AP Chemistry Page 1 of 13Mr. Markic

Chapter 8 - Periodic Properties of the Elements

Development of the Periodic TableMendeleev• Grouped elements by their properties• Predicted the properties of other elements that had yet to be discovered

Moseley • Atomic number increases in the same order as atomic mass.

Representative Elements (main group elements)• Elements in group 1A through 7A• All of which have incompletely filled s or p subshells of the higher principal quantum number (except

helium)

Transition Metals (d block)• Elements in groups 1B and 3B through 8B, which have incompletely filled d subshells, or readily produce

cations with incompletely filled d subshells • Group 2B elements are not transition nor representative elements, they have no special name (Zn, Cd, and

Hg)

Lanthanide and Actinide• f – block transition elements• Incompletely filled f subshells

Alkali Metals• Group 1A• ns1

Alkaline Earth Metals• Group 2A• ns2

Halogens • Group 7A• ns2np5

Sample ExerciseAn atom of a certain element has 15 electrons. Without consulting a periodic table, answer the following questions:

(a) What is the ground-state electron configuration of the element

(b) How should the element be classified?

(c) Is the element diamagnetic or paramagnetic?

AP Chemistry Page 2 of 13Mr. MarkicAn atom of a certain element has 20 electrons.

(a) Write the ground-state electron configuration of the element

(b) Classify the element

(c) Determine whether the element is diamagnetic or paramagnetic

Representing Free Elements in Chemical EquationsMetals and Metalloids• The same as the symbol for the elementNonmetals• BrINClHOF – diatomic• Phosphorous P4 (stable form)• Sulfur – S (commonly used) than S8 (stable)

Electron Configurations of Cations and Anions Of Representative Elements

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Na+: [Ne]Al3+: [Ne]

F-: 1s22s22p6 or [Ne]O2-: 1s22s22p6 or [Ne]

N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

Atoms lose electrons so that cation has a noble-gas outer electron configuration.

Atoms gain electrons so that anion has a noble-gas outer electron configuration.

AP Chemistry Page 3 of 13Mr. MarkicElectron Configurations of Cations of Transition MetalsWhen a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.

Fe: [Ar]4s23d6 Mn: [Ar]4s23d5

Fe2+: [Ar]4s03d6 or [Ar]3d6 Mn2+: [Ar]4s03d5 or [Ar]3d5

Fe3+: [Ar]4s03d5 or [Ar]3d5

Review of ConceptsIdentify the elements that fit the following descriptions:

(a) An alkaline earth metal ion that is isoelectronic with Kr

(b) An ion with a -3 charge that is isoelectronic with K+

(c) An ion with a +2 charge that is isoelectronic with Co3+

Periodic Variation in Physical Properties1. Effective Nuclear Charge2. Atomic Radius3. Ionic Radius

1. Effective Nuclear Charge (Zeff)• e- are attracted to the nucleus of the atom (electrostatic force) • The closer the e- are to the nucleus, the more strongly it is attracted• The more protons in a nucleus, the more attracted an e- becomes• Completed shells are very stable• Valence e- are added or subtracted in order to create complete shells

Shielding Effect• e- are repelled by other e- in an atom• If there are e- between the nucleus and the valence e-, the valence e- will be less attracted to the nucleus• Filled inner shells shield outer electrons more effectively than electrons in the same subshell shield each

other

Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.

Zeff = Z – s 0 < σ < Z (σ = shielding constant)

Zeff » Z – number of inner or core electrons

AP Chemistry Page 4 of 13Mr. Markic

Z Core Zeff Radius (pm)

Na 11 10 1 186

Mg 12 10 2 160

Al 13 10 3 143

Si 14 10 4 132

2. Atomic Radius• One-half the distance between the two nuclei in two adjacent metal atoms• The larger the effective nuclear charge, the stronger the hold of the nucleus on

these electrons, the smaller the radius• Effective nuclear charge increases steadily, while the principal quantum number

remains constant• Across a period (left to right) radius decreases• Down a group radius increases with increasing atomic number

(a) In metals such as beryllium, the atomic radius is defined as one-half the distance between the nuclei of two adjacent atoms

(b) For elements that exist as diatomic molecules, such as iodine, the radius of the atom is defined as one-half the distance between the nuclei

Periodic Trends in Atomic Radii

i. Atomic radius increases moving down a group • Protons and e- are added• Valence e- get farther away from the nucleus due to the shielding effect

ii. Atomic radius decreases moving from left to right across a period • Protons and e- are added• Protons increase the force of attraction to the valence e- • e- are added to the same shell at the same distance• There is less shielding effect

Increasing Zeff

I n c r e a

AP Chemistry Page 5 of 13Mr. MarkicSample Exercise• Referring to a periodic table, arrange the following atoms in order of increasing atomic radius: P, Si, N

• Arrange the following atoms in order of decreasing radius C, Li, Be

Review of ConceptsCompare the size of each pair of atoms listed here:

(a) Be, Ba (b) Al, S (c) C12, C13

3. Ionic Radius• The size of an ion depends on its nuclear charge, number of e-, and the valence e-

• All ions increase in size as we go down a group in the periodic table

i. Cations are smaller than atoms • e- - e- repulsions are reduced• Valence e- move closer to the nucleus• Less shielding effect• Radius decreases with increasing nuclear

charge• More electrons lost, the smaller the ion

ii. Anions are larger than atoms • e- - e- repulsions increase• Valence e- move farther apart (increase radius)• Increases shielding effect• More electrons gained, the bigger the ion

Comparison of Atomic Radii with Ionic Radii(a) Alkali metals and alkali metal cations

(b) Halogens and halide ions

Sample ExerciseFor each of the following pairs, indicate which one of the species is larger:

(a) N3- or F- (b) Mg2+ or Ca2+ (c) Fe2+ or Fe3+

Select the smaller ion in each of the following pairs:

(a) K+ or Li+ (b) Au+ or Au3+ (c) P3- or N3-

AP Chemistry Page 6 of 13Mr. MarkicReview of ConceptsIdentify the spheres shown here with each of the following: S2-, Mg2+, F-, Na+

Ionization Energy• The minimum energy (kJ/mol) required to remove an e- from a gaseous atom in its ground state • It takes energy to remove an e- • An atom in the gaseous phase is virtually uninfluenced by its neighbors (no intermolecular forces) • The first ionization energy (I1) is the energy needed to remove the 1st e- • The atom becomes + charged• The greater the ionization energy, the more difficult it is to remove an e-

• There is a large increase in ionization energy when e- are removed from its noble gas core

Periodic Trends in 1st Ionization energies1. Ionization energy increases with increasing atomic

number • Alkali metals show low ionization energy• Noble gases show highest ionization energy

2. Ionization energy decreases moving down a group • The pull of the nucleus on the e- is reduced due to the shielding effect• Valence e- are easily pulled away

3. The subsequent ionization energy is greater than the previous ionization energy • e- - e- repulsion decreases when an e- is removed• Remaining valence e- move closer to the nucleus• Attractive force between e- and nucleus increases (increasing the ionization energy)• Ex: ionization energy is higher for Na+3 than Na+2

I1 + X (g) X+(g) + e- I1 first ionization

energy

I2+ + X (g) X2+

(g) + e- I2 second ionization energy

I3+2 + X (g) X3+

(g) + e- I3 third ionization energy

I1 < I2 < I3

Increasing First Ionization Energy

AP Chemistry Page 7 of 13Mr. MarkicPractice ExerciseWhich of the following elements will have: • Which atom should have a smaller first

ionization energy: oxygen or sulfur?

• Which atom should have a higher second ionization energy: lithium or beryllium?

• Which of the following atoms should have a larger first ionization energy: N or P?

• Which of the following atoms should have a smaller second ionization energy: Na or Mg?

Review of ConceptsLabel the plots shown here for the first, second, and third ionization energies for Mg, Al, and K.

Photoelectron Spectroscopy (PES) Provides data for ionization energy trends and applications Each peak is relative to the others. This indicates the relative number of electrons. If the peak is twice as

big, there are twice as many electrons. Gap is due to increased energy of the orbital (decreased amount of energy to remove the electron) Depending on the size of the table, 1s may be intentionally cut out of view because it’s too far away and

makes the graph too long Remember IE is about REMOVING electrons, which means they are removed from the OUTSIDE to the

INSIDE, and NOT in reverse order of energy! For example, 4s is removed BEFORE 3d.

AP Chemistry Page 8 of 13Mr. Markic1. Which element could be represented by the

complete PES spectrum below?

a. Lib. Bc. Nd. Ne

2. Which of the following best explains the relative positioning and intensity of the 2s peaks in the following spectra?

a. Be has a greater nuclear charge than Li and more electrons in the 2s orbital

b. Be electrons experience greater electron-electron repulsions than Li electrons

c. Li has a greater pull from the nucleus on the 2s electrons, so they are harder to remove

d. Li has greater electron shielding by the 1s orbital, so the 2s electrons are easier to remove

3. Given the photoelectron spectra above for phosphorus, P, and sulfur, S, which of the following best explains why the 2p peak for S is further to the left than the 2p peak for P, but the 3p peak for S is further to the right than the 3p peak for P?a. S has a greater effective nuclear charge than

P, and the 3p sublevel in S has greater electron repulsions than in P.

b. S has a greater effective nuclear charge than P, and the 3p sublevel is more heavily shielded in S than in P.

c. S has a greater number of electrons than P, so the third energy level is further from the nucleus in S than in P.

d. S has a greater number of electrons than P, so the Coulombic attraction between the electron cloud and the nucleus is greater in S than in P.

4. Looking at the spectra for Na and K above, which of the following would best explain the difference in binding energy for the 3s electrons?a. K has a greater nuclear charge than Nab. K has more electron-electron repulsions than

Nac. Na has one valence electrond. Na has less electron shielding than K

AP Chemistry Page 9 of 13Mr. Markic

5. Looking at the spectra for Na and K above, which of the following would best explain the difference in signal intensity for the 3s electrons?

a. K has a greater nuclear charge than Nab. K has more electron-electron repulsions than

Nac. Na has one valence electrond. Na has less electron shielding than K

6. Given the photoelectron spectrum above, which of the following best explains the relative positioning of the peaks on the horizontal axis?a. O has more valence electrons than Ti or C,

so more energy is required to remove themb. O has more electron-electron repulsions in

the 2p sublevel than Ti and Cc. Ti atoms are present in a greater quantity

than O than C in the mixture.

d. Ti has a greater nuclear charge, but the 2p sublevel experiences greater shielding than the 1s sublevel.

7. Given the photoelectron spectrum of Scandium above, which of the following best explains why Scandium commonly makes a 3+ ion as opposed to a 2+ ion?

a. Removing 3 electrons releases more energy than removing 2 electrons.

b. Scandium is in Group 3, and atoms only lose the number of electrons that will result in a noble gas electron configuration

c. The amount of energy required to remove an electron from the 3d sublevel is close to that for the 4s sublevel, but significantly more energy is needed to remove electrons from the 3p sublevel.

d. Removing 2 electrons alleviates the spin-pairing repulsions in the 4s sublevel, so it is not as energetically favorable as emptying the 4s sublevel completely.

8. On the photoelectron spectrum for magnesium given above, draw the spectrum for aluminum.

Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.

X (g) + e- → X-(g)

F (g) + e- → F-(g) ΔH = -328 kJ/mol EA = +328 kJ/mol

O (g) + e- → O-(g) ΔH = -141 kJ/mol EA = +141 kJ/mol

AP Chemistry Page 10 of 13Mr. MarkicWhy are the electron affinities of the alkaline earth metals, shown in Table 8.3, either negative or small positive values?

Is it likely that Ar will form the anion Ar-?

Review of ConceptsWhy is it possible to measure the successive ionization energies of an atom until all of the electrons are removed, but it becomes increasingly difficult and often impossible to measure the electron affinity of an atom beyond the first stage?

Metals Luster Malleable and ductile Good conductors of heat and electricity Have low ionization energies (forms cations

easily) Most metal oxides are basic Compounds of metals and nonmetals tend to

be ionic Going down a group – melting and boiling

points decrease Francium – highest metallic character

Nonmetals• No luster• Brittle and soft• Poor conductors of heat and electricity• Most nonmetal oxides form acidic solutions• Nonmetals gain e- when they react with

metals• High electron affinity• Going down a group – melting and boiling

points increase

Group 1A Elements (ns1, n ³ 2)

M M+1 + 1e-

2M(s) + 2H2O(l) 2MOH(aq) + H2(g)

4M(s) + O2(g) 2M2O(s)

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AP Chemistry Page 11 of 13Mr. Markic

Group 2A Elements (ns2, n ³ 2)

M M+2 + 2e-

Be(s) + 2H2O(l) No Reaction

Mg(s) + 2H2O(g) Mg(OH)2(aq) + H2(g)

M(s) + 2H2O(l) M(OH)2(aq) + H2(g) M = Ca, Sr, or Ba

Group 3A Elements (ns2np1, n ³ 2)

4Al(s) + 3O2(g) 2Al2O3(s)

2Al(s) + 6H+(aq) 2Al3+

(aq) + 3H2(g)

Group 4A Elements (ns2np2, n ³ 2)

Sn(s) + 2H+(aq) Sn2+

(aq) + H2 (g)

Pb(s) + 2H+(aq) Pb2+

(aq) + H2 (g)

Group 5A Elements (ns2np3, n ³ 2)

N2O5(s) + H2O(l) 2HNO3(aq)

P4O10(s) + 6H2O(l) 4H3PO4(aq)

Group 6A Elements (ns2np4, n ³ 2)

SO3(g) + H2O(l) H2SO4(aq)

Group 7A Elements (ns2np5, n ³ 2)

X + 1e- X-1

X2(g) + H2(g) 2HX(g)

Group 8A Elements (ns2np6, n ³ 2)

Completely filled ns and np subshells. Highest ionization energy of all elements.

Increasing Reactivity

Increasing Reactivity

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AP Chemistry Page 12 of 13Mr. Markic

No tendency to accept extra electrons.

Properties of Oxides Across a Period

• Oxygen combines with almost all elements• Oxides can be classifies as acidic, basic, or amphoteric• Amphoteric – display both acidic and basic properties• Oxides containing metallic elements are basic• Oxides containing nonmetallic elements are acidic• High atomic numbers tend to be more basic with oxides than lighter ones

Sample ExerciseClassify the following oxides as acidic, basic, or amphoteric:

(a) Rb2O (b) BeO (c) As2O5

Classify the following oxides as acidic, basic, or amphoteric:

(a) ZnO (b) P4O10 (c) CaO

Big Idea 1: The chemical elements are fundamental building materials of matter, and all matter can be understood in terms of arrangements of atoms. These atoms retain their identity in chemical reactionsDuration: 1 week (Late November/Early December)Textbook Chapter: 8

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AP Chemistry Page 13 of 13Mr. Markic

Enduring Understanding Essential Knowledge1.B: The atoms of each element have unique structures arising from interactions between electrons and nuclei.

1.B.1: The atom is composed of negatively charged electrons, which can leave the atom, and a positively charged nucleus that is made of protons and neutrons. The attraction of the electrons to the nucleus is the basis of the structure of the atom. Coulomb’s Law is qualitatively useful for understanding the structure of the atom.1.B.2: The electronic structure of the atom can be described using an electron configuration that reflects the concept of electrons in quantized energy levels or shells; the energetics of the electrons in the atom can be understood by considerationof Coulomb’s Law.

1.C: Elements display periodicity in their properties when the elements are organized according to increasing atomic number. This periodicity can be explained by the regular variations that occur in the electronic structures of atoms. Periodicity is a useful principle for understanding properties and predicting trends in properties. Its modern-day uses range from examining the composition of materials to generating ideas for designing new materials.

1.C.1: Many properties of atoms exhibit periodic trends that are reflective of the periodicity of electronic structure.1.C.2: The currently accepted best model of the atom is based on the quantum mechanical model.

1.D: Atoms are so small that they are difficult to study directly; atomic models are constructed to explain experimental data on collections of atoms.

1.D.1: As is the case with all scientific models, any model of the atom is subject to refinement and change in response to new experimental results. In that sense, an atomic model is not regarded as an exact description of the atom, but rather a theoretical construct that fits a set of experimental data.

1.E: Atoms are conserved in physical and chemical processes.

1.E.1: Physical and chemical processes can be depicted symbolically; when this is done, the illustration must conserve all atoms of all types.1.E.2: Conservation of atoms makes it possible to compute the masses of substances involved in physical and chemical processes. Chemical processes result in the formation of new substances, and the amount of these depends on the number and the types and masses of elements in the reactants, as wellas the efficiency of the transformation.

Learning Objectives1.5 The student is able to explain the distribution of electrons in an atom or ion based upon data.1.6 The student is able to analyze data relating to electron energies for patterns and relationships.1.7 The student is able to describe the electronic structure of the atom, using PES data, ionization energy data, and/or Coulomb’s Law to construct explanations of how the energies of electrons within shells in atoms vary.1.9 The student is able to predict and/or justify trends in atomic properties based on location on the periodic table and/or the shell model.1.10 Students can justify with evidence the arrangement of the periodic table and can apply periodic properties to chemical reactivity.1.11 The student can analyze data, based on periodicity and the properties of binary compounds, to identify patterns and generate hypotheses related to the molecular design of compounds for which data are not supplied.