be2
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Why do positive ions become smaller?

• Two Reasons
• The electron lost from the atom will always be a
  valence electron smaller radius.
• The lost electron no longer shields the other
  electrons from the positive nucleus, so they are
  pulled closer to the nucleus.

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Ionic Radii

• Atoms that gain electrons always become larger.
• Why?
• Additional electron causes other orbitals to be
  filled.
• Increased shielding causes other electrons to be
  farther away from the nucleus.

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Ionic Radii
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Ionic Radii Trends

• What is the pattern in Periods?
• Smaller until 5A.
• What is the pattern in Groups?
• Gradual increase in size.

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Size of Isoelectronic ions

• Positive ions have more protons so they are
  smaller.

N-3
O-2
F-1
Ne
Na1
Al3
Mg2
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Ionization Energy
http//www.iun.edu/cpanhd/C101webnotes/modern-ato
mic-theory/images/ionization-energy.jpg
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Ionization Energy
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Can I Remove More Than One Electron?

• A second, third, etc, electron can be removed
  from an atom.
• The ionization energies are termed accordingly
• 2nd Ionization energy to remove the 2nd electron.
• 3rd Ionization energy to remove the 3rd electron.

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2nd and 3rd Ionization Energies

• Do you think they are higher values or lower
  values than the 1st Ionization energy?
• Usually the values are higher since the atom
  holds onto the remaining electrons even tighter.

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Ionization Energies in kJ/mol
1 2 3 4 5 6 7 8
H 1312
He 2372 5250
Li 520 7297 11810
Be 899 1757 14845 21000
B 800 2426 3659 25020 32820
C 1086 2352 4619 6221 37820 47260
N 1402 2855 4576 7473 9442 53250 64340
O 1314 3388 5296 7467 10987 13320 71320 84070
F 1680 3375 6045 8408 11020 15160 17860 92010
Ne 2080 3963 6130 9361 12180 15240
Na 496 4563 6913 9541 13350 16600 20113 25666
Mg 737 1450 7731 10545 13627 17995 21700 25662
http//www.shodor.org/chemviz/ionization/students/
background.html
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Trends of Ionization Energy

• Within Periods
• Increase as you move left-to-right.
• Due to increase in nuclear charge and a tight
  hold on electrons.
• Within Groups
• Generally decreases as you move down a group.
• Electrons are farther away from nucleus.

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Ionization Energy

• What happens when sodium loses an electron?
• What is its electron configuration?
• Na 1s22s22p63s1
• Na1 1s22s22p6
• The octet rule states that atoms tend to gain,
  lose, or share electrons in order to acquire a
  full set of eight valence electrons.

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Electronegativity
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Yet Another Trend!

• Electronegativity ability of an element to
  attract an electron in a chemical bond.
• How badly does it want another electron?

Why no Values?
http//college.hmco.com/chemistry/intro/zumdahl/in
tro_chemistry/5e/students/protected/periodictables
/pt/pt/table/t_e2.gif
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Electronegativity

• The tendency for an atom to attract electrons to
  itself when it is chemically combined with
  another element.
• How fair it shares.
• Big electronegativity means it pulls the electron
  toward it.
• Atoms with large negative electron affinity have
  larger electronegativity.

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Group Trend

• The further down a group the farther the electron
  is away and the more electrons an atom has.
• So as you go from fluorine to chlorine to bromine
  and so on down the periodic table, the electrons
  are further away from the nucleus and better
  shielded from the nuclear charge and thus not as
  attracted to the nucleus. For that reason the
  electronegativity decreases as you go down the
  periodic table.

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Period Trend

• Electronegativity increases from left to right
  across a period
• When the nuclear charge increases, so will the
  attraction that the atom has for electrons in its
  outermost energy level and that means the
  electronegativity will increase

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Period trend

• Electronegativity increases as you go from left
  to right across a period.
• Why? Elements on the left of the period table
  have 1 -2 valence electrons and would rather give
  those few valence electrons away (to achieve the
  octet in a lower energy level) than grab another
  atom's electrons. As a result, they have low
  electronegativity. Elements on the right side of
  the period table only need a few electrons to
  complete the octet, so they have strong desire to
  grab another atom's electrons.

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Group Trend

• electronegativity decreases as you go down a
  group.
• Why? Elements near the top of the period table
  have few electrons to begin with every electron
  is a big deal. They have a stronger desire to
  acquire more electrons. Elements near the bottom
  of the chart have so many electrons that loosing
  or acquiring an electron is not as big a deal.
• This is due to the shielding affect where
  electrons in lower energy levels shield the
  positive charge of the nucleus from outer
  electrons resulting in those outer electrons not
  being as tightly bound to the atom.

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Shielding

• Shielded slightly from the pull of the nucleus by
  the electrons that are in the closer orbitals.
• Look at this analogy to help understand

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Electronegativity Trends
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Overall Trends!
http//campus.ru.ac.za/full_images/img05206111510.
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Melting Points of Group 1
Element Melting Point (K)
Li 453
Na 370
K 336
Rb 312
Cs 301
Fr 295
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Metallic bonding

• Collective bond, not a single bond
• Strong force of electromagnetic attraction
  between delocalized electrons (move freely).
• This is sometimes described as "an array of
  positive ions in a sea of electrons

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• Why does the melting point decrease going down
  the alkali metals family?
• Atoms are larger and their outer electrons are
  held farther away from the positive nucleus.
• The force of attraction between the metal ions
  and the sea of electrons thus gets weaker down
  the group.
• Melting points decrease as less heat energy is
  needed to overcome this weakening force of
  attraction.

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Melting Points for halogens
Element Melting Point (K)
Fluorine 85
Chlorine 238
Bromine 332
Iodine 457
Astatine 610
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Why does melting point increase going down the
halogens?

• The halogens are diatomic molecules, so F2, Cl2,
  Br2, I2
• As the molecules get bigger there are more
  electrons that can cause more influential
  intermolecular attractions between molecules.
• The stronger the I.A, the more difficult it will
  be to melt. (more energy needed to break the I.A)

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What are these I.A?

• van der Waals forces (London dispersion)
• Electrons are mobile, and although in a diatomic
  molecule they should be shared equally, it is
  found that they temporarily move and form
  slightly positive end and negative end.
• Now that one end is and the other -, there can
  be intermolecular attractions between the
  opposite charges of the molecules

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van der Waals forces
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• IB requires knowledge specifically for halogens.
  Check out this site for more detail.
• http//www.chemguide.co.uk/inorganic/group7/proper
  ties.html

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Period 3 melting point trends
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Explanation

• M.P rise across the 3 metals because of the
  increasing strength of the metallic bonds.
• Silicon has a giant covalent structure just like
  diamond which makes its structure remarkably
  strong and therefore takes more energy to break
  apart.

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• The atoms in each of these molecules are held
  together by covalent bonds (except Ar)
• They would have weak I.A affecting the amount of
  energy needed to melt them.
• Ar has extremely weak forces of attraction
  between its atoms, so its easiest to melt.

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3.3 Chemical properties

• 3.3.1 Discuss the similarities and differences
  in the chemical properties of elements in the
  same group.
• 3.3.2 Discuss the changes in nature from ionic
  to covalent and from basic to acidic of the
  oxides across period 3

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Reactivity of alkali metals

• Generally group 1 metals become more reactive as
  you go down a group.
• The valence electron of group 1 are found further
  from the nucleus as you go down the group.
• It is easier to remove an electron from francium
  than from lithium

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Alkali metal water

• Li(s) H2O (l) ? LiOH(aq) H2 (g)
• (Li and OH- in solution)
• The metal reacts with water to form the hydroxide
  of the metal (strong base) and bubbles off
  hydrogen gas.
• The larger the alkali metal, the more vigorous
  the reaction. Sometimes the H2 gas actually
  lights itself (exothermic reaction, releases
  heat) causing the H2 to burn.

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MUST KNOW!

• Na (s) H2O (l) ? NaOH (aq) H2(g)
• K (s) H2O (l) ? KOH (aq) H2(g)

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Alkali metals halogens

• 2Na (s) Cl2(g) ? 2NaCl (s)
• Halogens are good oxidizing agents, which means
  they cause electrons to be lost from another atom
  (the reducing agent)
• Halogens are 1 electron from stable octet and
  will try to remove electrons from valence
  electrons of other metallic atoms.

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MUST KNOW!

• 2K (s) Br2(l) ? 2KBr (s)
• 2Li (s) I2(g) ? 2LiI (s)

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Halogens reacting with halides

• Halogens want an electron and even will remove
  electrons from other soluble salts, we refer to
  as halides.
• When a salt dissolves it forms both of its ions
  in solution.
• Ex NaCl (aq) ? Na(aq) and Cl- (aq)
• So halides are easily available for reactions

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Done in aqueous systems

• Chlorine is stronger OA (oxidizing agent) than
  bromine because its found higher on the periodic
  table, so Cl2 will remove the electron from Br-,
  making Cl- and Br2
• Cl2 (aq) 2Br- ? 2Cl- Br2 (aq)
• Cl2 (aq) 2I- ? 2Cl- I2 (aq)
• Br2 (aq) 2I- ? 2Br- I2 (aq)

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Properties of Metals

• Shiny (lustre)
• Good conductors of heat and electricity
• Malleable and ductile (change shape and make
  wires)
• Tend to lose electrons
• Metal oxides form basic solutions in water (pH
  greater than 7)

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Properties of non-metals

• Brittle
• Poor conductors of heat and electricity
• Tend to gain electrons
• Non-metal oxides tend to be basic when dissolved
  in water (pH less than 7)

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Across Period 3 metallic to non-metallic oxides

• Basic solution from metallic oxide.
• Na2O(s) H2O (l) ? 2 NaOH (aq)
• MgO (s) H2O (l) ? Mg(OH)2 (aq)
• Hydroxides of group 1 and 2 generally considered
  strong.
• Acidic solution from non-metallic oxide.
• SO3(g) H2O (l) ? H2SO4 (aq)
• P4O10 (s) 6H2O (l) ? 4 H3PO4 (aq)
• Aqueous hydrogen involved with acidity

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Properties of metalloids

• Based on chemical and physical properties
• Tend to have semi-conductive properties and form
  amphoteric oxides.
• Considered metalloids are
• Boron (B)
• Silicon (Si)
• Germanium (Ge)
• Arsenic (As)
• Antimony (Sb)
• Tellurium (Te)
• Polonium (Po

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Amphoteric

• Behave as an acid or a base depending upon the
  reaction it is involved with.
• Also called amphiprotic (donate or accept a
  proton, H)
• Aluminums oxide is amphoteric.
• Al2O3(s) 3HCl (aq)? AlCl3 (aq) 3H2O (l)
• Reacts with a strong acid to make a to make a
  salt with water.
• Al2O3(s) NaOH (aq) ? NaAl(OH)4 (aq)
• Reacts with a strong base to form sodium
  aluminate