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 z=20.
3.1.4 Apply the relationship between the highest occupied energy
level for an element and its position in the periodic table.
 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
Atomic Size
 The electron cloud doesn’t 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/essentialch
emistry/flash/atomic4.swf
Group trends
 As we go down a group
(each atom has another
energy level) the atoms
get bigger, because more
protons and neutrons in
the nucleus
H
Li
Na
K
Rb
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.
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
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.
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.
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.
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.
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.
Li+1
N-3
B+3
Be+2
C+4
O-2
F-1
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.
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.
Ionic Radii
http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/medialib/media_portfolio/text_images/CH08/FG08_13.JPG
Ionic Radii: Trends
 What is the pattern in Periods?
 Smaller until 5A.
 What is the pattern in Groups?
 Gradual increase in size.
Size of Isoelectronic ions
 Positive ions have more protons so they are smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
Ionization Energy
http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/images/ionization-energy.jpg
Ionization Energy
http://images.google.com/imgres?imgurl=http://www.webelements.com/webelements/properties/media/tables/cityscape-x/ionization-e
nergy-1.jpg&imgrefurl=http://www.webelements.com/webelements/properties/text/image-cityscape/ionization-energy-1.html&h=1365&
w=2048&sz=808&tbnid=lq6myNPljopuCM:&tbnh=100&tbnw=150&prev=/images%3Fq%3Dionization%2Benergy&start=1&sa=X&oi=im
ages&ct=image&cd=1
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.
nd
2
and
rd
3
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.
Ionization Energies in kJ/mol
1
2
3
4
5
6
7
8
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
H
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.
Ionization Energy
 What happens when sodium loses an electron?
 What is its electron configuration?
 Na 1s22s22p63s1
 Na+1 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.
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/intro_chemistry/5e/students/protected/periodictables/pt/pt/table/t_e2.gif
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.
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.
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
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.
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.
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
Electronegativity Trends
Overall Trends!
http://campus.ru.ac.za/full_images/img05206111510.jpg
Melting Points of Group 1
Element
Melting Point (K)
Li
453
Na
370
K
336
Rb
312
Cs
301
Fr
295
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
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.
Melting Points for halogens
Element
Melting Point (K)
Fluorine
85
Chlorine
238
Bromine
332
Iodine
457
Astatine
610
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)
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
van der Waals forces
 IB requires knowledge specifically for halogens.
Check out this site for more detail.
http://www.chemguide.co.uk/inorganic/group7/propert
ies.html
Period 3 melting point trends
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.
 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.
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
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
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.
MUST KNOW!
 Na (s) + H2O (l)  NaOH (aq)+ H2(g)
 K (s) + H2O (l)  KOH (aq)+ H2(g)
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.
MUST KNOW!
 2K (s) + Br2(l)  2KBr (s)
 2Li (s) + I2(g)  2LiI (s)
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
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)
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)
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)
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
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
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+)
 Aluminum’s 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