14-1

CHEMICAL PERIODICITY

• PERIODIC TABLE• MENDELEEV• PREDICTION OF PROPERTIES

• Chapters 8.4; 9.5; 14.1-14.10

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Goals & Objectives

• See the following learning objectives on pages 322-323, 356, 563.

• Understand the Concepts:• 8.9-12; 9.13-15; 14.-8, 10, 14, 20, 23.

• Master these Skills: 8.6.

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CHEMICAL PERIODICITY

• THE PERIODIC TABLE

– 1869 Russian chemist Dmitri Mendeleev published a periodic table based on chemical properties of the elements known at the time

– 1869 German chemist Julius Lothar Meyer published a similar table based on physical properties

– Greater credit was given to Mendeleev because he was able to predict the properties of several undiscovered elements

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Mendeleev’s Periodic Table

Elements were grouped in order of atomic mass; if the next known element did not fit, Mendeleev left a space.

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Pictures of Mendeleev & Meyer

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Mendeleev’s Predictions

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The Periodic Law

• The properties of the elements are periodic functions of their atomic number

• This change is based on the work of British chemist Henry G. J. Moseley, who showed that the order of the periodic table is not based on atomic mass but rather on atomic number.

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Henry Moseley picture

• During the first World War, Moseley was killed while serving as a pilot in the British army at the age of 26. Thus a brilliant career was lost to science

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Terminology

• Group(family) of element - vertical column• Period(row) of elements - horizontal row• Group 1 metals - alkali metals• Group 2 metals - alkaline earth metals• Group 1 + Group 2 - representative metals• Group 16 - chalcogens• Group 17 - halogens• Group 18 - rare, inert or noble gases

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Classification of the Elements

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Classification of the Elements

Transitions metals Post trans.

Metals

Nonmetals

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This is a link to the Element Song.

The Elements" (1959) is a song by musical humorist Tom Lehrer, which recites the names of all the chemical elements known at the time of writing, up to number 102, nobelium. It can be found on his albums Tom Lehrer in concert, More Songs by Tom Lehrer and An Evening Wasted with Tom Lehrer. The song is sung to the tune of the Major General's Song from The Pirates of Penzance by Gilbert and Sullivan.

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This is a link to the New Element Song. The New Periodic Table Song

(2013) is a song by ASCAP Science, which recites the names of all the chemical elements (in order) including the new elements: Flerovium with the symbol "Fl" and atomic number 114 and Livermorium with the symbol “Lv” and atomic number 116.

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Chapter 14

Periodic Patterns in the Main-Group Elements

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Hydrogen

• Hydrogen has a very simple structure:– the nucleus has a single positive charge, and has 1 electron.

• Hydrogen is the most abundant element in the universe.

• Hydrogen exists as a diatomic gas, H2.– H2 is colorless and odorless with very low melting and boiling

points.

• H is abundant in combination with oxygen as H2O.

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Hydrogen and Group 1

• Like the Group (1) elements, H has a half-filled valence level.

• H is similar to the other Group 1 elements in terms of– ionization energy,

– electron affinity,

– electronegativity, and

– bond energies.

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Hydrogen and the Halogens

• Like the halogens or Group (17), hydrogen– exists as a diatomic molecule and

– needs only 1 electron to fill its valence shell.

• Unlike the halogens– H has a much lower electronegativity than any halogen,

– H lacks the three valence e- pairs that halogens have, and

– halide ions (X-) are common and stable, but the hydride ion (H-) is rare and reactive.

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Table 14.1 Trends in Atomic, Physical, and Chemical Properties of the Period 2 Elements.

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Table 14.1 Trends in Atomic, Physical, and Chemical Properties of the Period 2 Elements.

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Table 14.1 Trends in Atomic, Physical, and Chemical Properties of the Period 2 Elements.

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Table 14.1 Trends in Atomic, Physical, and Chemical Properties of the Period 2 Elements.

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Table 14.1 Trends in Atomic, Physical, and Chemical Properties of the Period 2 Elements.

Trends in atomic radius, ionization energy, and electronegativity across Period 2.

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Group (1): The Alkali Metals Family Portrait

KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONSCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Properties of the Alkali Metals

• Alkali metals are the largest elements in their respective periods and their valence electron configuration is ns1.– The valence e- is relatively far from the nucleus, resulting in

weak metallic bonding.

• Alkali metals are unusually soft for metals. They can be cut easily with a knife.

• Alkali metals have lower melting and boiling points than any other group of metals.

• Alkali metals have lower densities than most metals.

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Periodic Properties

• Electron configurations• Li [He]2s1

• Na [Ne]3s1

• K [Ar]4s1

• Rb [Kr]5s1

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Lithium floating in oil floating on water. Alkali metals have low densities.

Potassium reacting with water. Alkali metals are very reactive.

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Demonstration of the Reactivity of Alkali Metals

• The following YouTube video shows the reactivity of the alkali metals.

• Brainiac Alkali Metals

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Group (2): The Alkaline Earth Metals

• The oxides of Group (2) elements form basic solutions and melt at extremely high temperatures.

• Group (2) elements have higher ionization energies than Group (1) elements– due to their higher effective nuclear charge and smaller size.

• Group (2) elements are strong reducing agents.

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Group (2): The Alkaline Earth Metals Family PortraitKEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Group (13): The Boron Family Family Portrait

KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONSCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Influence of Transition Elements on Group (13)

The larger 13 elements have smaller atomic radii and larger ionization energies than electronegativities than expected.

These properties influence the physical and chemical behavior of these elements.

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Group (14): The Carbon Family Family PortraitKEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Allotropes

Allotropes are different crystalline or molecular forms of the same element.One allotrope of a particular element is usually more stable than another at a particular temperature and pressure.

Carbon has several allotropes, including graphite, diamond, and fullerenes.

Tin exhibits two allotropes; white β-tin and gray α-tin.

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Figure 14.9 Phase diagram of carbon.

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Figure 14.10 Crystalline buckminsterfullerene and a buckyball (A) and nanotubes (B).

Nanotubes

Crystals of buckminsterfullerene (C60)

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Carbon in Organic Chemistry

Catenation is the process whereby carbon bonds to itself to form stable chains, branches, and rings.

Since C is small, the C-C bond is short enough to allow effective side-to-side overlap of p orbitals. C readily forms double and triple bonds.

The large number and wide variety of organic compounds is due to the ability of C to bond to itself, and to form multiple bonds.

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Figure 14.12 Three of the several million known organic compounds of carbon.

Lysine, one of 20 amino acids that occur in proteins

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Inorganic Compounds of Carbon

Carbon forms two common gaseous oxides, CO and CO2, which are molecular. Other Group (14) elements form network-covalent or ionic oxides.

Carbon halides have major uses as solvents and in structural plastics.

Carbon bond with oxygen to form carbonates. Metal carbonates such as CaCO3 are abundant in minerals.

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Figure 14.13 Freon-12 (CCl2F2), a chlorofluorocarbon.

Chlorofluorocarbons (CFC’s or Freons) are chemically and thermally stable, nontoxic, and nonflammable. They are excellent cleaners, refrigerants, and propellants, but they decompose extremely slowly near the Earth’s surface. They readily enter the stratosphere, where UV radiation causes them to release free Cl atoms that damage the ozone layer.

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Group (15) Elements

• Nitrogen is a diatomic gas (N2) with a very low boiling point, due to its very weak intermolecular forces.

• Phosphorus exists most commonly as tetrahedral P4 molecules. It has stronger dispersion forces than N2.

• Arsenic exists as extended sheets of As atoms covalently bonded together. The covalent network structure gives it a high melting point.

• Antimony also has a covalent network structure.

• Bismuth has metallic bonding. Its melting point is lower than that of As or Sb.

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Group (15): The Nitrogen Family Family PortraitKEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Figure 14.16 Two allotropes of phosphorous.

White phosphorous (P4) Strained bonds in P4

Red phosphorous

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Periodicvideos.com

• This is the periodic video for Phosphorus.• Phosphorus Video

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Patterns of Behavior in Group (15)

• N gains 3 electrons to form the anion N3-, but only in compounds with active metals.

• The higher elements in the group are metallic and lose electrons to form cations.

• Oxides change from acidic to amphoteric to basic as you move down the group.

• All Group 5A(15) elements form gaseous hydrides with the formula EH3.

– All except NH3 are extremely reactive and toxic.

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Oxides of Nitrogen

• Nitrogen forms six stable oxides. DHf for all six oxides is positive because of the great strength of the NΞN bond.

• NO is produced by the oxidation of ammonia:– 4NH3(g) + 5O2(g) → 4NO(g) +6H2O(g)

– This is the first step in the production of nitric acid.

• NO is converted to 2 other oxides by heating:

– This type of redox reaction is called disproportionation.

• NO2 is a component of photochemical smog.

3NO(g) N2O(g) + NO2(g)Δ

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Table 14.3 Structures and Properties of the Nitrogen Oxides

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Group (16) Elements

• Oxygen, like nitrogen, occurs as a low-boiling diatomic gas, O2.

• Sulfur, like phosphorus, occurs as a polyatomic molecular solid.

• Selenium, like arsenic, commonly occurs as a gray metalloid.

• Tellurium, like antimony, displays network covalent bonding.

• Polonium, like bismuth, has a metallic crystal structure.

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Group (16): The Oxygen Family Family Portrait

KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONSCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Allotropes in the Oxygen Family

Oxygen has two allotropes:

- O2, which is essential to life, and

- O3 or ozone, which is poisonous.

Sulfur has more than 10 different forms, due to the ability of S to catenate. S–S bond lengths and bond angles may vary greatly.

Selenium has several allotropes, some consisting of crown-shaped Se8 molecules.

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Figure 14.20 The cyclo-S8 molecule.

top view side view

At room temperature, the sulfur molecule is a crown-shaped ring of eight atoms. The most stable S allotrope is orthorhombic α-S8, which consists of cyclo-S8.

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Hydrides of the Oxygen Family

• Oxygen forms two hydrides:– water (H2O) and hydrogen peroxide (H2O2).

– H2O2 contains oxygen in a -1 oxidation state.

• The hydrides of the other 16 elements are foul-smelling, poisonous gases.– H2S forms naturally in swamps from the breakdown of organic

matter and is as toxic as HCN.

• H2O and H2O2 can form H bonds, and therefore have higher melting and boiling points than other H2E compounds.

• Hydride bond angles decrease and bond lengths increase down the group.

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Halides of the Oxygen Family

Except for O, the Group 16 elements form a wide range of halides.Their structure and reactivity patterns depend on the sizes of the central atom and the surrounding halogens.

As the central atom becomes larger, the halides become more stable.

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Highlights of Sulfur Chemistry

• Sulfur forms two important oxides:– SO2 has S in its +4 oxidation state. It is a colorless, choking gas

that forms when S, H2S or a metal sulfide burns in air.

– SO3 has S in the +6 oxidation state.

• Sulfur forms two important oxoacids.– Sulfurous acid (H2SO3) is a weak acid with two acidic protons.

– Sulfuric acid (H2SO4) is a strong acid, and is an important industrial chemical. It is an excellent dehydrating agent.

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Figure 14.21 The dehydration of sugar by sulfuric acid.

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Group (17): The Halogens Family PortraitKEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Reactivity of the Halogens

A halogen atom needs only one electron to fill its valence shell. Halogens are therefore very reactive elements.

The halogens display a wide range of electronegativities, but all are electronegative enough to behave as nonmetals.

A halogen will either- gain one electron to form a halide anion or- share an electron pair with a nonmetal atom.

The reactivity of the halogens decreases down the group, reflecting the decrease in electronegativity.

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Group (18): The Noble Gases Family PortraitKEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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Noble Gases

• Noble gases have a full valence shell.

• The noble gases are the smallest elements in their respective periods, with the highest ionization energies.

• Atomic size increases down the group and IE decreases.

• Noble gases have very low melting and boiling points.

• Only Kr, Xe, and Rn are known to form compounds.– Xe is the most reactive noble gas and exhibits all even oxidation

states from +2 to +8.

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Figure 14.26 Crystals of xenon tetrafluoride (Xe(F4).

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Chapter 8

Electron Configuration and Chemical Periodicity

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Periodic Properties

• Electronic configurations

spd

f

f

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Electron Configurations

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Electron Configuration

• Give the symbol for the element having the outermost electron configuration 3s2

• Give the symbol for the element having the outermost electron configuration 4p4

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Trends in Atomic Size

Atomic size increases as the principal quantum number n increases.As n increases, the probability that the outer electrons will be further from the nucleus increases.

Atomic size decreases as the effective nuclear charge Zeff increases.As Zeff increases, the outer electrons are pulled closer to the nucleus.

For main group elements:

atomic size increases down a group in the periodic table

and decreases across a period.

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Figure 8.13

Atomic radii of the main-group and transition elements.

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Figure 8.14 Periodicity of atomic radius.

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Sample Problem 8.3 Ranking Elements by Atomic Size

PROBLEM: Using only the periodic table (not Figure 8.15), rank each set of main-group elements in order of decreasing atomic size:

(a) Ca, Mg, Sr (b) K, Ga, Ca

(c) Br, Rb, Kr (d) Sr, Ca, Rb

PLAN: Locate each element on the periodic table. Main-group elements increase in size down a group and decrease in size across the period.

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Sample Problem 8.3

SOLUTION:

(a) Sr > Ca > Mg

Ca, Mg, and Sr are in Group 2A. Size increases down the group.

(b) K > Ca > Ga

K, Ga, and Ca are all in Period 4. Size decreases across the period.

(c) Rb > Br > Kr

Rb is the largest because it has one more energy level than the other elements. Kr is smaller than Br because Kr is further to the right in the same period.

(d) Rb > Sr > Ca

Ca is the smallest because it has one fewer energy level. Sr is smaller than Rb because it is smaller to the right in the same period.

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Atomic Radii

size decreases

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Atomic Radii

• Give the symbol for the larger of the two atoms As or N

• Give the symbol for the atom having the smaller radius of Na or P

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Ionization Energy (IE)

• The energy in kilojoules required for the complete removal of one mole of electrons from one mole of gaseous atoms or ions

• Pulling an electron away from the nucleus requires energy to overcome the attraction of the nucleus.

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Trends in Ionization Energy

Ionization energy (IE) is the energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions.

Atoms with a low IE tend to form cations.

Atoms with a high IE tend to form anions (except the noble gases).

Ionization energy tends to decrease down a group and

increase across a period.

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Figure 8.15 Periodicity of first ionization energy (IE1).

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Figure 8.16 First ionization energies of the main-group elements.

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Ionization Energy

IE increases

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Ionization Energy

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Ionization Energy

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Sample Problem 8.4 Ranking Elements by First Ionization Energy

SOLUTION:

PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn

(c) K, Ca, Rb (d) I, Xe, Cs

PLAN: Find each element on the periodic table. IE1 generally decreases down a group and increases across a period.

(a) He > Ar > Kr

Kr, He, and Ar are in Group 8A. IE1 decreases down the group.

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Sample Problem 8.4

SOLUTION:

(b) Te > Sb > Sn

Sb, Te, and Sn are in Period 5. IE1 increases across a period.

(c) Ca > K > Rb

K has a higher IE1 than Rb because K is higher up in Group 1A. Ca has a higher IE1 than K because Ca is further to the right in Period 4.

(d) Xe > I > Cs

Xe has a higher IE1 than I because Xe is further to the right in the same period. Cs has a lower IE1 than I because it is further to the left in a higher period.

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Ionization Energy

• Give the symbol for the element in the same group as Sn having the largest IE.

• Give the symbol for the element in the same period as Si having the smallest IE.

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Electron Affinity (EA)

• Electron Affinity is the energy change in kilojoules accompanying the addition of one mole of electrons to one mole of gaseous atoms or ions.

• In most cases, energy is released when the first electron is added because it is attracted to the atoms nuclear charge.

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Trends in Electron Affinity

Electron Affinity (EA) is the energy change that occurs when 1 mol of electrons is added to 1 mol of gaseous atoms or ions.

Atoms with a low EA tend to form cations.

Atoms with a high EA tend to form anions.

The trends in electron affinity are not as regular as those for atomic size or IE.

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Figure 8.18 Electron affinities of the main-group elements (in kJ/mol).

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Behavior Patterns for IE and EA

Reactive nonmetals have high IEs and highly negative EAs. These elements attract electrons strongly and tend to form negative ions in ionic compounds.

Reactive metals have low IEs and slightly negative EAs. These elements lose electrons easily and tend to form positive ions in ionic compounds.

Noble gases have very high IEs and slightly positive EAs.These elements tend to neither lose nor gain electtons.

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Electron Affinity

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Electron Affinity

EA increases

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Electron Affinity

• Give the symbol for the representative metal in the same period as As having the larger electron affinity.

• Give the symbol for the metalloid in the same family as N having the smaller EA.

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Electronegativity (EN)

• Electronegativity is the relative ability of a bonded atom to attract the shared electron pair

• It was developed over 50 years ago by American Chemist Linus Pauling, who is the only person to win the Nobel Prize in two different categories – Chemistry in 1954 & Peace in 1963. He died in 1994 at age 93.

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Pictures of Linus Pauling

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Trends in Electronegativity

In general electronegativity decreases down a group as atomic size increases.

In general electronegativity increases across a period as atomic size decreases.

Nonmetals are more electronegative than metals.

The most electronegative element is fluorine.

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Figure 9.22 Electronegativity and atomic size.

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Electronegativity

EN increases

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Electronegativity

• Give the symbol for the nonmetal in the same period as Al having the greatest EN.

• Give the symbol for the element having the greatest electronegativity.

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Ionic Size vs. Atomic Size

Cations are smaller than their parent atoms while anions are larger.

Ionic radius increases down a group as n increases.

Cation size decreases as charge increases.

An isoelectronic series is a series of ions that have the same electron configuration. Within the series, ion size decreases with increasing nuclear charge.

3- > 2- > 1- > 1+ > 2+ > 3+

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Ionic Radii• Cations are always smaller than their

corresponding neutral atoms

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Ionic Radii• Anions are always larger than their

corresponding neutral atoms

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Ionic Radii

• Give the symbol for the element with the largest ionic radius between Cl and Cl-1.

• Give the symbol for the element with the smallest ionic radius between Na and Na+1.

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Binary Compounds

A. Salt of metal cation with only one charge and nonmetal anion

MgCl2

metal + nonmetal--metal in Group 2 metal + stem of nonmetal + ide

magnesium chloride

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Binary Compounds

Charges on metals Group 1 +1 Group 2 +2 Aluminum +3 Zinc (Zn) +2 Silver(Ag) +1 Cadmium(Cd) +2 See also Table 2.3 on page 63.

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Binary Compounds

lithium sulfide barium chloride aluminum bromide AgCl CdO Al2O3

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Binary Compounds

B. Salt of metal cation with more than one possible charge and nonmetal anion

All metals not covered in A

Name must include charge on metal

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Binary Compounds

IUPAC System FeCl3, FeCl2

metal + stem of nonmetal+ide charge on metal in Roman numerals

iron(III) chloride, iron(II) chloride

See also Table 2.4 on page 65.

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Binary Compounds

NiO CrCl3

SnCl4

Copper (II) bromide Iron (III) oxide Cobalt (III) nitride

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Binary Compounds

Older system “ic “ suffix designates the higher of two

possible charges “ous” suffix designates the lower of two

possible charges

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Binary Compounds

Memorize these four ions for the older system

Fe2+ - ferrous Fe3+ - ferric Cu+1 - cuprous Cu2+ - cupric Sn2+ - stannous Sn4+ - stannic Hg2

2+ - mercurous Hg2+ - mercuric See also Table 2.4 on page 65.

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Binary Compounds

ferrous chloride ferric sulfide cuprous iodide cupric bromide mercuric chloride mercurous chloride stannous fluoride