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Trends of the Periodic Table

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Nahreen M

on 29 June 2013

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Transcript of Trends of the Periodic Table

WHY IS THE PERIODIC TABLE
ARRANGED THE WAY IT IS?

TRENDS OF THE PERIODIC TABLE
ATOMIC SIZE
ATOMIC RADIUS DECREASES ACROSS PERIODS
ELECTRONEGATIVITY
CORE CHARGE
FIRST IONISATION ENERGY
First ionisation energy is the
minimum amount of energy
(kJ mol^-1) required to
remove the highest energy electron
from an atom (valence electrons)
PATTERNS WITHIN THE
PERIODIC TABLE

PERIODS VS. GROUPS
REFERENCES
The Periodic Table is arranged according to the
properties
of each element
Electron configuration, number of valence electrons (reactivity),
atomic number
and mass, the chemical and physical behaviour, etc all dictate the position of an element within the Periodic Table
The Periodic Table is divided into four blocks called
s, p, d
and
f
which reflects the subshell occupied by the valence electrons of an element

The blocks contain groups of elements with common names (s-block: group 1 contains
alkali metals
while group 2 contains
alkali earth metals
, p-block: groups 13-18 contains non-metals where group 17 contains
halogens
and group 18 contains
noble gases
, d-block: groups 3-12 contain
transition metals
, f-block:
lanthanides
and
actinides
)
The table is further split into
periods and groups

Grouping elements in these ways
allows predictions to be made about the physical and chemical properties
or
bonding behaviour
of an element based on its position in the Period Table
Periods =
rows
of the Periodic Table
The row or period can determine the number of
electron shells or orbitals
of an atom (e.g. 1st period = 1 shell, 2nd period = 2 shells, etc)
The maximum number of electron shells an atom can have is
seven
The number of occupied shells remain the
same
across a period so valence electrons are not pushed further away from the core
However, core charge
increases
along a period, drawing electrons in closer to the atom's core
This causes the atom to condense;
decreasing
in radius but
increasing
in density
ATOMIC RADIUS INCREASES DOWN GROUPS
The number of occupied shells
increase
where valence electrons are pushed further away from the core of the atom
However, core charge remains the
same
down group
This causes the atom to
increase
in radius but
decrease
in density
Groups =
columns
of the Periodic Table
The group can determine the number of
valence electrons
of an atom, although this does not apply to transition metals (e.g. Group 1 = 1 valence electron, Group 2 = 2 valence electrons, etc but group 17 = 7 valence electrons)
Atoms of the same groups have
similar chemical properties
There are
eighteen
groups in the Periodic Table
Both groups and periods are similar as they are designed to
organise elements
and be able to
demonstrate trends and properties of an element
based on their position in the Periodic Table
E.g. Reactivity of an element can be predicted where elements situated further left along a period and down a group are more reactive
Atomic radius is generally considered the
total distance
between the
nucleus
and the
outer most electron shell
or orbital of an atom
Determining the atomic radius can be difficult when various
bonds
between atoms are involved as the position of the outer most electron is not clear
Therefore, the radius is determined by the
distance between the nuclei of two atoms
However, this means there is
no exact radius
of an atom but is rather
dependant on the bond an atom forms
COVALENT RADIUS
IONIC RADIUS
METALLIC RADIUS
NON-METAL REACTIVITY
METAL & NON-METAL CHARACTERISTICS
METALLIC REACTIVITY
Covalent radii can be determined when
covalent bonds
are present
When two atoms of the same element are bonded this way, the covalent radius can be determined by
halving the distance between the two nuclei
whereas the whole distance gives the diameter
Covalent radii follow the
same trend as atomic radii
80pm 80pm
Ionic radii can be determined when
ionic bonds
or
ions
are present
As ionic bonds contain atoms of varying sizes, where cations are generally smaller in size while anions are much larger, the
distance between the two nuclei must be divided according to the respective sizes
of each atom
If we knew the atomic radius of atom A, we could determine the radius of atom B by substracting the radius of atom A from the total distance between the two atoms' nuclei
E.g. Compound CaSe has a total distance of 278pm between the nuclei of atoms Ca and Se, then the atomic radius of Ca would be
278pm (total distance) - 178pm (radius of Se) = 100pm (radius of Ca)
CATIONS HAVE SMALLER RADII THAN THEIR NEUTRAL ATOMS
Cations are atoms with a positive charge, meaning they have
less electrons than protons
Due to this
loss
of electrons, there are less electrons to pull towards the nucleus, resulting in a
stronger
pull

so that the atom is forced to condense and
decrease
in radius
ANIONS HAVE LARGER RADII THAN THEIR NEUTRAL ATOMS
Anions are atoms with a negative charge, meaning they have
more electrons than protons
Due to this
gain
of electrons, there are more electrons to pull towards the nucleus, resulting in a
weaker
pull

so the atom expands and
increases
in radius
Metallic radii can be determined when
metallic bonds
are present
The metallic radius is
half of the total distance between two nuclei
of adjacent atoms
As these groups of atoms will be of the same element, the distance of each atom will be the same
Metallic radii follow the
same trend as atomic radii
http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/atomic-radii.html
http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Atomic_Radii
http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Atomic_Radii
http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Atomic_Radii
University of California 2012, Atomic Radii, <http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Atomic_Radii> accessed 8 December 2012
University of California 2012, Periodic Trends, <http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Trends> accessed 12 December 2012
Lukins, N et al. 2006, Heinemann Chemistry 1, 4th edn. Pearson Education Australia, Australia
ENERGY REQUIRED INCREASES ACROSS PERIODS
The energy required to remove valence electrons
increases
as the strength of the attraction between the electrons and nucleus
increases
along a period
Elements at the beginning of a period (such as metals)
require less energy
to remove an electron as such elements
want to give up electrons
in order to be more stable
As you move along a period, elements (such as non-metals)
require more energy
to remove electrons as such elements
want to accept
rather than give up electrons to make a full shell and be more stable
Noble gases especially require a great amount of energy to remove a valence electron as they already have full shells/are stable
ENERGY REQUIRED DECREASES DOWN GROUPS
As atomic radius increases down groups, valence electrons are situated further from the nucleus and are
less affected by core charge
, hence easier to remove
http://ibchem.com/IB/ibnotes/brief/ato-hl.htm
Electronegativity is a measure of an
atom's ability to attract electrons
ELECTRONEGATIVITY INCREASES
ACROSS PERIODS
An atom's electron-attracting ability
increases
due to the
increasing core charge
as you move along a period
ELECTRONEGATIVITY DECREASES
DOWN GROUPS
Electrons are more weakly attracted to atoms as
valence electrons are more distant
due to the atoms increasing in size down a group
http://mageechem11.blogspot.com.au/2012/05/hello-d-today-well-look-at.html
Core charge is the
attractive force
between the
core of an atom
and its
valence electrons
The core charge of an atom can be calculated by
subtracting the number of non-valence electrons from the number of protons
so the charge will always be positive
CORE CHARGE INCREASES
ACROSS PERIODS
The number of protons in the nucleus of an atom
increases
across a period (where the number of non-valence electrons is generally the same) so the overall positive charge/core charge
increases
CORE CHARGE STAY THE SAME
DOWN GROUPS
Elements of the same group have the
same number of valence electrons
, which essentially gives the core charge of the atom
Valence electrons
do not feel the full attraction
of the nuclear charge as the two inner electrons act as a
shield
which is why the core charge does not change down a group despite the increasing distance between valence electrons and the core of an atom
Reactive metals are those which readily react with other elements to
give up
their electrons (valence) to form
positive ions
REACTIVITY DECREASES ACROSS PERIODS
REACTIVITY INCREASES DOWN GROUPS
As first ionisation energies
increase
across periods, the strength of the attraction between the outer electrons and nucleus
increases
so it is
more difficult
to remove these electrons
As first ionisation energies
decrease
down groups, the strength of the attraction between valence electrons and the nucleus is much
weaker
so it is
less difficult
to remove these electrons
Reactive non-metals are those which readily react with other elements to
accept
electrons to form
negative ions
REACTIVITY INCREASES ACROSS PERIODS
REACTIVITY DECREASES DOWN GROUPS
As electronegativity
increases
across periods, the atoms' electron-attracting abilities
increase
so they are able to accept electrons
more easily
As electronegativity
decreases
down groups, the atoms' electron-attracting abilities
decrease
so they accept electrons
less easily
Metals typically display qualities such as

solid at room temperature
conducts electricity/heat
shiny
malleable/ductile
Non-metals display qualities such as

usually gas, liquid or low-melting solid at room temperature
poor conductors of electricity/heat
dull in colour/lustre
brittle
METALLIC CHARACTER DECREASES ACROSS PERIODS & INCREASES DOWN GROUPS
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