# Chapter 7 – Periodic Properties of the Elements

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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.

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 = 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 Where Can We Find Nonbinding 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

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 O22- 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

Reactions between hydrogen and nonmetals tend to be VERY exothermic
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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