Transcript Trends in the Periodic Table

Trends in the Periodic Table
trend: direction or pattern
Zumdahl2 p. 324-335
s, p, d and f blocks of the
periodic table
“Long form” periodic table
groups
periods
• Electronic structures are related to the
position of the elements on the periodic
table
– s-block: s orbitals are filled
– p-block: p orbitals are filled, etc.
Zumdahl2 p. 322
Atomic number
• Increases from left to right (across periods)
• Increases from top to bottom (down
groups)
• Period number: number of energy levels
containing electrons
• Group: same valence shell electron
configuration
Group names
• 1: alkali metals (1 valence electron)
• 17: halogens (7 valence electrons)
• 18: noble gases (all electron shells filled,
little chemical reactivity)
• periodicity: repeating of similar properties
because of similar valence electron
configuration
Atomic and Ionic Radii
• Atomic radius: half the distance between
the nuclei of two touching atoms
• Increases going down a group
– Additional filled energy levels of electrons
• Decreases going across a period
– More electrons = increased attraction to
positive nucleus
Atomic Radius Increases
Atomic Radius Increases
• Radius of an atom always decreases
when it loses an electron (becomes a
cation +)
– A whole energy level may be lost, or
– There is less electron-electron repulsion
(pushing away) between the electrons in
different energy levels
• The radius of an atom always increases
when it gains an electron (becomes an
anion -)
Ionization energy
• The amount of energy that is required to
remove an electron from a gaseous atom
• Decreases going down a group
– Valence electrons are further from the nucleus,
less “pull” from the protons
Effective Nuclear Charge
-
-
-
-
-
-
+15
-
-
-
+15
+5
-
+13
-
-
• Charge exerted on each electron by the
positively charged nucleus
Ionization Energy
• Increases going across a period
– Electrons on the same energy level are more
strongly “pulled” by the nucleus (which is
increasing in positive charge)
+11
+12
Ionization Energy Increases
Ionization Energy Increases
Electronegativity
• How strongly an atom attracts other
electrons in a chemical bond (electron
affinity)
• Decreases going down a group
– Valence electrons are further from the nucleus,
less “pull” from the protons
• Increases going across a period
– Electrons on the same energy level are more
strongly “pulled” by the nucleus
Electronegativity Increases
Electronegativity Increases
Other Physical Properties
•
•
•
•
Melting Point
Boiling Point
Density
Types of bonds formed
• Homework: handout
• Quiz tomorrow:
– Orbital notation
– Electron configuration notation
– Noble gas shorthand
Chemical Properties
• Elements in the same group have similar
chemical properties
Increasing atomic and
ionic radii
Increasing ionization
energy
Increasing
electronegativity
Increasing reactivity
Alkali metals
• Soft, malleable (can be shaped) metals
• Low melting points
– Can only contribute one electron to a bond –
easily broken
• Low density
– Largest atomic radius in the period
• Very chemically reactive
– One valence electron, easily lost, + ion
• Tarnish quickly
• Combine with O, Cl, Br to form ionic compounds
sodium
potassium
cesium (l)
rubidium
• All react with water to form a solution of
metal hydroxide and hydrogen
• 2M(s) + 2H2O(l)  2M+(aq) + 2OH-(aq) + H2(g)
sodium
Reaction is
alkaline
(base)
• M = alkali metal
potassium
• What trend in reactivity did you see as we
moved down the group?
Increasing ionization
energy
Increasing
electronegativity
Increasing reactivity
Increasing atomic and
ionic radii
Halogens
• Very reactive non-metals
– Need only one electron to fill valence shell
• All exist as diatomic molecules
– Cl2, Br2, I2 (all colored)
• Slightly soluble in water – non-polar
bonds
Halogens – colored diatomic
molecules
fluorine
pale yellow gas
chlorine
yellow-green
gas
bromine
red-brown liquid
iodine
black-purple solid
purple gas
• X2(aq) + H2O(l)  H+(aq) + X-(aq) + HOX(aq)
– X = halogen
– HOX = acid
• Ex: Chlorine: HOCl (HClO: chloric acid),
used as a bleach, toxic to microbes, treats
water
• all quite electronegative (high electron
affinity)
• easily gain electrons to form anions
– halide ions
• Ability to gain electrons decreases going
down a group
• reactivity decreases going down a group
• Halogens combine with metals to produce
ionically bonded salts containing a halide
ion.
– white and soluble in water  colorless solutions
– insoluble: lead and silver compounds
– lead(II) iodide: bright yellow precipitate
• Test for a halide ion by adding nitric acid, then a
solution of silver nitrate
– a precipitate indicates Cl-, Br-, or Icompound
color
AgF
no precipitate
AgCl
white, then purple/black in
sunlight
off-white
pale yellow
AgBr
AgI
silver chloride
used to print black
and white photos silver bromide
used to print black
and white photos
silver iodide
used to “seed the
clouds” to make it
rain
• Oxidant: In a reaction, a higher halogen
will replace a lower halogen.
• NaCl(aq) + Br- 
• NaBr(aq) + Cl-  NaCl(aq) + Br-(aq)
Atomic Radius Increases
Ionization Energy Increases
Electronegativity Increases
Metallic Character Increases
Your questions…
Answered!
• Q: Why don’t the electrons crash into the
nucleus?
• A: Electrons have lots of their own energy.
E=hf due to their position around the
nucleus. Electrons are constantly moving,
very fast.
• This kinetic energy overcomes the positive
attraction of the nucleus.
• A: How does temperature affect ionization
energy?
• Q: Temperature has no affect on ionization
energy. Heat is only powerful enough to change
kinetic energy of a particle or molecule.
• Microwaves and radio waves can affect nuclear
spin. Gamma rays and X rays can effect the
nucleus and the inner electrons.
• Electricity does have an affect on ionization
energy.
• Q: How does temperature affect the
movement of the subatomic particles,
specifically electrons?
• A: Temperature does not have enough
energy to affect subatomic particle
movement, only molecule movement.
• Q: Where does ionization energy come
from?
• A: Electricity
• Q: If there is just one electron in a px orbital, can it be
located everywhere, or just on one side of the orbital?
•A: It can be located
anywhere. An orbital is an
area of probability inside of
which the electron will be
found. Electrons are
constantly moving very fast,
and can be anywhere within
their orbitals.
• Q: Why does ionization energy increase
going across a period? Why, oh Why?!?
• A: The effective nuclear charge DOES
change on each valence electron!
ENC=+2
ENC=+1
+11
+12