chapter 7 – periodic properties of the elements homework: 9, 10, 12, 13, 14, 15, 17, 19, 21, 25,...
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Chapter 7 – Periodic Properties of the Elements
Homework: 9, 10, 12, 13, 14, 15, 17, 19, 21, 25, 27, 28, 30, 33, 35, 37, 39, 41, 43, 47, 51, 54, 55, 57, 59, 63, 65, 67, 69,
71, 75, 78, 105
7.2 – Effective Nuclear Charge
Because electrons are – charged, they are attracted to nuclei (which are + charged)
Many properties of atoms depend not only on their electron configurations, but also on how strongly their outer electrons are attracted to the nucleus.
Coulomb’s law tells us the strength of this interaction depends on the signs and magnitude of their charges, as well as the distance between them.
Thus, the force of the attraction between electron and nucleus depends on the:
Net nuclear charge acting on the electron Average distance between the nucleus and the electron
This force of attraction increases as the nuclear charge increase and decreases as the electron moves farther from the nucleus.
In many-electron atoms, each electron is simultaneously attracted to the nucleus and repelled by other electrons. Generally, there are so many electron-
electron repulsions that we cannot analyze the situation exactly.
We CAN estimate the net attraction of each electron to the nucleus by looking at how the electron interacts with its average environment
Environment created by the nucleus and the other electrons in the atom.
Effective Nuclear Charge By looking at the average situation, we
treat each electron individually as if it were moving in a net electric field created by the nucleus and the electron density of other electrons. We view this net electric field as if it results
from a single positive charge located at the nucleus.
This charge is called the effective nuclear charge, Zeff
The Zeff for an electron is less than the actual nuclear charge, because it takes the electron-electron repulsions into account. So if Z = actual nuclear charge Zeff < Z
Finding Zeff
A valence electron is attracted to the nucleus, but repelled by other electrons in the atom.
In particular, the electron density that is due to the inner electrons are very effective at partially canceling the attraction of the valence electron to the nucleus. We say the inner electrons partially shield or
screen the outer electrons from the attraction in the nucleus.
We can write a relationship between the effective nuclear charge and the number of protons in the nucleus (Z) Zeff = Z – S
Where S is a number called the screening constant
S represents how much the inner electrons shield the valence electron from the nucleus. The value of S is usually close to the
number of core electrons in an atom. Valence electrons don’t shield each
other, only core electrons shield.
Expected Zeff Values What would we expect the Zeff value for
Sodium to be? [Ne]3s1
So Z would be 11 (11 protons) We would expect S to be 10 (10 inner, or core,
electrons) So the expected Zeff value would be +1.
Is this what the Zeff actually is? No Because the 3s orbital has a very small change of
being close to the nucleus (figure 6.18(c)), there is a change of the 3s electron will have a greater attraction than normal.
Implications? This effective nuclear charge explains
energy amounts for the different l values within the same energy level (same n) 2s sublevel has greater probability to be closer
to nucleus than the 2p. This leads to slightly greater attractive force
between nucleus and 2s Which leads to lower energy for 2s than for 2p.
This rule applies to all of the energy differences between sublevels.
Trends in Zeff for Valence Electrons The effective nuclear charge as we move
across a row (period) increases. This is because the core electrons remain the same,
but the total nuclear charge increases. In other words, Z increases, but S remains the same.
Going down a column (family), the effective nuclear charge changes much less than moving across.
The expected nuclear charge should stay the same But we actually experience a slight increase in
effective nuclear charge as we move down a family But this slight increase is MUCH less than the
increase in moving across a period.
7.3 – Sizes of Atoms and Ions
An important property of atoms or ions is their size But don’t think of an atom or ion as a
hard, spherical object Remember, there are no sharply defined
boundaries at which electrons are, then aren’t
But, we can still define atom size in several ways, based upon the distance between individual atoms in different situations.
Ways to Determine Radius Imagine a bunch of argon atoms at
room temperature When they collide, the ricochet apart like
pool balls. This happens because the electron clouds of
each atom can’t penetrate the other atom Electrons repel
The closest distance the nuclei get during such a collision determines the apparent radii of the argon atoms.
We call this radius the nonbinding atomic radius
Now imagine a chlorine molecule (Cl2) A chemical bond holds each chlorine
atom to the other This attractive interaction brings the
atoms closer than they would be in a nonbinding collision.
This distance is called the bonding atomic radius, and is shorter than the nonbinding atomic radius.
Space-Filling Models Space-filling atomic models,
like the one below are based on the non-binding radii (also called the van der Waals radii) and binding radii Non-binding radii used to
determine size of each sphere Binding atomic radius used to
determine how deep each sphere penetrates the other
Chemists use different methods to measure the distance separating nuclei in molecules. In the I2 molecule
The distance separating the nuclei is found to be 2.66Å (1 Å = 10-10 m)
We find the bonding atomic radius to be ½ the bond distance, so 1.33Å.
Generally, the bonding atomic radius is ½ the nonbinding radius
Where Can We Find Nonbinding Radii?
pg. 266
Figure 7.6
Periodic Trends in Atomic Radii Within each column (family) atomic
radius tends to increase as you go down the column. This is primarily because of the increase in
the principal quantum number Electrons have a greater probability of being
farther from the nucleus. Within each row (period), atomic radius
tends to decrease from left to right This is primarily because of the increase in
the Zeff, which draws valence electrons closer to the nucleus
Causing atomic radius to decrease
Periodic Trends in Ionic Radii
The radii of ions are based upon the distances between ions in ionic compounds.
Like the size of an atom, the size of an ion depends on its nuclear charge, on the # of electrons it has, and on the orbitals in which the valence electrons are found.
Cation
The formation of a cation empties the larger occupied orbitals in an atom, and also decrease the number of electron-electron repulsions.
Cations will be smaller than their parent atoms
Anions
The formation of an anion begins to fill the larger occupied orbitals in an atom, and also increases the number of electron-electron repulsions.
Anions will be larger than their parent atoms
Similar Charges
For ions with the same charge size increases as we go down in a column of the periodic table
Isoelectric Ions An isoelectric series is a group of ions
that have the same number of electrons O2-, F-, Na+, Mg2+, and Al3+ all have 10
electrons In any isoelectric series, we can list the
members in order of increasing atomic number (nuclear charge) Because the number of electrons remains
the same, the radius of the ions decrease with increasing atomic number, as the electrons are more strongly attracted to the nucleus
7.4 – Ionization Energy
The ease with which electrons are removed from an atom or ion has a major effect on chemical behavior.
The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground state of the gaseous atom or ion.
The first ionization energy, or I1, is the energy needed to remove the first electron from a neutral atom Na(g) Na+
(g) + e-
The second ionization energy, or I2, is the minimum energy needed to remove the second electron from an ion that has already lost one electron Na+(g) Na2+
(g) + e-
The greater the ionization energy, the more difficult it is to remove an electron.
Variations in Successive Ionization Energies
As successive electrons are removed from an atom, the ionization energy increases for that ion. I1 < I2 < I3 and so on
This is because it is harder to pull an electron away from a more positive and positive ion.
There is a sudden increase in ionization energy when an inner shell electron is removed. Silicon
[Ne]3s23p2 I1 = 786 kJ kJ/mol I4 = 4360 kJ/mol I5 = 16,100 kJ/mol
This is because the 2p electrons are closer to the nucleus than the four 3rd shell electrons, and thus the 2p electrons experience a greater Zeff.
Every element shows a large increase in ionization energy when electrons are removed from its noble-gas core. This lends support to the idea that only
the outermost electrons (beyond the noble-gas core) are involved in bonding.
The inner electrons are too tightly bound to the nucleus to be lost or even shared with another atom.
Periodic Trends in the First Ionization Energies
Within each row (period), I1 will generally increase with increasing atomic number. So alkali metals show the lowest, and noble
gases show the highest. Within each column (group, the
ionization energy generally decreases with increasing atomic energy. Helium has higher ionization energy than
xenon.
The representative elements show a larger range of values of I1 than transition-metals do. Generally, transition metals increase
slowly as we proceed from left to right in a period.
The f-block metals only show a very small increase in ionization energies.
In general Smaller atoms have higher ionization
energies. Because the same factors that effect atomic
size also influence ionization energies. The energies needed to remove an electron
from the outermost shell depends on both the effective nuclear charge and the distance from the nucleus.
As Zeff increases, ionization energies increase.
There are occasional irregularities within a given row. Explained by energy levels Higher energy level generally = lower
ionization energies.
Electron Configuration of Ions When electrons are removed from
an atom to form a cation, they always come from the orbitals of the highest energy level first. Iron example Fe Fe2+
[Ar]3d64s2 [Ar]3d6
If we removed a 3rd electron, it would now be removed from the 3rd energy level, now that the 4th is empty.
7.5 – Electron Affinities
The first ionization energy of an atom measures the energy change associated with the removal of an electron from the atom to form a positively charged ion. Cl(g) Cl+(g) + e- ΔE = 1251 kJ/mol The positive ΔE tells us energy must
be put into the atom to remove the electron.
Most atoms can gain electrons to form negatively charged ions.
This energy change that occurs when an electron is added to a gaseous atom is called the electron affinity Measures the attraction, or affinity, of the
atom for the added electron. For most atoms, energy is released
when an electron is added.
Example
Cl Cl-
Cl(g) + e- Cl-(g) ΔE = -349 kJ/mol
Ionization Energy ≠ Electron Affinity
Ionization energy measures the ease with which an atom LOSES an electron
Electron affinity measures the ease with which an atom GAINS an electron
The greater the attraction between a given atom and an added electron, the more negative the atom’s electron affinity. For some elements (like the noble
gases), the electron affinity has a positive value.
Means the anion is higher in energy than the separated atom and electron
Means we must add energy to it to add an electron
Trends in Electron Affinity The halogens will have the highest electron
affinity Gaining an electron gives them a noble-gas
configuration The electron affinities of the noble-gases, Be,
and Mg are positive Any more electrons and we’d have to start filling up
in a currently-empty higher-energy subshell Interesting notes
(N, P, As, Sb) The added electron gets put into an orbital that is
already occupied (p subshell), leading to increased electron-electron repulsion
As we move down a group electron affinity doesn’t change greatly. Although the electrons would be more
spaced out (less electron-electron repulsion) the distance to the nucleus is greater (lower electron-nucleus attraction).
7.6 – Metals, Nonmetals and Metalloids
The elements can be broadly grouped into the categories of metals, nonmetals and metalloids.
Most elements are metals, found in the left and middle portion of the periodic table.
Nonmetals are found at the top right corner
Metalloids are between the metals and nonmetals
Metals Most have a shiny luster Conduct heat and electricity Malleable All (except mercury) are solid at room
temperature Low ionization energies
Tend to form cations relatively easily Characterized as an “oxidizer” because
they tend to lose electrons during chemical reactions.
Metals and Others Compounds of metals with nonmetals
tend to be ionic substances Oftentimes produce oxides, a group metal
ions where a metal bonds with an oxygen NiO, Na2O
Most metal oxides are basic When they dissolve in water, they react to
form metal hydroxides The metal hydroxides are basic The basicity of metal oxides are due to the
oxide ion.
Nonmetals Nonmetals vary greatly in appearance Not lustrous Generally poor conductors of heat and
electricity Melting points tend to be lower than
those of metals Seven nonmetals exist as diatomic
molecules H2, N2, O2, F2, Cl2, Br2, I2
Nonmetals can be solids, liquids or gases at room temperature
Chemical Reactions Because of their electron affinities,
nonmetals tend to gain electrons when they react with metals.
Compounds made entirely of nonmetals are molecular substances.
Most nonmetal oxides are acidic When they dissolve in water they will
react to form acids
Metalloids B, Si, Ge, As, Sb, Te, At are the metalloids Have properties somewhere between
metals and nonmetals Silicon looks like a metal, but is brittle and is a
poor conductor of heat and electricity than most metals.
Compounds of metalloids have characteristics of the compounds of metals or nonmetals, depending on the compound.
7.7 – Group Trends for the Active Metals
We’re going to look at the chemistry of the alkali metals (1A)and the alkaline earth metals (2A)
Group 1A: The Alkali Metals Soft metallic solids
Silvery metallic luster and high thermal and electrical conduciveness
Named from the Arabic word meaning “ashes” Originally isolated from wood ashes by early
chemists Low densities and melting points
Melting points decrease as we move down the group
Densities increase as we move down the group
Since they have the lowest I1 values, they can have electrons very easily removed Very reactive So reactive, they exist in nature only
as compounds. Combine directly with most
nonmetals
React vigorously with water, producing H2 gas and a solution of an alkali metal hydroxide M stands for any alkali metal 2M(s) 2H2O(l) 2MOH(aq) + H2(g) This reaction is very exothermic
Often enough energy is released to ignite the H2
This reaction is more violet with lower alkali metals, since they more easily lose their electron, and are more reactive
Other reactions
When oxygen reacts with metals, metal oxides (containing O2- ion) are formed 4Li(s) + O2 2Li2O When dissolved in water, Li2O and
other soluble metal oxides react with water to form OH- ions from the reaction of O2- with H2O
Other alkali metals all react with oxygen to form metal peroxides, which contain the O2
2- ion 2Na(s) + O2(g) Na2O2(s)
K, Rb, Cs also form compounds that contain the O2
- ion, called superoxide ion. K(s) + O2(g) KO2(s)
Generally
Generally we only worry about the oxide formation Don’t worry about the peroxides or
superoxides
Flame Tests
Although alkali metal ions are colorless, each emits a color when put into a flame. The ions are reduced to gaseous
metal atoms, then the flame exists the valence electrons.
As the valence electrons fall back, they emit light.
Reaction Example
Write a balanced equation for the reaction of cesium metal with water.
Group 2A: The Alkaline Earth Metals Like alkali metals, all solid at room
temperature and have typical metallic properties
Tend to be harder, more dense and have higher melting points than the alkali metals
I1 is low, but higher than the alkali metals, so less reactive.
Reactions Be doesn’t react with water Mg won’t react with liquid water, but
will react with steam Mg(s) + H2O(g) MgO(s) + H2(g)
Ca, Sr, Ba and Ra will react with water (but slower than the alkali metals) Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)
In general, tends to lose their two outer s electrons to form 2+ cations.
Burning!
Like the alkali metals, the heavier alkaline earth metals give off colors when strongly heated
7.8 – Group Trends for Selected Nonmetals
Hydrogen Because it has a 1s1 electron, usually
positioned with the alkali metals But H really doesn’t belong in any
group Not an alkali metal, because it’s usually a
colorless diatomic gas Though under extreme pressure, hydrogen
can be metallic Like found at the core of Jupiter or Saturn
Because there is no nuclear shielding of its sole electron, the ionization energy for H is quite high
Actually closer to the ionization energies for nonmetals such as O or Cl
Less likely to lose an electron than the alkali metals Tends to share its electrons and form molecular compounds
Reactions between hydrogen and nonmetals tend to be VERY exothermic
Reacts with metals to form solid metal hydrides Contains the hydride ion, H-
This property of hydrogen gaining electrons makes it more like a halogen than an alkali metal
But hydrogen can lose its electron to form a cation H+(aq) ion is relatively common, but we’ll study it later
Group 6A: The Oxygen Group
Going down 6A, we change from nonmetallic to metallic elements Oxygen, sulfur and selenium are
nonmetals Tellurium is a metalloid Polonium is a metal
All except oxygen are solids at room temperature
Oxygen 2 common forms of oxygen
O2 and O3 O2 commonly referred to as just “oxygen” O3 referred to as “ozone.” O2 is more stable
These two forms are examples of allotropes
An allotrope are different forms of the same element in the same state
Sulfur
Most stable of the sulfur allotropes is S8
Even though solid sulfur contains S8, we usually write it as S(s) to simplify stoichiometric coefficients.
Like oxygen, tends to gain electrons to form sulfides, which form the S2- ion Most sulfur in nature is actually in the form
of metallic sulfides. Though less likely to form sulfides than
oxygen is to form oxides.
Group 7A: The Halogens Astatine is usually left out of this group,
because it’s unstable and many of its properties not yet well known.
Typical nonmetals Melting and boiling points increase with
increasing atomic number F and Cl are gasses, Br is a liquid, and I is a
solid All are diatomic
Chemical Properties
Highly negative electron affinities Tend to gain electrons from other
atoms to form halide ions, X-
X being used to represent any of the halogens
F and Cl more reactive than Br and I
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