Transcript Chapter 7 - GEOCITIES.ws

Periodic Properties of
the Elements
Chapter 7
The Periodic Table
 Developed independently by German Julius Lothar
Meyer and Russian Dmitri Mendeleev (1870”s).
 Found that similar chemical and physical
properties recur periodically when the elements
are arranged in order of increasing atomic number.
 At this time they didn’t know about atomic
numbers, so used masses which generally increases
as atomic # increases.
 Had blanks in the table and used periodicity to
guess about the characteristics of the missing
elements.
Moseley
 After Rutherford proposed the nuclear
model of the atom, Henry Moseley
developed the concept of atomic numbers.
 Bonding Atomic Radius- based on
the distance separating atoms when they
are chemically bonded to one another
 This radius is shorter than the non bonded
radius due to the nuclear attraction
between the two atoms.
 The bonding atomic radius decreases as
you go across the period, and increases as
you go down a group. (Rb> F)
Bond length
 Predict which will be greater, the P-Br
bond length in PBr3 or the As-Cl bond
length in AsCl3
 P-Br
Radial Electron Density
 The probability of finding an electron with
respect to the nucleus
 The 1s subshell in Ar is much closer to the
nucleus than the 1s subshell of He. This is
because of the Zeff
Ionic Radius
 Cations are formed when metallic atoms
lose valence electrons.
 These ions have smaller radii than their
parent atoms
 Anions are formed when nonmetallic atoms
gain electrons
 These ions are larger than their parent
atoms due to the extra repulsions of
another electron
Periodic Trends
 How easily an electron will be removed from
an atom is an important indicator of the
chemical nature of that atom.
 Ionization energy is the energy required to
remove an electron from the ground state a
gaseous atom
 The greater the ionization energy, the more
difficult it is to remove an electron
Ionization Energy
 Highest energy electron removed first
(outermost).
 First ionization energy (I1) is that required
to remove the first electron.
 Second ionization energy (I2) - the second
electron
 etc. etc.
Trends in ionization energy
 for Mg
• I1 = 735 kJ/mole
• I2 = 1445 kJ/mole
• I3 = 7730 kJ/mole
 The effective nuclear charge increases as you remove
electrons.
 It takes much more energy to remove a core electron
than a valence electron because there is less
shielding.
 This trend is because the positive nuclear charge that
provides attractive forces remains the same, while
the number of electrons which provide repulsive
forces decreases.
Explain this trend
 For Al
•
•
•
•
I1 = 580 kJ/mole
I2 = 1815 kJ/mole
I3 = 2740 kJ/mole
I4 = 11,600 kJ/mole
Ionization Trends
 Generally from left to right, I1 increases
because there is a greater nuclear charge
with the same shielding. (Generally, the
alkali metals show the lowest ionization
energies in a row, and the noble gases the
highest.
 As you go down a group I1 decreases
because electrons are farther away.
It is not that simple
 Zeff changes as you go across a period, so
will I1
 Half filled and filled orbitals are harder to
remove electrons from.
 here’s what it looks like.
Atomic
number
First Ionization
energy
Atomic
number
First Ionization
energy
Atomic
number
First Ionization
energy
Electron Affinities
 The energy change that occurs when an
electron is added to a gaseous atom to
form a negative ion. (A measure of the
affinity or attraction for the added
electron.
 For most atoms the electron affinity is
negative because energy is released when
an electron is added.
Different entities
 Remember!!!
 Ionization Energy is the desire to lose an
electron (+)
 Electron Affinity is the desire to gain an
electron
Electron Affinity Trends
 Generally becomes increasingly negative as
you go toward the halogens. For the noble
gases, the electron affinity is positive,
meaning the ion will not form because that
would mean that the gas would have to go
to a higher energy sub shell which is
energetically unfavorable.
 Any time the value is zero, the ion will not
form. The bigger the negative, the more
likely that the ion will form.
 Chlorine has the highest electron affinity.
Parts of the Periodic Table
Metals, Non metals, and Metalloids
Metals
 Roughly 3/4 of the elements are metals.
 Properties of metals include luster,
malleability, ductility, god conductors of
heat & electricity, form cations in an
aqueous solution.
 The more an element shows properties of
metals, the greater it’s metallic character
 (Increases down a row, decreases across a
period)
More Properties of Metals
 All metals are solid except for mercury
 Metals tend to have low ionization energies
which is why they are oxidized (lose
electrons) to form a cation when they
undergo a chemical reaction.
 When the outermost electrons are lost,
the ion achieves a noble gas configuration
 Many of the transition metals have the
ability to form more than one ion
Metal Oxides
 Metal-nonmetal compounds are said to be
ionic
 Oxides are especially important (oxygen is
everywhere!)
 Most metal oxides are basic, which means
they dissolve in water to form bases
 Metal oxides can also react with acids to
form salt + water
Nonmetals
 Poor conductors of heat & electricity
 Vary in appearance
 Have lower melting points than metals
 Several exist as diatomic molecules
 Tend to gain electrons in a chemical
reaction to fill their outer p subshell
completely giving a noble gas configuration
More properties of Nonmetals
 Nonmetal bonded to a nonmetal makes a
molecular substance
 These molecules tend to be gases, liquids,
or low-melting solids.
 Non-metal oxides are generally acidic
which means they combine with water to
form an acid. (This is why carbonated
water is acidic)
 These acidic nonmetal oxides will combine
with a base to produce salt & water
Metalloids
 Have properties of both metal and
nonmetal!
 Silicon- looks like a metal, brittle like a
nonmetal, semiconductor used in computer
chips
Group Trends
Active Metals
The Alkali Metals
 Group 1
 Soft metallic solids
 Doesn’t include hydrogen- it behaves as a
non-metal
 Down the group-decrease in IE
 Down the group-increase in radius
 Decrease in density
 Decrease in melting point
Alkali Metals
 For each row, the alkali has the lowest
ionization energy
 All very reactive and lose 1 electron to
form +1 cation
 Exist in nature only as compounds
 Electrolysis used
 React vigorously with water to produce
hydrogen gas and metal hydroxides
 Exothermic enough to ignite Hydrogen
 Also extremely reactive with oxygen
 Stored in kerosene
 Do not produce a colored solution because
no electron to excite
Alkaline Earth Metals
 All solids
 Harder , more dense, melt at higher temps than
alkali
 Slightly higher ionization energies, thus slightly
less reactive (compared to Alkali).
 Increasing reactivity as you go down the group
that accounts for why Berylium does not react
with water but Calcium and everything below it do.
 Tend to lose 2 electrons and form +2 cation
Alkaline Earth Metals
 Highly reactive so usually found in nature
as part of a compound
Group Trends
Nonmetals
Hydrogen
 Nonmetal that occurs as a colorless
diatomic gas
 Since there is no shielding, has an
extremely high ionization energy
 Usually combines with other non-metals to
form molecular compounds
 Reacts with active metals to form metal
hydrides (H is -)
Oxygen’s Group
 Changing trends as you go down the group
 Oxygen usually found in two molecular
forms oxygen and ozone (allotropesdifferent forms of the same element in
the same state)
 Oxygen makes up 21% of air
 Ozone is toxic and smelly
 Oxygen usually present as the oxide ion
Oxygen’s Group
 Sulfur (exists as eight membered rings of
sulfur atoms
 Most sulfur is found as metal sulfides
 Can be burned in oxygen to produce sulfur
dioxide (pollutant)
The Halogen Family
 (Astatine omitted because extremely rare,
radioactive and unknown)
 All typical non-metals
 Melting and boiling point increase as you go
down the group
 All diatomic
 Tend to gain electrons and form -1 anion
 Have highly negative electron affinities
 Fluorine and Chlorine most reactive
The Noble Gases
 All non-metal gases
 All monoatomic
 Rn too highly radioactive to study
 Completely filled s and p subshells
 Large ionization energies which decrease
as you go down the group
 Inert gases because thought to be unable
to form compounds