Transcript Explaining Periodic Trends

Explaining Periodic Trends
Atomic Radius
Atomic Radius – The distance from the centre of an atom to the
boundary within which the electrons spend approximately 90% of
their time.
 The size of an atom is usually reported in terms of its atomic radius.
Atomic Radius
To obtain the atomic radius of an atom, the centre – to – centre
distance is measured and then divided by 2.
Atomic Radius Trends Within a Period
 The size of atoms decreases going left to right across a period.
WHY?
As you move left to right on the periodic table:
 The number of protons (positive charges) in the nucleus increases.
 The number of electrons (negative charges) increases as well, but
each additional electron is added to the same outer shell.
 The increase of both the negative and positive charges increases the
overall attractive forces between the nucleus and the outermost shell.
Thus, the outermost shell gets pulled in closer to the nucleus and the
atom is smaller than the one before it.
Atomic Radius Trends Within a Group
 The size of atoms increase as you move down a group in the
periodic table.
Why?
 The atomic number is higher (more protons in the nucleus) as you
move down a group.
 The number of electron shells increases as well.
 Thus, the inner shells shield the increase in positive charges form the
nucleus. This is called effective nuclear charge – the outermost
electrons do not exhibit as strong an attraction to the nucleus as they
would if the inner shell electrons were not there.
Effective Nuclear Charge Example
Atomic Radii Trends
Ionization Energies
Ionization Energy – The amount of energy required to remove
the outermost electron from the atom or ion in the gaseous state.
Example:
A(g) + energy  A+(g) + e- (first ionization energy)
 After one electron has been removed, it is still possible to remove
more electrons. This leaves an ion with a larger positive charge.
Example:
A+(g) + energy  A+2(g) + e- (second ionization energy)
First Ionization Energy Trends
Within a Period
 There is an increase in first ionization energies of atoms
going left to right across a period.
WHY?
 As you move across a period, the outermost electron shells get
closer to being filled (very stable). Thus, it will take more energy
to remove them. Ex. The noble gases have outer electron shells
completely filled and are very stable.
First Ionization Energy Trends
Within a Group
 There is a decrease in first ionization energies of atoms
moving down a group.
Why?
 The atomic radius increases (by an electron shell) as you go down a
group. Thus, the outermost electrons will be farther away from the
positive nucleus. This results in less attractive “pull” by the nucleus
and the outermost electron.
 Less attraction means that there will be less energy required to
remove the outermost electron the larger the atom.
First Ionization Energies of
Elements
Electron Affinity
Electron Affinity – The energy absorbed or released when an
electron is added to a neutral atom.
Example: A(g) + e-  A-(g) + energy
 Those atoms that release energy when they gain an electron have
electron affinity values that are negative. These atoms form ions that
are more stable than the atom (the more negative the value, the more
stable the ion).
 Those atoms that absorb energy when they gain an electron have
electron affinity values that are positive. . These atoms form ions that
are less stable than the atom (the more positive the value, the less
stable the ion).
Electron Affinity General Trends
 The general trend is that electron affinity values become
increasingly negative as you move from left to right across a
period (NOTE: the noble gases DO NOT form ions. Thus,
they ALL have positive values).
 The general trend is that electron affinity values tend to become
decreasingly negative as you move down a group in the periodic
table.
Electron Affinity General Trends
Electronegativity
Electronegativity – An indicator of the relative ability of an atom to
attract shared electrons.
 The atom must be bonded to another atom to make this
determination.
 Electronegativity is relative and, therefore, has no units.
Electronegativity Trends Within a
Period
 There is an increase in electronegativity values in atoms going
left to right across a period.
WHY?
 The atomic radius/size gets smaller as you move from left to right in
the periodic table.
 Thus, the positive nucleus of one atom is closer to the outer
electrons of the other atom that is involved in a covalent bond. This
enables the smaller atom to exhibit stronger attractive forces on the
shared electrons within the bond.
Electronegativity Trends Within a
Group
 There is a decrease in electronegativity values of atoms moving
down a group.
Why?
 The atomic radii/sizes get larger as you move down a group.
 This results in the positive nucleus being farther away from the
electron pair involved in the bond. Therefore, there is less attractive
force to the positive nucleus.
Electronegativity Trends