Periodicity
Topic 3 Chapter 3 page 75
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 electron arrangement of element and
their position in the periodic table up to Z = 20.
3.1.4
Apply the relationship between the number of electrons in the
highest occupied energy level for an element and its position in the
periodic table.
THE PERIODIC TABLE – THE
CHEMIST`S , MOST IMPORTANT
TOOL
Periodic Table
Metals and non-metals are separated by the zigzag line on most
periodic tables.
 Metallic behaviour changes gradually among the elements; the
most metallic acting element is Fr on the bottom left of the
periodic table.

•Periodicity
•
•
Using the simple metal/non-metal classification is
possible but it is really too general to be of much
use.
When the elements are ordered by increasing
atomic number a repeating patterns of physical and
chemical properties known as periodicity can be
observed.
The Periodic Table 1871
• This
was first discovered by Dmitri Mendeleev
(Russian) who is considered the father of the
periodic table because, unlike those before him:
•
•
he included 63 known elements and left spaces or gaps
which he predicted would be filled in one day by newly
discovered elements (three of these gallium, scandium and
germanium were found in his lifetime)
his table allowed for the accommodation of future findings
including noble gases (new column in 1884-1901), the rare
earth elements, Bohr atom and electronic structure in 1913
(energy levels), and the discovery of synthetic elements 1939 to present
•Groups vs. Periods
•
•
rows are called periods
• the period number = number of electron shells
columns are called groups or families
• the group number = number of valence electrons
•
groups have similar chemical and physical properties.
IB numbering system1-7,0
Activity
Arrangement of Groups and Periods
•Important Groups
–Colour on your periodic table
Group 1- Alkali metals
Group 2- Alkaline Earth Metals
Group 3 – Boron Family
Group 4 – Carbon Family
Group 5 – Nitrogen Family
Group 6 – Oxygen Family
Group 7 – Halogens
Group 0 – Nobel Gases
Assigned Work
•
•
Use the class notes Periodicity to complete Introduction to
Matter and Periodicity Practice Questions
Response to Video
3.2
3.2.1
3.2.2
3.2.3
3.2.4
Physical Properties
Define the terms first ionization energy and electronegativity.
Describe and explain the trends in atomic radii, ionic radii, first
ionization energies, electronegativities, and melting points for
the alkali metals (Li → Cs) and the halogens (F→I).
Describe and explain the trends in atomic radii, ionic radii, first
ionization energies, and electronegativities for elements across
period 3.
Compare the relative electronegativity value of two or more
elements based on their position in the periodic table.
Five Physical Trends
For each of the following be able to i) define the term ii) state the
trend (if there is an decrease or increase) as you go across a
period and down a group, and iii) explain why the trend occurs.
1. atomic radii
2. ionic radii
3. first ionization energies
4. electronegativities
5. melting points
1. Atomic Radii
•
•
Atoms do not have definite boundaries; the boundary is where
you are likely to find the outer electrons that are moving
around the nucleus. Since this is a moving target...
We measure the covalent atomic radius (the distance of half
way between the nuclei of two atoms in a diatomic molecule).
Atomic Radii Trends
•
Within each period, the atomic radius tends to decrease with increasing
atomic number.
•
Within each group, the atomic radius tends to increase with the period
number in part because there are more layers (energy levels) and in part
because shielding (next slide) increases.
•
Thus, the largest atom is cesium and the smallest fluorine.
Nuclear Charge and Shielding
•
Shielding electrons are the electrons in the energy
levels between the nucleus and the valence electrons. They are
called "shielding" electrons because they "shield" the valence
electrons from the force of attraction exerted by the positive
charge in the nucleus.
shielding
electrons
valence
electron
Nuclear Charge and Shielding
•
As you go down a group the shielding effect increases since there are more
energy levels (shields). This results in the outer electrons being held less
tightly.
•
The strength of the attraction between the nucleus and the outermost
electrons helps determines the atomic radius. The stronger the attraction,
the smaller the size of the atom. The weaker the attraction (such as from
shielding), the larger the atom or ion.
Shielding Analogy- A concert
Nuclear Charge and Shielding
•
Effective nuclear charge (ENC) is the charge felt by the valence electrons
after you have taken into account the number of shielding electrons that
surround the nucleus. Look at your periodic table.
ELEMENT
nuclear charge
shielding
(number of protons) electrons
valence
electrons
ENC
protons-shielding e-
Li
3
2
1
3-2 = +1
O
8
2
6
8-2 = +6
F
9
2
7
9-2 = +7
Ne
10
2
8
10-2 = +8
•
As you go across a period, the number of protons and effective nuclear
charge (i.e. pull on valence e-) increases while shielding does not change.
This results in outer electrons being pulled in closer to the nucleus and
smaller atoms being found on the right side of the periodic table.
•
Nuclear charge and shielding must be considered together.
Nuclear Charge and Shielding
•
•
As you go down a group, the increase in the nuclear charge is
cancelled out by the increase in shielding electrons and the effective
nuclear charge (ENC) stays the same.
The extra layers of electrons result in the outer electrons being
further away from the nucleus and the lack of a stronger ENC to pull
them closer to the nucleus results bigger atoms being found at the
bottom of a group
ELEMENT
•
nuclear charge
(number of
protons)
shielding
electrons
valence
electrons
Effective Nuclear
Charge
protons-shielding e-
F
9
2
7
9-2 = +7
Cl
17
10
7
17-10=+7
Br
35
28
7
35-28 =+7
As ENC ↓ size ↑; as ENC ↑ Size ↓
2. Ionic Radii Trends
•
•
Cations are always smaller than their neutral atoms. Loss of
outer electrons results in an increase in nuclear charge on the
remaining ones.
Anions are always larger than their neutral atoms. Effective
nuclear attraction does not change for an increased number of
electrons.
11-10= +1
11-2 = +9
(now there
are 8 valence)
17- 10 =+7
17-10 = +7
(but there is one
more electron in valence)
Ionic Radii Trends cont’d
•
•
•
Across a period the cations
decrease in radius.
With each increase in charge,
(i.e. Na1+, Mg2+, etc.), the
atoms contract because of
the greater attractive force of
the nucleus.
When you reach the nonmetals (anions) there is an
abrupt increase in radii and
then the radius again
decreases.
3. Ionization Energy
•
•
•
•
The minimum energy needed to remove the outermost electron from a
neutral atom in the gaseous state is called the first ionization energy.
Equation for first ionization energy:
Ca(g) + energy (E1) —> Ca+ (g) + e‾
As you go across a period the ionization energies increase (effective
nuclear charge is increasing but the shielding effect is constant, so more
energy is needed to pull the electrons from the atom).
Nobel gases do not form ions. They have a full outer valence of 8 e- and
are stable.
4. Electronegativity
•
•
•
The electronegativity of an element is the tendency for an atom to
attract a bonding pair of electrons to itself when it is chemically
combined (covalently bound) with another element.
It is a measure of how strongly an atom can attract an electron to itself.
Electronegativities are expressed in arbitrary units on the Pauling
Electronegativity scale (1-4). We will see more of electronegativity in the
next unit.
Electronegativity Trends
•
Electronegativity increases going across the period as the attraction of the
nucleus to the electrons increases, and decreases down each group.
•
F is the most electronegative element and Cs is the least.
Summary
Summary
•atomic radius decreases
• ionic radius decreases
• ionization energy increases
• electronegativity increases
• metallic character decreases
INCREASING
EFFECTIVE
NUCLEAR
CHARGE
5. Melting Points Period 3 Trends
•
•
•
Physical properties, such as melting point, boiling point and density
depend on the nature of the bonding between the particles of the
element which we will look at in the next section.
The more metallic part of the period will have a higher
melting/boiling point due to stronger bonds.
The nonmetallic tend to have a lower melting/boiling point but the
pattern is more complex.
Melting Points/Boiling Points Trends
Left
• At the left of the period the elements (Li, Na, K) are solid metals
whose atoms are held together tightly by strong bonds (metallic
bonds). As you go across the across the period table, the strength of
the metallic bonding increases which results in increasing melting
points from left to right
Melting Points/Boiling Points
Centre
• At the centre of period 2 and 3, we find C and Si, giant covalently
bonded structures which have the strongest bonds of all substances
(Cn is a diamond), and therefore the highest melting points.
Melting Points/Boiling Points
Right
• Following C and Si, the melting points suddenly drop
as the elements (N, P) have forces holding one molecule
(eg. Cl2) to another molecule (another Cl2).
• Melting points are low, and depend on mass and size of the
molecules P4, S8, or Cl2. This is further emphasised with the
noble gases (He, Ne, Ar), which exist as single atoms with
very, very weak forces between the gas molecules and very
low melting points.
Assigned Work
•
•
Work on Physical Periodic Trend Practice Questions
Spreadsheet Lab-40 min
|
PERIODIC TRENDS OF
ALKALI METALS,
HALOGENS, AND PERIOD 3
ELEMENTS
3.3 Chemical Properties
3.3.1
Discuss the similarities and differences in the chemical properties of
elements in the same group.
Know the following reactions:
•
Alkali metals (Li, Na, and K) with water
•
Alkali metals (Li, Na, and K) with halogens (Cl2, Br2 and I2)
•
Halogens (Cl2, Br2 and I2) with halide ions (Cl¯, Br¯, and I¯)
Review – Balancing Chemical
Equation
•
•
•
•
•
See handbook on line under General Info
This slide show is on blog
Worksheet Balancing Chemical Reactions
Handout Classifying Chemical Reactions (focus on single and
double displacement and writing ionic equations)
Chemical Reactions Lab
Alkali Metals
•
•
•
•
•
Most metallic character
Highly reactive (only have 1 outer electron to lose)-therefore
not found in nature in elemental state; reacts with O2, to form
oxide coating and H2O to form basic (alkali) solutions.
Increase in reactivity down the group (Fr is most reactive metal).
Softest of all metals (can be cut with a knife) and decrease in
melting point down the group.
Decrease in density down the group (atoms are larger and less
can be packed in a space).
Alkali Metals Chemical Trends
1.
React exothermically (give off heat) with H2O to form a metal hydroxide
solution and hydrogen gas:
•
e.g. 2Na (s) + 2HOH(l) →2NaOH(aq) + H2(g)
•
These elements are called alkali metals because they produce OH¯ when
combined with H2O which makes resulting solution alkaline (basic) NaOHaq)
→ Na+(aq) + OH-(aq)
2. alkali metals are donate electrons while halogens accept electrons. Therefore
these metals combine directly with halogens (Cl2, Br2 and I2), to form ionic salts).
3.
•
2Na(s) + Cl2(g) → 2NaCl(s)
•
2K(s) + Br2(l) →2KBr(s)
react to O2 in air to form a solid, white or greyish
surface coating of metal oxide
4Li(s) + O2(g) →2Li2O(s)
Rust, corrosion and oxides
•
•
Al is an extremely important metal since it is light and
strong (can be used to make airplanes) and cheap enough
to be found in every kitchen.
Its big advantage over steel (also very strong but contains
iron) is that is does not rust. It is extremely reactive with
air (even more so than iron) but unlike iron (III) oxide
(rust), aluminum oxide is a tough protective coating that
is harder than the Al itself and prevents the metal from
further corrosion.
Al(s) + 3O2(g) →2Al2O3(s)
Rust, corrosion and oxides
•
Fe2O3 is a red flaky powder that soon falls off and exposes the metal
so fresh corrosion until the Fe is `rusted away`.
4Fe(s) + 3O2(g) → 2Fe2O3(s)
•Halogen Properties
•
•
•
•
•
Typical non-metal character
Most reactive of non-metals (only need to gain 1 electron to fill outer
shell) - not found free in nature (always found combined with other
elements or is diatomic).
Decrease in reactivity down the group (F2 is most reactive non-metal).
Increase in melting point down the group (go from gas → liquid →
solid) because forces holding diatomic molecules together are
stronger.
Distinctive colours, toxic and corrosive
Halogen Colours
View Periodic Videos
• F2
yellowish gas
• Cl2 greenish
• Br2
• I2
gas
reddish brown liquid
purplish black solid (SATP)
-gas is violet
Reactions of Halogens
1. Form ionic compounds by combining with metals e.g. NaF
Halogen means salt forming.
2. Form covalent bonds and molecular compounds by binding
with other non-metals e.g. F2O
3. Halogens with halogen salts -a halogen higher in group 7 will
displace a halogen ion that is lower down the group from
solution. (i.e. halogens lower down are less reactive than
those above them).
Cl2(g) + 2NaBr(aq) →Br2(l) + 2NaCl (aq)
Br2(l) + 2NaCl(aq)→ no reaction
HALOGEN
Br2
HALOGEN
ION Cl-
Reactions of Halogens
4. Reactions between halide ions and silver solutions;
formation of silver halides (AgX)
Ag+ (aq) + X ¯(aq) → AgX (s)
where X = Cl, Br, or I
AgNO3 (aq) + NaCl(aq) → AgCl (s) + Na(NO3)(aq)
•
spectator ions
Colours of silver halides:
Silver halides are used
in black and white film
photography.
•
AgCl
white
AgBr
cream
AgI
yellow
Summary of some halogen reactions
•Overview of Alkali Metals and Halogens
Alkali Metals
•
•
•
•
•
Most metallic character
Highly reactive (only have 1 outer
electron to lose)-therefore not
found in nature in elemental
state; reacts with O2, to form
oxide coating (tarnish) and H2O to
form basic (alkali) solutions.
Increase in reactivity down the
group (Fr is most reactive metal).
Softest of all metals (can cut with
knife) and decrease in melting
point down the group because
atoms are larger (Cs is largest
atom)
Decrease in density down the
group (atoms are larger and less
can be packed in a space)
Halogens
•
•
•
•
•
Typical non-metal character
Most reactive of non-metals (only have 1
electron to gain to have 8 in valence) - not
found free in nature (combined with
other elements or diatomic). Halogen
means salt former; form salt when they
bond to alkali metals.
Decrease in reactivity down the group (F2
is most reactive non-metal).
Increase in melting point down the group
because of forces holding diatomic
molecules together are stronger.
Distinctive colours, toxic and corrosive:
•
•
•
•
F2: yellowish gas
Cl2: greenish gas
Br2: reddish liquid
I2: purplish, black solid (sublimes to purple
gas)
•Alkaline Earth Metals and Nobel Gases
Group 2- Alkaline Earth
Metals
•
Similar to the Alkali
metals but less reactive
(but still too reactive to
be found in elemental
state in nature).
Group 0 -Nobel Gases
•
•
•
All colourless, odourless
gases at SATP
Least reactive of all
elements; only the
heavier ones (Kr, Xe Rn)
react with F2 under
controlled conditions
All except radon (which
is radioactive) are nontoxic.
•Assigned Work
• Worksheet
Chemical Properties of Alkali Metals
and Halogens
• Chemical Periodicity Practice MC
• Chemical Reactions Lab
• Activity Series Lab