Download Trends of period 3

Periodicity
Objectives:
1.Explain the reason behind the
arrangement of elements of the periodic
table.
2. Show the link between atomic structure,
oxidation state and the position of elements.
3. Predic the properties of unknown
elements.
4.Identify trends of Group II, VII.
5.Identify trends of Group 3, specifically
metallic and non-metallic characteristics.
The Periodic table:
There are 118 elements of the Periodic table which are
arranged by increasing atomic number. These elements are
presented as vertical columns or Groups and horizontal
columns or rows.
Groups contain the same
number of outer shell electrons
Be 2,2
Mg 2,8,2
Ca 2,8,8,2
Sr 2,8,18,8,2
Ba 2,8,18,18,8,2
All elements of the same Group (vertical column)
have the same outer shell configuration.
Groups contain the same
number of outer shell electrons
Elements with the same number of outer shell electrons
have similar properties and behaviour. E.g all elements of
group 1 are highly reactive metals. Elements of group 7 are
highly reactive non-metals.
Elements of the same group
have similar properties
Lithium
Sodium
Potassium
Rubidium
All of these elements are highly reactive metals.
Trends of Group IIA:
1. Atomic radius
2. Ionization energy
3. Melting Point
4. Boiling Point
5. Relative density
Group 2 trends:
1. Atomic Radius
Atoms have electrons arranged as rings or
levels at fixed distances from the nucleus.
Each ring or shell holds a specific number of
electrons and the greater the numbers of
electrons, the more occupied shells there are
and the larger the radius of the atom.
-Atomic radius increases down the group.
And decreases left to right across the period.
2. Ionization energy:
Decreases down the group
Atoms consist of a nucleus whose positive
charge holds negatively charged electrons
at specific levels or shells around the
nucleus much like the moon orbiting the
earth . The farther from the nucleus an
electron is the weaker the attractive force
exerted by the nucleus upon that electron
and the less work or energy it would take to
liberate that electron from the atom. The
amount of energy it takes to remove an
electron from an atom is called ionization
energy. As the atoms become bigger lower
down the group their outer electrons move
farther from the nucleus and the ionization
energy of that element decreases.
3. Melting Point & Boiling point:
Increases Up the group
To melt a substance the forces pulling one molecule or atom
from another Should overcome the attractive forces (bonds)
holding the molecules together. As atoms become larger
they possess increasing amounts of Protons and the
attraction between the nucleus of one atom and the shells of
electrons around neighbouring atoms becomes stronger and
so elements with high atomic numbers are usually liquids
and solids at room temperature. This also applies to boiling.
-Melting and Boiling point increases Up the group.
3. Melting Point & Boiling point:
decreases down the group
4. Density:
Increases down the group
Be
Ra
`
*electrons are not shown.
Moving from top to bottom along group 2
represents and increase of atomic number
and mass of the elements, the atom of each
element having a bigger nucleus with more
protons and neutrons than the element
above it. The amount of matter that each
atom contains increases going down the
group. Materials made up elements lower
down the group will contain more mass pe
unit volume than members of the group
higher up, as each of their atoms contains
Why does an increase of
more protons and neutrons.
E.g
cubic
electrons
not 1count
meter of Beryllium weighs
tonnes
towards1.85
the density
of anwhile
element?
1 cubic metre of Radon weighs
5.5 tonnes.
Introduction to Group VIIA:
The halogens
Group
Fluorine
Chlorine
Bromine
Iodine
The members of group 7 are
a highly
reactive
class
Members
of group
7 lower down
the of
group
are solids,
those“salt
such asformer”.
Bromine
elements called halogens, from
thewhile
Greek
near the middle are liquid and Fluorine &
Due to their reactivity theseChlorine
elements
near theare
top never
are gasesfound
at room
Explain
this graded change
of
naturally as pure deposits temperature.
but exist
as
compounds
with
state from solids to gas?
other elements or as ions. Members higher up on the group
form Diatomic gases at room temperature, while Bromine
near the middle is liquid and Iodine near to the bottom is a
solid as room temperature.
Trends of Group VII:
1.Atomic radius
2. Ionization energy
3. Melting Point
4. Boiling Point
5. Density
1.Atomic radius
Similar to that of Group 2,
atomic radius increases down
the group. Fluorine is the
smallest member of Group 7 with
Iodine and Astatine being the
largest.
2. Ionization energy
Ionization energy decreases going down the group
from Fluorine to Iodine and Astatine.
The trend of ionization energy is also similar to that of Group II. However the
ionization energies of Group 7 are much larger than those of group II. This is because
ionization energy is a measure of the amount of energy it would take to remove
electrons from an atom. However members of group 7 have outer shells with 7
electrons it is far easier to lose 1 or 2 electrons than it is to lose 7 and so more energy
or work is done to remove an electron from a group 7 element. However it is still easier
to remove the electrons from larger atoms than from smaller ones and so the general
trend remains the same: Ionization energy decreases going down the group from
Fluorine to Iodine and Astatine.
Element:
Ionization Energy:
Fluorine
1690
Chlorine
1260
Bromine
1150
Iodine
1020
Melting and Boiling point:
Group 7 shows identical trends to group 2 with
melting and boiling points increasing going down the
group.
Boiling
point and
melting
point
increase
down the
group
Members of group 7 lower down the group
are solids, while those such as Bromine
near the middle are liquid and Fluorine &
Chlorine near the top are gases at room
temperature. Explain this graded change
of state from solid to gas?
Density:
Decreases down the group
Identical to that of Group 2 Density of materials
made of group 7 elements increases with elements
lower down the group than those above.
• Density increases down the group.
Introduction to Period 3
Trends of period 3:
Period 3 contains the three different types of
elements: metals, metalloids and non-metals
and refers to the elements left to right from
Sodium to argon.
Trends of period 3:
1.Atomic radius
2.Electronegativity
3.Metallic properties
4.Reactivity
5. Bonding
6.Melting point
7. Boiling point
8. Density
1.Atomic radius:
Each element of period 3 has one more proton
and one more electron than the one before it.
Each extra electron going to the outer shell of
the element while the proton resides at the
nucleus. As one goes across the period the
nucleus experiences an increase of positive
charge while it’s electrons are still at the same
fixed distance away. This increase of nuclear
positive charge without an increase of distance
between the nucleus and electrons causes the
atomic radius to shrink as one goes across the
period. T’s as if the nucleus is pulling the atom
together harder because of the extra “pulling”
force of the protons upon the orbiting electrons.
So the elements at the beginning of period 3 are
the largest while those at the end are the
smallest.
2.Electronegativity:
Electronegativity is the ability of an element to attract electrons to itself and is a balance between 2
factors:
1. The size of the nucleus and
2. The number of occupied electron shells.
The positive charge from the protons of the nucleus of each element emanates outwards to attract
negatively charged bodies. This force however is decreased by
1. The distance over which it acts; the longer the distance between the nucleus and the electron the
weaker the attraction.
2. The Number of occupied shells that the atom has. Each electron orbiting the atom is capable of
absorbing the positive attractive force produced by the nucleus. Much of the attraction is able to
propagate to the outside of the atom but where ever an electron intersects with the positive rays
from the nucleus the positive attractive force is absorbed. . An effect referred to as “shielding” as the
occupied shells serve to shield outside electrons from the pull of the nucleus.
Very large atoms with large distance between electrons outside the nucleus and many occupied
shells of electrons absorbing positive force from the nucleus are so well shielded that they often will
not accept an electron unless forced. Additional electrons do not affect the positive forces coming
from the nucleus as much as the addition of entire shells.
2.Electronegativity:
electron
Both Chlorine and Sodium have the same numbers of shells, smaller atomic
radius. However the Chlorine atom is more electronegative element due to 1.
Larger Nucleus with more protons.
Electronegativity is a balance between the pull of positively charged Protons,
Shielding effect of electrons and the distance between the nucleus and electrons
outside.
2.Electronegativity:
Periodic table showing electronegativity values of all elements(Pauling scale).
Electronegativity increases up the group and left to right
across the period. Fluorine is the most electronegative
element.
3.Metallic properties
All metals react by losing electrons to produce positive ions, the more readily an ion loses
electrons, the more metallic it's nature. The first 3 elements of period 3 consist of sodium,
magnesium and Aluminum, which have 1,2 and 3 valence electrons respectively. it is
easier to lose 1,2 or 3 electrons than to gain 7,6 or 5 electrons so these elements react by
losing electrons. This behaviour means that these elements are true metals. Silicon which
has 4 valence electrons ad neither loses nor gains but shares electrons as covalent bonds.
Phosphorous ,Sulphur, Silicon and Chlorine readily gain electrons instead of losing them
and are the least metallic elements of period 3.
Metallic nature decreases from right to left along the period 3.
4.Reactivity
The reactivity of an element is defined as how readily an element will partake of chemical
reactions. Chemical reactions take place when electrons are lost, gained or exchanged.
Metals react by losing electrons, so metals which have a very weak hold on the outer
electrons react readily.
Nonmetals react by gaining electrons, so non-metals whose nuclei have a very strong
attraction to electrons are very reactive.
Several factors affect the reactivity of an element
1. Electronegativity
2. Atomic radius
3. Chemical nature: whether metal or non metal.
Metals with 1 valence electron lose their electrons to chemical reactions more easily than
those with 2 or 3 electrons, so reactivity of metals increases from right to left.
Reactivity of non-metals also varies based on the electronegativity of the element, the
more however the reactivity increase from left to right instead.
4. Reactivity:
• Reactivity of metals increase from right to left
across period 3, sodium being the most reactive.
• Reactivity of non-metals increases from right to
left across period 3, Chlorine being the most
reactive.
Reactivity of metals
Reactivity of Non-metals
4. Reactivity:
Question: Sodium explodes on contact with air and is
stored under mineral (paraffin) oil, while magnesium and
aluminium can be stored safely using an ordinary reagent
bottle. Why?
Question: Fluorine is so reactive that when stored with
glass containers it begins to react with the silicon of the
vessel walls. explain Fluorine's highly reactive nature.
5. Bonding:
Na 2,8,1
Mg 2,8,2
Al 2,8,3
Si 2,8,4
P 2,8,5
S 2,8,6
F 2,8,7
Ar 2,8,8
Metals react by losing electrons and are said to be
Electropositive. Metals of group 1 and 2 lose electrons to form ions
with a charge of +1 or +2. When an element bonds with metals of
group 1 or 2 it accepts the electrons which the metal loses and it
itself becomes negatively charged. The oxidation state of an
element is equal to it’s charge so elements which are bonded to
group 1 or 2 metals usually have negative oxidation states.
When non-metals of group 6 or 7 bond react they gain
electrons and become negatively charged taking on oxidation
states of -2 or -1 respectively. When 2 non-metals bond the more
electronegative partner carries the electrons and has a negative
oxidation state, while the less electronegative partner has a
positive oxidation state. The exception is Fluorine which always
has a negative oxidation state of -1. Oxygen also always has a
negative oxidation state of -2 except when bonded to fluorine.
Elements of Group 4 and 5 may have positive or negative
oxidation states according to what it they are bonded to. When
bonded to elements less electronegative than themselves, e.g
metals, they have negative oxidation states as they accept the
electrons from the metals. When bonded to more electronegative
elements, e.g group 6 and 7 non-metals, they lose electrons and
have a positive oxidation state.
5. Bonding:
Question: No matter which
element is bonded to fluorine,
fluorine will always have a
negative oxidation state, while
it’s partner will take on a
positive oxidation state. Why?
6.Melting Point & 7.Boiling Point:
Na
Mg
Al
Si
P
S
Cl
Ar
Melting
Point
98
650
660
1423
44
117
-101
-189
Boiling
Point
890
1120
2447
2680
280
445
-34
-186
Both the melting and Boiling point increase
to a maximum at silicon and then decrease
with the exception of Sulphur.
Silicon has 4 valence
electrons, why is it the
element with the highest
melting and boiling point?
8: Density
Density
(gcm-3)
Na
Mg
Al
Si
P
S
Cl
Ar
0.97
1.74
2.70
2.33
1.82
1.96
2.99x10 -3
1.66x10-3
Density follows the same trend as melting
and boiling points. Increasing up to silicon
and then decreasing to argon, with the
exception of sulphur.
Physical Properties of period 3
elements:
Element
Atomic
Number
Description
Sodium (Na)
11
Soft metal; low density; very/reactive
Magnesium
(Mg)
Aluminum
(Al)
Silicon (Si)
12
14
Harder metal than sodium; low density; less reactive than
sodium
Metal as hard as magnesium; low density less reactive than
magnesium
Brittle non metal; not very reactive
Phosphorus
(P)
Sulphur (S)
15
Non-metal; low melting point; white solid; reactive
16
Non-metal; low melting point; yellow solid; moderately reactive
Chlorine (Cl)
17
Non-metal; green gas; extremely reactive .
Argon (Ar)
18
Non-metal; colorless gas, unreactive
13
Oxides of Group 3 Elements:
When dissolved with water the oxides of group 1 and group 2 metals produce Bases, Aluminium
oxide is Amphoteric, and when dissolved reacts with both acids and bases. Silicon oxide reacts
with bases at high temperature. The oxides of Phosphorous are acidic and react with bases to
form silicates. All sulphur oxides are acidic , Sulphuric acid is a solution of Sulphur trioxide.
Chlorine Oxides are very rare.
Sodium Oxide - Basic
Magnesium Oxide - Basic
Aluminium Oxide -Amphoteric (reacts with both acids and
bases)
Silicon Oxide - Basic ( reacts only at high temperature)
Phosphorus Oxide - Acidic
Sulphur Oxide - Definitively acidic
Chlorides of some group 3
elements:
Sodium Chloride and Magnesium Chloride are
ionic compounds, when molten or dissolved
with water they conduct electricity. All other
Chlorides are Co-valent.