Unit 1 Chp 7 Notes B and L.ppt

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Periodic Relationships

Among the Elements

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General Periodic Trends

to know:

• Atomic size (atomic radius)Atomic size (atomic radius)

• Ionic size (ionic radius)*Ionic size (ionic radius)*

• Ionization energyIonization energy

• Electron Affinity*Electron Affinity*

• ElectronegativityElectronegativity

• Metallic CharacterMetallic Character

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Half of the distance between nuclei in covalently bonded

diatomic molecule

"covalent atomic radii"

Determination of Atomic Radius:

1. Atomic Radius

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Periodic Trends in Atomic Radius

Radius inc down a group

Addition of principal quantum levels (larger “n”) and increase in shielding

Reason?

Radius dec across a period

Reason?

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Effective Nuclear Charge?

Z_eff = Z – S

Z = atomic number (number of protons)

S = screening constant (# of core electrons)

Core electrons are those not in the valence shell. They help to shield the nucleus from the other electrons in the cloud.

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Nuclear Charge for Na?

Z_eff = 11 – 10 = +1

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Table of

Table of

Atomic

Radii

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2. Ionic Radii

Cations

Positively charged ions

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Ionic Radii

Anions

Negatively charged ions

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Cation is always smaller than atom from which it is formed.

Anion is always larger than atom from which it is formed.

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Table of Table of Ion Sizes Ion Sizes

*Ionic Radii follows same trend as

Atomic radii – if only comparing ions

(cations and anions separately)

*If comparing atoms and ions,

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Examples:

1. Arrange Na, Be and Mg in order of increasing atomic radius.

Be < Mg < Na

2. Arrange K+, Cl-, Ca2+ and S2- in order of

increasing ionic radius.

Ca2+ < K+ < Cl- < S

2-3. Arrange Mg2+, Ca2+ and Ca in order of

increasing radius.

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Isoelectronic Series

Isoelectronic: ions all containing the same numbers of electrons.

 Increasing Nuclear Charge 

O2- F- Na+ Mg2+ Al3+ 1.26 Å 1.19 Å 1.16 Å 0.86 Å 0.68 Å

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AP Practice Question

The elements in which of the following have most nearly the same atomic radius?

A.Be, B, C, N

B.Ne, Ar, Kr, Xe C.Mg, Ca, Sr, Ba D.C, P, Se, I

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AP Practice Question

Which of the following best helps to account for the fact that the F- ion is smaller than the

O2- ion?

A.F- has a larger nuclear mass than O2- has.

B.F- has a larger nuclear charge than

O2- has.

C.F- has more electrons than O2- has.

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 Increases for successive electrons taken from

the same atom

3.

Ionization Energy

- the energy required to remove an electron from a

gaseous atom

Ionization of Mg

Mg + 738 kJ Mg+ + e-

Mg+ + 1451 kJ _ Mg2+ e-

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AP Practice Question

The ionization energies for element X are listed in the table above. On the basis of the data,

element X is most likely to be A. Na

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 Tends to increase across a period

Electrons in the same quantum level do not shield as effectively as electrons in inner levels, thus nuclear charge increases

Irregularities at half filled and filled

sublevels due to extra repulsion of

electrons paired in orbitals, making them easier to remove

 Tends to decrease down a group

Outer electrons are farther from the

Nucleus, less energy to remove, electron shielding, Coulomb’s Law

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IE is endothermic (heat E

needed to remove e-)

Can be written as an equation:

Example for atom of sodium: 1s22s22p63s1

Na(g)  Na+(g) + e- ΔH = 496 kJ mol-1 Na+(g)  Na2+(g) + e- ΔH = 4562 kJ mol-1

*Why is the 2nd IE (I

2)so much higher? Core electrons being removed

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Table of 1

Table of 1stst

Ionization Energies

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Can you spot the discontinuities?

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AP Practice Questions

Which of the following represents an electron configuration that corresponds to the valence electrons of an element for which there is an especially large jump between the second and third ionization energies? (Note: n represents a

principal quantum number equal to or greater than 2.)

A.ns2

B.ns2_np1

C.ns2np2

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AP Practice Question

For element X represented above, which of the following is the most likely explanation for the large difference

between the second and third ionization energies? A.The effective nuclear charge decreases with

successive ionizations.

B.The shielding of outer electrons increases with successive ionizations.

C.The electron removed during the third ionization is, on average, much closer to the nucleus than the first two electrons removed were.

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AP Practice Question

Which of the following correctly identifies which has the higher first-ionization energy, Cl or Ar, and supplies the best justification?

A.Cl, because of its higher electronegativity B.Cl, because of its higher electron affinity

C.Ar, because of its completely filled valence shell

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 Affinity tends to increase across a period  Affinity tends to decrease as you go down

in a group.

Electrons farther from the nucleus experience less nuclear attraction *Some irregularities due to repulsive forces in the relatively small p orbitals

4. Electron Affinity

- the - the energy changeenergy change associated with the

associated with the additionaddition of an of an electronelectron

*Opposite meaning than IE, but same trend.

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Electron Affinity is Exothermic

(E is released by attracting e-)

Cl(g) + e-  Cl-(g) ΔH = -349 kJ mol-1

*Ionization E and Electron affinity are opposites

•IE measures the ease to lose e- to become

cation (+)

•Affinity measures the ease to gain e- to

become anion (-)

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Trends in Electron Affinity

There are again,

however, two discontinuities in this trend.

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Trends in Electron Affinity

• The first occurs

between Groups IA and IIA.

– The added electron must go in a p orbital, not an s orbital.

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Trends in Electron Affinity

• The second

discontinuity occurs between Groups IVA and VA.

– Group VA has no empty orbitals.

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Which of the following best helps to explain why the electron affinity of Br has a greater magnitude than that of I?

A.Br has a lower electronegativity than I does. B. Br has a lower ionization energy than I does. C. An added electron would go into a new shell in Br but not in I.

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5. Electronegativity

Is a measure of the

ability of an atom in

a molecule to attract

electrons to itself.

Concept proposed by

Concept proposed by

Linus Pauling

1901-1994

Concept proposed by Concept proposed by

Linus Pauling Linus Pauling

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Trend in Electronegativity?

_{Electronegativities tend to}increase

across a period WHY???

stronger nuclear charge and more

attraction of the electrons to the nucleus

_{Electronegativities tend to}decrease

down a group or remain the same WHY???

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Periodic Table of Electronegativities Highest

Electronegativity

Decreases as elements get

FURTHER

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Which of the following best helps explain why the electronegativity of Cl is less than that of F?

A.The mass of the Cl atom is greater than the mass of the F atom. B.The Cl nucleus contains more protons than the F nucleus

contains.

C.When Cl and F form bonds with other atoms, the Cl bonding electrons are more shielded from the positive Cl nucleus than the F bonding electrons are shielded from the positive F nucleus. D.Because Cl is larger than F, the repulsions among electrons in the valence shell of Cl are less than the repulsions among

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Trends Review

Which period 3 element could be

represented by the ionization energy data collected below? Provide an explanation.

Ionization Energy (kJ/mol)

1st ₇₃₈

2nd ₁₄₅₁

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Rank the Following:

Rank the following ions in order of

increasing ionic radius and explain your answer.

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Isoelectronic Series

Isoelectronic: ions all containing the same numbers of electrons.

 Increasing Nuclear Charge 

O2- F- Na+ Mg2+ Al3+ 1.26 Å 1.19 Å 1.16 Å 0.86 Å 0.68 Å

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Properties of Metal, Nonmetals,

and Metalloids

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“Special” Groups

• Group 1: Alkali metals

• Group 2: Alkaline earth metals

• Group 3-12: Transitional elements

• _{Group 17: Halogens}

• Group 18: Noble gases

• Elements 58-71: Lanthanides

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Metals

• Located on the left-hand side of the stairs • Good conductors of electricity and heat • Easily lose electrons to form cations

• Malleable • Ductile

• Luster

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Nonmetals

• Located on the right-hand side of the stairs • Poor conductors of electricity and heat

• Tend to gain electrons to form anions • Brittle

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Metalloids

• Located in purple below (make up the stairs)

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Metals vs Nonmetals

• Metals tend to form cations.

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6. Metallic and 7. Nonmetallic

Character

Metallic character: Decrease across the PT and increases down PT

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Summary of Periodic Trends

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You Should also know

generalities

• Mass increases (for the most part) from L to R and Top to Bottom

• PT organized by increasing atomic number • HONCl BrIF elements (diatomic gases)

• Monatomic gases (Noble gases)

• Nuclear Charge, E- shielding, PEL, Coulomb’s Law

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AP Practice Questions

Which of the following correctly compares periodic properties of two elements and provides an accurate explanation for that difference?

A.The first ionization energy of Al is greater than that of B because Al has a larger nuclear charge than B does.

B.The first ionization energy of F is greater than that of O because O has a higher electronegativity than F has.

C.The atomic radius of Ca is larger than that of Mg because the valence electrons in Mg experience more shielding than the

valence electrons in Ca do.

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Trends with Valence Electrons

• Valence electrons: electrons in the highest outer most principle

energy level of an atom • Across a period the

valence electrons increase by 1

• Down a group the

valence electrons stay the same. Why?

Bohr Model for

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Remember this?

*Elements in the same

group end with the same # of electrons in each

sublevel

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Importance of Valence e-?

• Properties of elements and

compounds/molecules dictated by the

number of valence electrons.

• Similar properties of all elements in groups, similar formulas for

compounds/mole cules in those

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Goal number of Valence

Electrons?

• Which group of elements is stable?

• How many VE to all elements have in this group?

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Trends with Ions?

• Atoms will gain or lose as many electrons needed to have 8 valence electrons

• Lose electrons: positive ions, metals

• Gain electrons: negative ions, nonmetals

*Because of ions, formulas for compounds should be similar with elements in the

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Modern Application

• Solar panels

*n-type: adding an atom w/ one more e- than Si (n = negative charge carrying) *p-type: adding an atom w/

one less e- than Si (p = positive charge carrying) *Junctions between the two

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Which of the following ions has the same number of electrons as Br - ?

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-AP Practice Question

All the chlorides of the alkaline earth metals have similar

empirical formulas, as shown in the table above. Which of the following best helps to explain this observation?

A.Cl₂(g) reacts with metal atoms to form strong, covalent double bonds.

B.Cl has a much greater electronegativity than any of the alkaline earth metals.

C.The two valence electrons of alkaline earth metal atoms are relatively easy to remove.

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AP Practice Question

RbCl has a high boiling point. Which of the following

compounds is also likely to have a high boiling point, and why?

A.NO, because its elements are in the same period of the periodic table.

B.ClF, because its elements are in the same group of the periodic table.

C.Cl₂O, because its elements have similar

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AP Practice Question

If Na reacts with chlorine to form NaCl, which of the following elements reacts with Na to form an ionic compound in a one-to-one ratio, and why?

A.K, because it is in the same group as Na.

B.Mg, because its mass is similar to that of Na. C.Ar, because its mass is similar to that of Cl.