Transcript Going across a period

Chapter 6:
Periodic Trends
Development of the
Periodic Table
Dmitri Mendeleev designed the
first periodic table in 1869 by
grouping elements with similar
chemical & physical properties in
rows according to Atomic mass
Henry Moseley rearranged the
table in 1913 according to atomic
number
The Modern Periodic Table
Organized by groups (columns) and periods (rows)
Classifying Elements by
Electron Configurations
Representative Elements- Outermost s and p are filling
 Alkali metals – Outermost s has 1 e
 Alkaline earth metals – Outermost s is full
 Halogens – Outermost s is full, outermost p has 5 e
 Noble Gases- Outermost s and p are totally full
Transition Metals -Outermost s is full and d is filling or full
Inner Transition Metals- Outermost s is full and f is filling
or full
Periodic Trends
The location of an
element on the
Periodic Table gives
a lot of information
about key properties
of that element
compared to the
other elements
As you move across the
periodic table, you keep
periodically running
across the same
properties...
Hence the
name
periodic
table.
Trend #1: Atomic Size
An atoms size is
determined by its
atomic radius
r
Atomic radius is
defined as half the
distance between
the nuclei of two
like atoms
Atomic Size Trend
Going down a group, atoms get bigger because the
number of energy levels increases.
Going across a period, atoms get smaller. No new
energy levels are added, but more protons and
electrons are. The increase in protons and
electrons result in a greater pull towards the
nucleus. (much like magnets)
Trend #2: Atomic Size of IONS
Cations (positive ions) are always
smaller than their parent atom
because they lose electrons and an
energy level. A strong attraction
forms between the electrons left
over and the nucleus.
Anions (negative ions) are always
larger than their parent atom.
Gaining electrons causes less
attraction to the nucleus – the new
electrons aren’t as attracted and
are free to move around!
Comparison of atomic vs. ionic size
Answer this . . .
Which is larger: S or S-2. Why?

S-2 is larger – Gaining electrons causes less of an attraction to the
nucleus – the new electrons are free to move around.
Which is smaller: Fe or Fe+4. Why?

Fe+4 is smaller – Losing electrons causes more of an attraction to the
nucleus – the remaining electrons are pulled in tighter.
Which is smaller: Na+ or Al3+? Why

They both lose an energy level BUT Al3+ loses MORE e-, nucleus
pulls in the remaining electrons due to the more drastic change..
Which is smaller: Be2+ or Na+. Why?

They both lose an energy level BUT Be2+ loses MORE e- & goes
down to a smaller energy level.
An atom walks into a ‘restaurant’
and proclaims, "Hey! Somebody
just stole one of my electrons!"
The bartender says, "Are you
sure?“ The atom replies, "Yes I'm positive!"
A neutron walks into a
‘restaurant’ and says,
"Hey bartender give
me a drink."
The bartender gives
him one and says,
“No charge for you"
Trend #3: Electronegativity
ELECTRONEGATIVITY - the ability for an element
to attract an electron
Going down a group, electronegativity decreases
because the added energy levels ‘shield’ the
power of the nucleus to attract electrons
Going across a period, electronegativity increases
because nuclear charge increases - the closer an
atom is to having a full outer shell, the greater
the desire to completely fill that shell by gaining
electrons.
Which element would be the most
electronegative? Why?
Fluorine - It is located at a low energy level
with a high number of electrons and protons.
There is a strong attraction between the
electrons and the nucleus. It wants an electron
BADLY!
How does ‘Shielding’ work?
SHIELDING EFFECT: The process of the inner
electrons shielding (repelling) the outer electrons - it
causes the outer electrons to be less attracted to the
nucleus.
Shielding increases as you go down the periodic table
because more electrons in more energy levels are added
Shielding remains constant as you go across because all
the electrons in a period are in the same energy level.
Why aren’t noble gases assigned
electronegativity numbers?
They are happy the way they are – they don’t
need any more electrons 
Trend #4: Ionization Energy
IONIZATION ENERGY- The energy needed to REMOVE an
electron from the outside shell.
Going down a group, ionization energy decreases because
more energy levels are added. This ‘shielding’ causes
the electrons to be less attracted to the nucleus little energy is needed to remove them.
Going across a period, ionization energy increases because
the electrons are closer to the nucleus. More energy is
needed to remove them the closer they are to the
nucleus.
First vs. Second
First ionization energy is the energy required to
remove the first electron
Second ionization energy is the energy required
to remove the second electron (and so forth)
First ionization energy is always smaller
(compared to the second or third ionization
energy) because each successive electron
removed is closer to the nucleus and strongly
attracted to it.
Trend #5: Metallic Properties
First off, what’s meant by ‘metallic’?
Differences between metals and nonmetals
Metals
Solid at room temp
Malleable (Can be hammered
into sheets) and ductile (Can be
drawn into wires)
Nonmetals
Most are gases at room
temp, some are solid or
liquid.
Not malleable nor ductile –
mostly brittle or powdery
if solid
Conduct electricity and heat
well
Lose electrons to become
cations
Conduct electricity and
heat poorly
Gain electrons to become
anions
Metallic properties increase as you go down a
group, and decrease as you go across a period.
X
Fr is the most metallic!