Download chapter 7- periodic properties of the elements

CHAPTER 7- PERIODIC PROPERTIES OF
THE ELEMENTS
• Periodic table- One of the most significant tools
chemists have for organizing and remembering chemical
facts
• In Chapter 6- what did we say was primarily responsible
for the properties/reactivities of elements?
• Electron configuration- repeating pattern on periodic
table
• Responsible for periodic nature of the elements
• Elements in same column (group)= same number of
valence electrons
• What were valence electrons again?
• Elements in same group- what else is special about them?
• Why do elements in the same group have similar
properties?
• What you haven’t heard- elements in same group share
properties, but also have differences
• Ex: Oxygen vs. Sulfur- both in group 6A- bond/react similarly
• Oxygen: colorless gas at room temp
• Sulfur: yellow solid at room temp
• These differences can also be attributed to electron
configuration
• What shell (value of n) are the valence electrons of O in?
• What shell are the valence electrons of S in?
SEC. 7.1
• Mendeleev- created one of the first widely accepted
periodic table
• Inspired by his favorite card game, solitaire, Mendeleev
created a set of cards of the elements- one card for each
element currently discovered.
• Each element card had information on it: the element’s
name, atomic mass, properties, and compounds they
formed with other elements
• Mendeleev organized his periodic table in order of
increasing atomic mass- he spread all the cards out,
placed them in a big line according to increasing mass
• When ordered this way, he noticed a pattern arise among
the elements- similar properties actually repeated
themselves in a regular pattern (a periodic pattern)
• Periodic Table Build!
• You are given a set of element cards- all info on these
cards should be used in order to organize the
elements you are given into the first periodic table
(correctly)
• Discovery of chemical elements has been an ongoing process
since ancient times
• Though many elements are stable- tend to form compounds
• Not many elements found in elemental form naturally- why’s
that?
• Many elements went undiscovered due to the fact that they
naturally form compounds
• Early 1800s- advances in chem allow for isolation of elements
from their compounds
• These advances= many elemental discoveries
• With number of elements increasing, the demand for a
method/way of organizing them increased
• Many scientists tried many different classifying schemes
• In steps Mendeleev with his crazy table, rambling on about how
his way of organizing the elements was the best way
• How did he organize it?
• Why did you end up leaving holes in your periodic table?
• What did you end up doing with the holes that were left?
• Mendeleev ended up with the same holes in his periodic tablehe insisted that elements with similar properties go in the same
families- this forced him to leave holes in his table
• Using his awesome table, Mendeleev boldly went were no man
had gone before and decided he would predict the existence of
these undiscovered elements- he went as far as predicting their
atomic mass, melting point, color, density, specific heat, boiling
point, and the compounds it would form with oxygen and
chlorine
• He called the element hole under Aluminum “eka- Aluminum”
and the element hole under Silicon “eka- Silicon”
• When Gallium and Germanium were discovered, they very
closely matched the properties Mendeleev has predicted
• This proved how useful Mendeleev’s periodic table was- it can be
used to predict properties of undiscovered elements
Property
Mendeleev’s prediction
for eka-Silicon (1871)
Germanium properties
(1886)
Atomic weight
72
72.59
Density
5.5 g/mL
5.35 g/mL
Specific heat
0.305 (J/gK)
0.309 (J/gK)
Melting Point
High
947 °C
Color
Dark gray
Grayish white
Oxide Formula
XO2
GeO2
Chloride Formula
XCl4
GeCl4
Boiling Point of chloride
compound
A little under 100 °C
84 °C
• Even though Mendeleev’s table was the bees knees, it had
its problems
• Do you remember which elements broke the trend of
increasing atomic mass? Why did you place them out of
order?
• After Rutherford’s model of the atom, a physicist named
Henry Mosely decided he would shoot high energy
electrons at atoms
• These atoms emitted X-rays of specific and unique
frequencies- frequency emitted generally increased as
atomic mass increased
• Assigned each element a specific whole number values
based on the frequency of X-ray emitted- atomic number
• Stated atomic number= number of protons in nucleus
• Ordered periodic table in order of increasing atomic number- this is
how the modern periodic table is arranged
• Ordering in increasing atomic number- clarified the problems seen in
Mendeleev’s table
• This new arrangement also allowed for the identification of
holes/undiscovered elements- property predictions as well- led to
discovery of even more elements
SEC. 7.2- EFFECTIVE NUCLEAR CHARGE
• Electrons- negative charge
• Nuclei- contain protons- positive charge
• Electron attracted to nucleus- atom properties depend on
electron configurations and how strongly their outer electrons are
attracted to the nucleus
• Coulomb’s Law= strength of interaction between 2 electrical
charges depends on magnitudes of charge and distance between
them
• Attractive force between electrons and nucleus increases as nuclear
charge increases
• Decreases as electrons move further away
• Many electron atoms- all electrons attracted to nucleus
and repelled by one another at the same time
• Due to number of repulsions, it is difficult to measure all
the forces
• However, we can estimate the net attraction of each
electron to the nucleus- must consider how single
electron interacts with “average” environment created by
nucleus and other electrons
• This allows us to treat each electron as if it were moving in a net
electric field created by nucleus and electron density of other
electrons
• Effective Nuclear charge (Zeff)- the nuclear charge felt by an electrontakes into consideration the “shielding” effects the inner shell
electrons have; this is the “net” nuclear charge felt by individual
electron
• Shielding- caused by the repulsions between electrons
• Effective nuclear charge will always be smaller than the actual nuclear
charge
• Valence electron- attracted to nucleus in atom, but also repelled
by inner shell electrons- the repulsions from inner shell electron
partially “cancel out” some of the attraction felt from the
nucleus (shielding)
• Zeff = Z – S
• Zeff= effective nuclear charge, Z= nuclear charge (# protons in nucleus),
S= shielding constant
• S is portion of nuclear charge that is “cancelled out” by
repulsion of inner shell electrons
• S is normally close to the value of inner shell electrons
• Electrons in same valence shell are not good at shielding
one another
ZEFF FOR A SODIUM ATOM
• Periodic trend: Zeff for valence electrons increases as we move
from left to right across a period
• Why do you think this is? What changes as we move from left to
right?
• Going down a group, Zeff slightly increases, but not by much- not
nearly as much as it does moving across a period
• Would expect the Zeff to be about the same for elements in same
group- why?
SEC. 7.3- SIZES OF ATOMS AND IONS
• Size of atoms and ions is a somewhat strange concept, but is also
extremely important
• Atoms- not hard, spherical objects- if you remember from
chapter 6, electron clouds are not really “hard” boundaries- no
sharp cut off for the boundary at which electron distribution
becomes zero
• We define atomic size in several ways- mostly based on the
distance between atoms in various situations
• Argon atoms- when these gas particle collide, they easily
bounce off one another without reacting/interacting all
that much
• This is due to the fact that the electron clouds of these gas
atoms will not mesh together to any significant extent –
no attraction because these atoms are not charged, and
the clouds will actually somewhat repel one another
• When atoms collide in this way, the collision is sometimes
referred to as a “non-bonding” collision- since the atoms
are simply ricocheting off one another
• Nonbonding atomic radius- distance between two nuclei
during a “non-bonding” collision
• When atoms are bonded to one another- like Cl2- there is some
sort of attractive force between the bonded atoms- hence why
they are bonded
• Due to this attractive force between bonded atoms, the nuclei
of the bonded atoms will be closer to one another than they
would be if they were not bonded- the bond kind of squeezes
their electron clouds closer together
• Bonding atomic radius- the distance between the
nuclei of bonded atoms
• How does bonding atomic radius compare to
nonbonding atomic radius?
• Through observation,/experiments, scientists have gathered data
on the bonding atomic radius for every element- each element
assigned its own bonding radius value
• With these values, we are able to tell how long a bond will be
between any combination of bonded atoms in molecules
• Ex: For C, bonding atomic radius= 0.77 Å, and Cl= 0.99 Å
• In the compound CCl4, we can estimate the bond length of a C-Cl bond to
be 1.76 Å (In reality, the bond length is 1.77 Å, which is pretty close to
the estimated value)
PERIODIC TRENDS IN ATOMIC RADII
• Two trends for atomic radius:
• 1- Atomic radius tend to increase from top to bottom (going
down) a group. Why do you think this is? What happens to
the value of n as we go down a group?
• 2- Atomic radius tends to decrease when moving from left to
right across a period. Major influence for this is the Zeff. Why
does Zeff have an effect on atomic radius?
PERIODIC TRENDS IN IONIC RADII
• Ionic radii- determined by distance between nuclei of ions
• Size of ion depends on effective nuclear charge, number of
electrons it possesses, and the orbitals in which its valence
electrons are located
• When a cation forms (positive ion), does this mean that
electrons were gained or lost?
• When an anion forms (negative ion), does this mean that
electrons were gained or lost?
• If cations form when electrons are lost, how do you think the
ionic radius of a cation compares to that of the regular atomic
radius?
• If anions form when electrons are gained, how does the ionic
radius of an anion compare to that of the regular atomic
radius?
• Ionic radius of cation < atomic radius
• Ionic radius of anion > atomic radius
• Sooo… cations are smaller than their “parent” atoms,
and anions are larger than their “parent” atoms
• Also, for ions with same charge, size increases as we
move down a group
• Based on periodic trends for atomic radius and ionic
radius, order to following atoms and ions in order of
increasing size (smallest to largest): Mg+2, Ca+2, and Ca
• Isoelectronic series: groups of ions all containing same number of
electrons
• Ex: O2-, F-, Na+, Mg2+, Al3+ … all have how many electrons
• If members of isoelectronic series are listed in order of increasing
atomic number, effective nuclear charge will increase. Why?
• If nuclear charge increases, what will happen to the ionic radius?
• List the ions above in order of increasing ionic radius
SEC. 7.4- IONIZATION ENERGY
• How easily electrons can be removed from atoms/ions has
major impacts on chemical behavior
• Ionization energy- minimum energy required to remove
an electron from the ground state of the isolated gaseous
atom or ion
• First ionization energy (I1)- energy required to remove the
first electron from a neutral atom
• Second Ionization energy (I2)- energy needed to remove
the second electron
• This pattern continues for third, fourth, fifth ionization
energies (continues until atom is out of electrons)
• Which ionization energy do you think is the smallest?
Why?
• First ionization energy for the sodium atom is the energy
required for the process:
• I2 for sodium is energy associated with this process:
• Ionization energies- always positive, due to the fact that energy
must be applied to the atom to pull an electron away
• First ionization energy is always the smallest- the first electron is
easiest to pull away- partially due to the fact that the atom is
neutrally charged at this point
• Higher ionization energies= more difficult to remove electrons
• This is due to the fact that after the first ionization, the atom
becomes positively charged- creating a stronger attraction
between the electrons and the nucleus
• There is also a trend when it comes to inner shell vs. valence
shell electrons- which electrons do you think would be easiest to
remove?
• There is a sharp increase in ionization energy once all valence
electrons have been removed- greater increase than from the
first to second ionization energy (if more than 1 valence electron)
• Inner shell electrons- much more difficult to remove than
valence electrons- why?
• The fact that valence electrons have such low ionization energies
than the inner shell electrons supports the fact that valence
electrons are mostly responsible for how something bonds/
chemically reacts
• The valence electrons will be the only ones willing to
transfer/share- form ionic/covalent bonds
• Inner electrons too tightly bound to nucleus to be lost from atom or
even be shared with another atom
• Predict which atom will have the largest second ionization energy:
Calcium, Sodium, and Sulfur
PERIODIC TREND OF FIRST IONIZATION ENERGY
• First ionization energy, I1, typically increases going left to
right across the same period (some exceptions to this
rule)
• Ionization energy typically decreases going down a groupas atomic number increases
• The s and p block element show a larger range in values of I1 than
the d block (transition metals) do. Generally, the ionization
energy of the transition metals increases slowly as we move from
the left to the right across a given period.
• In general, smaller atoms have higher ionization energies- why do
you think this happens?
• The energy needed to remove an electron depends on:
• Effective nuclear charge
• Distance of the electron from the nucleus
• Increasing effective nuclear charge or decreasing the distance
between the electron and nucleus will increase the attraction
between the electron and the nucleus
• As attraction to nucleus increases, it becomes more difficult to
remove the electron- more attraction = more required energy=
higher ionization energy
• Moving left to right across a period, what happens to:
• Effective nuclear charge?
• Atomic size?
• So energy needed to remove electron will:
• When moving down a group, what happens to:
• Effective nuclear charge?
• Atomic size?
• So, energy needed to pull electron away will:
• Arrange the following from smallest to largest first ionization
energy: Ne, Na, P, Ar, K
SEC. 7.5- ELECTRON AFFINITIES
• Electron affinity- energy change that occurs when an
electron is added to a gaseous atom
• Measures attraction (affinity) of an atom to an added electron
• Measures the ease in which an atom will gain an electron
• Electron affinities are typically negative energy values
(this indicates that energy is released when electron
added)
• If the electron affinity value is positive, this means the atom
will not gain an electron
• The more negative the electron affinity is, the more the atom
will want to gain an electron
• No real periodic trend when it come to electron affinitieshowever, some groups of elements have characteristically
low/high affinities
• The Halogens- Have the most negative electron affinity values…
why?
• The Noble Gases- do you think their electron affinities will be
negative or positive? Why?
SEC. 7.6- METALS, NONMETALS, AND METALLOIDS
• Elements can be broadly grouped into the categories of metals,
nonmetals, and metalloids
• Metalloid staircase- marked on your periodic tables
• Left of staircase- metals (any exceptions?)
• Right of staircase- Nonmetals
• Touching staircase on two sides- metalloids (Boron, silicon,
germanium, arsenic, antimony, tellurium)- form a border
between metals and nonmetals
• What category do most elements on the periodic table belong
to?
• Metallic character- a measure of how much an element exhibits
the physical and chemical properties of metals
• Metallic character- Increases moving down a group; decreases
moving left to right across a period
METALS
• Shiny luster; various colors- often silvery
• Good conductors of heat and electricity (this is due to
their metallic bonding)
• Most are Malleable- can be pounded into thin sheets
• Most are Ductile- can be pulled into a thin wire
• High melting points- all are solid at room temp except
one… which one?
• Tend to have low ionization energies- what does this mean
they tend to do with their valence electrons?
• Tend to form what charges?
• When metals undergo chemical changes, they typically oxidize
(Oxidation is Loss) due to low ionization energies/ electron loss
• For s block metals- they will lose electrons in the outermost s
shell- forming positive charges- either 1+ or 2+ charges form
• For p block metals- They will either lose just the outer p
electrons, or both the outer s and p electrons
• Al3+ - what block did the electrons this metal lost belong to?
• Sn2+ - what block did the electrons this metal lost belong to?
• Sn4+- what block did the electrons this metal lost belong to?
• Charges formed by transition metals do not follow a distinctive
pattern- often able to form more than one positive ion
• When metals bond with nonmetals, they typically form an ionic
compound
• When metals left outside, typically combine with the nonmetal
oxygen (abundant in our air) to form an ionic compound:
•
Ni +
O2 →
• Which one of these reactants- Ni or O2, will be easier to pull electrons
away from? How do you know this?
• Ionic bonding- one atom will lose electrons- the other will gain- which
element is losing electrons in this reaction? Which one is gaining?
• Whenever metals react with oxygen, we call the products metal
oxides- metals oxides are fairly common ionic compounds due to
the abundance of oxygen in nature
• Metal oxides have a special characteristic- most of them are basic
when in solution
• When metal oxides are able to dissolve in water (if soluble) they
react to form metal hydroxides:
• Metal oxide + water → metal hydroxide
• If you look at the net ionic equation, the basicity of metal
oxides is due to the oxide ion, which reacts with water as
follows:
• Metal oxides also display their basicity when reacting with
acids:
• Metal oxide + acid → salt + water
• What kind of reaction does that look like?
NONMETALS
• Nonmetals- vary greatly in their physical appearance
• No luster, and are poor conductors of heat and electricity
• Melting points- typically lower than those of metals
(Notice that all gases on the periodic table are nonmetals)
• Not malleable/ductile- may be very hard, brittle, or soft
• Nonmetals- how do their ionization energies compare to those
of metals? What about their electron affinities?
• So when nonmetals bond with metals, what will they typically
do?
• Ex: Aluminum reacting with bromine
• When nonmetals bond ionically (with other metals) they will gain
enough electrons to fill their outermost occupied p subshell
• When nonmetals bond with other nonmetals- this forms a
covalent bond (electrons are shared)
• Nonmetal oxides- acidic!
• When dissolved in water, nonmetal oxides react to form an acid
• Nonmetal oxide + water → acid
• Ex:
• Nonmetal oxides will also react with bases to form salt
and water
• Nonmetal oxide + base → salt + water
• Ex:
METALLOIDS
• Metalloids- kind of like the strange child of metal/nonmetal
• Metalloids have properties that are both metallic and
nonmetallic- depends on which metalloid you are dealing
with
• Ex: Silicon- looks like a metal, but when you hit it with a
hammer, it shatters (not malleable like metal)
• Several metalloids- most notably silicon- are
semiconductors and are used in computer chips/circuits
• Pure Silicon= electrical insulator… but if you add specific
impurities, its electrical conductivity dramatically
increases- this makes it possible to control electrical
conductivity by simply controlling chemical composition
SEC. 7.7- GROUP TRENDS FOR THE ACTIVE METALS
• We have seen that elements in the same group posses
general similarities
• There are trends found without groups- using periodic
table and our knowledge of electron configuration to
examine the chemistry of the alkali metals, and the
alkaline earth metals
GROUP 1A- ALKALI METALS
• Alkali Metals= metals in group 1A
• Soft, metallic solids, silvery, metallic luster, and good conductors
of heat/electricity
• They have a fairly low density and melting point (compared to
other metals)
• Moving down the alkali metals: melting point decreases, density
increases, what about atomic radius? First ionization energy?
• For each row (period) on the periodic table, the alkali metals
have the lowest I1 value- what does this reflect about its outer s
electron?
• Due to the low ionization energy- alkali metals lose electrons
very easily= very very very reactive
• All will form what charge?
• Alkali metals only exist as compounds- partially due to high
reactivity
• Will combine directly with most nonmetals
• Alkali metals will combine with hydrogen to form hydrides and
sulfur to produce sulfides:
Na + H2 →
Na + S →
• These reactions happen with all alkali metals
• When reacting with H2, alkali metals will give up their electron to
H, which will then form a hydride ion:
• All alkali metals are violently reactive with water
• When they combine with water, they produce hydrogen gas
and a metal hydroxide
• The reaction between alkali metals and water is extremely
exothermic- which means it releases a lot of heat- most of
the time this is enough heat to ignite the H2 gas that is
also produced… this makes a boom
• How do you think the violence of this reaction changes as
you go down a group?
• Alkali metal ions= colorless in compounds/ solutions- But,
when placed in flame, they all emit a characteristic color of
light
• When placed in flame, the alkali metal ions are reduced to
gaseous metal atoms, and the valence electron absorbs
energy from the heat of the flame
• What happens when the single valence electron of an alkali
metal absorbs energy?
• When electrons become excited, they will want to
immediately return to ground state
• Ground state?
• What will the excited valence electron have to do in
order to return to ground state?
• When releasing energy, the excited electron releases a photon
with unique energy- due to this unique energy, the photon
emitted will have a special wavelength- this will correspond
with a specific color of light
• Every metal ion releases a unique color of light when placed in
a flame- why?
• Ex: Sodium will release yellow light after being excited- valence
electron in 3s jumps to 3p when excited- when returning back
to ground state (3p to 3s) it releases a photon with a
wavelength of 589 nm (yellow light)
GROUP 2A: THE ALKALINE EARTH METALS
• Alkaline earth metals are harder than alkali metals, have
higher melting points, and are more dense
• Going down the alkaline earth metals group, melting points
decrease, density increases, atomic radius? First ionization
energy?
• How do you think their reactivity compares to the alkali
metals?
• What about reactivity going down the alkaline metal group?
• The way the alkaline earth metals react with water shows the
trend of reactivity (increases going down the alkaline earth
metals)- none are as reactive as the alkali metals
• Beryllium does not react with water, but every other alkaline
earth metal does
• Ex: Magnesium and water (all other alkaline earths- besides
beryllium- react with water the same way)
• Which electrons will the alkaline earth metals lose?
• What ion will all alkaline earth metals form?
• Alkaline earth metal with chlorine and oxygen:
• Heavier alkaline earths (lower in group) are even more
reactive to nonmetals
• Alkaline earth metals are similar to alkali metals when in the
presence of a flame
• When heated to very high temperatures, their valence
electrons will become excited- once excited they will release
photons with a characteristic color (energy/wavelength) to
return to ground state
• Flame test analysis can help us identify the type od metal in
an unknown compound
SEC. 7.8- GROUP TRENDS OF SELECTED NONMETALS
HYDROGEN
• Very strange- doesn’t know exactly what to do with its valence electronwill be happy losing it, or gaining 1 more
• Hydrogen does not lose its valence electron as easily as the alkali metals
• When bonding with other nonmetals, hydrogen will share its valence
electron instead of losing it
• When reacting with alkali metals, will gain an electron to form H(hydride ion)
• In aqueous chemistry, hydrogen loses its electron to form H+, hydrogen
cation (acid!)
GROUP 6A- THE OXYGEN FAMILY
• Not many shared physical characteristics- mix of nonmetals,
one metalloid, and a metal
• Oxygen- gas at room temp, the rest of the group are solids
• Oxygen, sulfur, selenium- typical nonmetals; telluriummetalloid; Polonium- radioactive metal
• All have 2 valence electrons and will typically form what type
of charge:
• Oxygen- found in 2 naturally occurring molecular forms
• O2= dioxide (typically referred to as just “oxygen”), and O3
= ozone
• O2 and O3 are allotropes- different forms of the same
element in the same state (O2 and O3 are gases)
• O2 is necessary for life- we use it for respiration
• O3 is found in upper atmosphere and polluted air- damaging to
health, but also necessary for life on Earth- ozone absorbs a large
part of the UV radiation emitted from the Sun
• O2 has a great tendency to attract electrons (oxidizes other
elements)- when combining with metal, forms oxide ion O2• Sulfur also has a tendency to gain electrons from other elements to
form sulfides containing the S2- ion (also does this in combination
with metals)
GROUP 7A- THE HALOGENS
• Group 7A= the Halogens (group beginning with fluorine)
• All are typical nonmetals (astatine is rare and radioactive,
so not much data)
• All (except astatine) are diatomic
• Going down the halogen group- what do you think
happens to melting point?
• All have different colored vapors- F2= pale yellow, Cl2=
yellow-green, Br2= reddish-brown, I2= violet vapor
• Highly negative electron affinities- meaning what?
• The halogens are the most reactive nonmetals due to
their extremely negative electron affinities
• What ion will all halogens form?
• Fluorine- most reactive, removes electrons from almost
any substance it comes into contact with
• Halogens will react with metals to form metal halides:
• Halogens also react with hydrogen to form gaseous
hydrogen halide compounds:
GROUP 8A- THE NOBLE GASES
• Group 8A= The Noble Gases- far right of the periodic table
• All are nonmetals and gases at room temperature
• All are monatomic- single atom, not a molecule (not
diatomic)
• Completely filled s and p subshells- this makes them
happy (full valence shells)
• All have very large ionization energies- meaning what?
• All have positive electron affinities- meaning what?
• Stable electron configuration= very unreactive
• Do not readily form compounds
• Radon- radioactive