Periodic Table and Periodic Trends

Test 2 MaterialAlong with Atomic Structure:

16 out of 75 MC questions

Free Response every year

Broad Periodic Table Classifications

Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)

Transition Elements: filling d orbitals (Fe, Co, Ni)

Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

Periodicity Trends in atomic radii, ionization

energies, and electron affinities are all tied to two basic structural features of the atoms: Effective nuclear charge (Zeff) Principal energy level of valence

electrons

Periodic Trends Explained in terms of:

Increasing/Decreasing nuclear charge (Z)

Increasing/Decreasing distance from nucleus (n)

Increasing/Decreasing shielding by core electrons

Ionization EnergyThe quantity of energy required to remove an electron from the gaseous atom or ion—form a positive ion.

Low ionization energy—electron is easy to remove

High ionization energy—electron is difficult to remove

Periodic TrendsFirst ionization energy:

increases from left to right across a period; higher nuclear charge

decreases going down a group; electrons farther from nucleus

Values of first ionization energy for the elements in the first six periods

Page 326

First Ionization Energy Notice that the trend is not

smooth. First ionization energy decreases

between s- and p- orbital filling and just after p-orbital is half full.

Extra stability at these points

Trends in ionization energies (kJ/mol) for the representative elements.

p. 326

Beyond First Ionization Each successive electron is harder

to remove—large + charge in relation to # of electrons

Large increase in ionization energy occurs when going to a lower energy level (n changes)

12Page 325

Electron AffinityThe energy change in energy associated with the addition of an electron to a gaseous atom.

X(g) + e X(g)

Electron Affinity More exothermic for atoms that

easily accept electrons; energy is more negative

Generally becomes more negative from left to right across a period—several exceptions such as adding to a 2p orbital with one electron already there (N vs C)

Electron affinity values for atoms among the first 20 elements that form stable, isolated X- ions.

Trend in Electron Affinity Becomes less exothermic down a

group

Periodic TrendsAtomic Radii:

decrease going from left to right across a period; more positive charge in nucleus

increase going down a group; more energy levels

The radius of an atom (r) = half the distance between the nuclei in a molecule consistingof identical atoms.

Atomic radii (picometers) for selected atoms.

p. 330

Trends in Transition Elements

Although there is a slight contraction at the beginning of the series, the atoms are all approximately the same size.

Sc Ti CoFeMnCrV Ni

Ions Formed Varies by group Very predictable for representative

elements +1 +2 +3 +/-4 –3 –2 –1 0 Lose or gain electrons to reach a

Noble gas configuration

Sizes of Ions Size increases down a group; more

energy levels In isoelectric ions, size decreases as

positivity increases; more + nucleus holds electrons closer.

Positive ions are much smaller than negative ions—nuclear attraction; less repulsion

All in the Family

Group 1A—Alkali Metals 1 valence electron Very reactive—very close to stable

configuration Lose electron to form +1 ion Donate electrons—reducing agents Reactivity increases down the group

Group 2A—Alkaline Earth Metals 2 valence electrons Very reactive, but less than alkali

metals—not as close to stable config. Lose 2 electron to form +2 ion Donate electrons—reducing agents Reactivity increases down the group

Group 3A 3 valence electrons (ns2, np1) Typically forms +3 ions Increasingly metallic down the

group Forms bonds with more covalent

character than groups 1 or 2

Transition Metals Show more similarity than rep. ele.

both within periods and groups Electrons added in d & f orbitals

are actually inner electrons & cannot bond as easily as the outer s & p electrons—chemical properties do not change as much.

Transition Metals—General Properties

Typical metals: shiny, good conductors, flexible (malleable & ductile)

Melting points, hardness, and reactivity vary, so transition metals have varied uses.

Uses forTransition Metals

Hg—low melting point + good conductivity—thermometers and thermostats

W—high melting point + ductility—filaments in lightbulbs

Uses forTransition Metals

Fe & Ti—strength + hardness—structural materials

Cu, Ag & Au—soft—jewelry and art

Cu—low resistance—wiring

Reactions with Transition Metals

Most react with O2 to form oxides Cr, Ni, Co—oxides protect surface Fe—oxide flakes off exposing

more metal to oxidation Au, Ag, Pt & Pd—Noble metals—

do not readily form oxides

Ionic Compounds with Transition Metals Often paramagnetic (unpaired e-s) Usually more than one oxidation

state Cations are often complex ions in

which ligands (Lewis base ions or molecules) surround metal ion

Complex ion

Ionic Compounds with Transition Metals

Most are colored because ions absorb specific wave-lengths of light.

Electrons in split d orbitals can rearrange.

Wulfenite—PbMoO4

Rhodochrosite—MnCO3

Electron Configuration Full (n)s orbital; filling (n-1)d

orbital; empty (n)p orbital Exception: Cu—4s1, 3d10

Breaking the Rules Since the energy of a 3d orbital is

less than that of a 4s orbital, ions formed from first row transition metals lose their 4s electrons first

Ex: Mn = [Ar] 4s2, 3d5

Mn2+ = [Ar] 3d5

Lanthanides & Actinides Technically part of the transition

metals Filling f orbitals with electrons Rare metals with few uses All actinides are radioactive—

nuclear fuels (U & Pu)

Bonding

8 out of 75 M/C Questions

Free Response—Every year

In General: All bonds occur because of

electrostatic attractions. Formation of molecules and the

state of matter of a substance depends on the attractions between electron clouds of one atom and nucleus of another atom.

Bonding—General Rules A metal and a nonmetal bond so

that charges on the ions cancel. When two nonmetals react to form

a covalent bond, they share electrons in a way that gives both atoms a Noble gas configuration.

Bonding—General Rules nonmetal + representative metal

binary ionic compound: ions form to give the nonmetal the valence electron configuration of the next noble gas atom, and valence orbitals of the metal are emptied

Bonding—General Rules Most bonds are combinations of

ionic/covalent character The more different the atoms

bonding are, the more ionic character of the bond.

More similar—more covalent

Word Association Covalent Sharing Molecules Structural

formula

Ionic Transfer /

charged Compounds Formula Unit

Coulomb’s Law Describes energy of interaction

between ions E = 2.31 x 10-19 J*nm(Q1Q2 / r) Q1 & Q2numerical ion charges r distance between ion centers Negative ans. means ion pair is

more stable than individual ions.

Bond Length Distance between bonding atoms

at which energy is minimized Atoms position themselves to

minimize repulsions and maximize attraction & thus achieve lowest possible energy.

Predicting Formulas for Ionic Compounds Metal—positive charge equal to # of

valence electrons Nonmetal—negative charge equal to

# of electrons away from next Noble gas

Compound—charges must cancel

Predict Formulas for: Potassium sulfide Barium chloride Aluminum oxide Magnesium phosphide

Exceptions to the Rule: Sn—forms both +2 and +4 ions Pb—forms both +2 and +4 ions Bi—forms +3 and +5 ions Tl—forms +1 and +3 ions “no simple explanation for this

behavior”

Lattice Energy Change in energy taking place

when separated gaseous ions are packed together to form an ionic solid

Energy released when an ionic solid forms from its ions

Calculating Lattice Energy The sum of energy changes in

each step Formation of lithium fluoride:

1) Sublimation of solid lithium: 161 kJ per mole

2) Ionization of lithium: 520 kJ per mole

Calculating Lattice Energy3) Dissociation of F2: 77 kJ per

mole of F atoms

4) Ionization of F: -328 kJ per mole (Electron affinity of F)

5) Formation of LiF: -1047 kJ per mole

Calculating Lattice Energy Sum:

161 + 520 + 77 – 328 – 1047 = -617 kJ per mole LiF

(negative sign means process is exothermic & product has

lower energy)

Calculating Lattice Energy Lattice energy can also be calculated

using a form of Coulomb’s Law:

Lattice energy = k(Q1Q2 / r)

k = proportionality constant dependent on solid’s structure & electron configuration of the solid

P. 365—energy changes in forming LiF and MgO

Lattice Energy Much larger changes in Mg and

O because of greater Q1 and Q2

Even though Mg and O must ionize twice (which requires lots of energy) the overall change is very exothermic

H from Bond Energies Analyze reaction for number of

bonds broken and number of bonds formed.

In general, breaking bonds uses energy (+), and forming bonds releases energy (-).

Find the sum.

Example Problem—p. 373

Polar Covalent Bonds Most bonds have both ionic and

covalent character Polar covalent bonds—unequally

shared electrons Not different enough to be purely

ionic nor similar enough to be purely covalent

Polar Covalent Bonds Strength of polarity depends on

electronegativity (ability of an atom to attract electrons to itself).

Higher electronegativity results in higher negative charge.

Partial charge results in a stronger bond & stronger inter-molecular attraction.

Dipoles + indicates direction of

dipole—points to negative end Dipolar molecules orient

themselves one way in an electric field.

Dipoles can cancel.

Canceling Dipoles If evenly spaced identical bonds

exist, they cancel each other. Linear—2 identical bonds separated

by 180* (CO2) Trigonal planar—3 identical bonds

separated by 120* (SO3) Tetrahedral—4 identical bonds

separated by 109.5* (CH4)

Ionic Bonds with Covalent Character No totally ionic bonds exist. Percent ionic character can be

calculated: (measured dipole moment of X-Y)

(calculated dipole moment of X+Y-)X 100%

Ionic Bonds with Covalent Character X-Y represents the molecule X+Y- represents the ions Graph—p. 367—Bonds with

more than 50% ionic character are classified as ionic

New Definition Ionic compound—any substance

that conducts electricity in its liquid state (melted, not dissolved)

Beyond the Bond—Intermolecular Forces & States of Matter

Intermolecular Forces

Weaker than chemical bond

Affect structure and state of matter

Dipole-Dipole Forces Positive and negative ends of polar

molecules attract each other. About 1% as strong as covalent or

ionic bonds Weaken as distance between

molecules increases

Hydrogen Bonding Especially strong dipole-dipole

force Occurs when H bonds to a strongly

electronegative atom—O, N, or F Very strong because 1) molecule is

very polar & 2) small size of H

H Bonding Example—water More pronounced in molecules

formed from small atoms (dipoles can come closer)

High boiling point

London Dispersion Forces Forces that exist among Noble gases

and nonpolar molecules Results from temporary dipoles

formed when electrons distribute themselves unevenly—can induce a dipole in a neighboring atom

VERY WEAK

States of Matter Gases—weak IM forces (like

London Dispersion Forces) Liquids—intermediate IM

forces Solids—strong IM forces

Liquids & IM Forces Surface tension—result of IM

forces that resist an increase in surface area

Capillary action—result of cohesive forces within liquid and adhesive forces between liquid and tube

Liquids (cont’d) Both effects are higher with

more polar molecules.

Solids (Crystals) Ionic solids—made of charged

particles; ions at lattice points Molecular solids—made of neutral

particles; molecules at lattice points Atomic solids—made of neutral

particles; atoms at lattice points; 3 types

An atomic solid, an ionic solid & a molecular solid

Ionic Solids Ions at lattice points Closest packed spheres Arranged to minimize repulsions

and maximize attractions Conducts only when melted

Molecular Solids Lattice positions occupied by

molecules Internal covalent bonds are strong, but

intermolecular forces are weak IM force: dipole/dipole if polar

covalent bond; London dispersion forces (larger in larger molecules)

Atomic Solids1. Network—directional covalent

bonds; forms giant molecules (diamond, graphite,and silicon); highest melting points

2. Metallic—delocalized covalent bonds; atoms have closest packing structure; high melting points

Atomic Solids3. Group 8A—Noble gases—

London dispersion forces only; low melting points.

Network Atomic Solids Strong, directional bonds Form giant “molecules” Typically brittle & poor conductors Examples—carbon and silicon

Carbon Network Follows a molecular orbital (not

atomic orbital) model

Diamond Tetrahedral--sp3 hybridized

bonds stabilize structure Large gaps exist between filled

and unfilled molecular orbitals—hard for electrons to move—no conductivity

Graphite Fused carbon rings form sheets Trigonal planar—sp2 hybridized

(1 p orbital remains unhybridized)

Delocalized electrons in orbital causes graphite to be conductive

Figure 10.22: The structures of diamond and graphite. In each case only a small part of the entire structure is shown.

Closest Packed Solids aba pattern—alternating layers—

atoms in 3rd layer lie directly above atoms in 1st layer—hexagonal unit cell

abca pattern—atoms in 1st and 4th layers are in line; 2nd & 5th layer; 3rd & 6th layer—face-centered cubic cell

aba Packing

abca Packing

Density of Closest Packed Solids

Figure out how many atoms are in each cell.

Use atomic radius to determine volume of cell.

Convert from atoms to moles and then from moles to grams to get mass.

Face-Centered Cubic Unit Cell

Bonding in Metals Strong, non-directional bonds Atoms are hard to separate but

easy to move. “Electron sea” model Mobile electrons carry heat or

electricity easily

Band Model or Molecular Orbital Model Electrons travel around metal

crystal in a molecular (instead of atomic) orbital

Result is a continuum of levels that eventually merge to form a band.

Figure 10.19: The molecular orbital energy levels produced when various numbers of atomic orbitals interact.

Band or MO Model Empty orbitals close in energy

exist. Electrons are very mobile into

and out of these similar-energy orbitals.

Alloys Introducing other elements into

metallic structure is easy. Substitutional—some “host” atoms

are replaced by atoms of similar size Interstitial—small atoms occupy

spaces between metal atoms (such as carbon used to harden steel)

Molecular Models

Localized Electron Model Assumes a molecule is composed

of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms

Lone pairs—localized on one atom Bonding pairs—between atoms

Three Aspects of Model Description of valence electron

arrangement (Lewis Structure) Prediction of geometry

(VSEPR) Description of types of orbitals

(Ch. 9--hybridization)

Localized Electron Model—Part 1

Lewis Structures

Lewis Structures“From experiment, chemists have

learned that the most important requirement for the formation of a stable compound is that the atoms achieve Noble gas electron configurations.”

Writing Lewis Structures Include only valence electrons!

1. Sum all valence electrons—doesn’t matter which atom they come from

2. Use pairs of electrons to bond atoms

3. Use remaining electrons to satisfy duet rule for H and octet rule for all others.

H O H O=C=O N N8 electrons 16 electrons 10 electrons

Exceptions to the Octet Rule1. B and Be often have fewer than 8

electrons when in compounds—very reactive

2. Third row and heavier elements can exceed the octet rule by using their empty d orbitals.

Summary: C, N, O, & F always obey Second row elements can never

exceed the octet rule When writing Lewis structures,

satisfy the octet rule first & then place extras on atoms with d orbitals

Resonance Structures Molecular structures with more than

one possible arrangement—usually relates to position of double bonds

Supported by experimental evidence—bonds are identical in length and strength

If bonds are localized, there would be a difference (DB shorter & stronger)

Localized Electron Model—Part 2

VSEPR Theory--Geometry

VSEPR Theory Valence Shell Electron Pair Repulsion Good predictions for approximate 3-D

arrangements of molecules Main postulate: The structure around

a given atom is mainly determined by minimizing electron pair repulsions.

What does it mean?? Binding atoms will arrange

themselves as far from other binding atoms as possible (maximum angle)

Steps for Using VSEPR1. Draw Lewis Structure2. Count electron pairs & arrange

them as far from each other as possible.

3. Determine positions of atoms from the way pairs are shared.

4. Name the shape.

TetrahedralTrigonal

Pyramidal Bent

Notice that all atoms have 4 influences on the shape (bonds and/or unshared pairs). More unshared pairs results in a smaller angle between bonded atoms.

Second Postulate of VSEPRLone pairs require more room

than bonding pairs and tend to compress the angles between the bonding pairs.

Figure 8.18: (a) In a bonding pair of electrons, the electrons are shared by two nuclei. (b) In a lone pair, both electrons must be close to a single nucleus and tend to take up more of the space around that atom.

Arrangement c is preferred because there are larger bond angles between the lone pairs.

Multiple Bonds & VSEPR Multiple bonds count as one

effective pair of electrons. Use any of a molecule’s

resonance structures to predict its geometry.

PRACTICE Draw a Lewis Structure and

predict the shape of PCl6-

Predict the shape of XeF4.

Octahedral structure when all electrons are included, but square planar when only bonded atoms are included.—Square planar geometry

Hybridization

Example CH4—H only has s orbitals; C has one s

orbital & 3 p orbitals Should have 2 different types of bonds (s

—s s—p) Should have 90* bond angles for p

orbitals experimental evidence for four identical

bonds at 109.5*

Carbon’s “native” orbitals

P orbitals are perpendicular to one another.

Explanation Carbon’s orbitals combine One s orbital and 3 p orbitals:

sp3 hybridization—four equal orbitals with one large and one small lobe

Carbon’s sp3 Hybridization—p. 415

Energy The energy of the hybrid

orbitals is between the energies for the native orbitals.

Fig. 9.24—p. 424

# of effective pairs

vs. arrangement

vs. hybridization

Types of Hybridization sp3—s and all p orbitals hybridize

—tetrahedron sp2—s and 2 p orbitals hybridize

—trigonal planar sp—s and 1 p orbital hybridizes--

linear

sp2 hybridization

sp hybridization—p. 419

Types of Hybridization dsp3—s, all p orbitals, and 1 d

orbital hybridize—trigonal bipyramid

d2sp3—s, all p orbitals, and 2 d orbitals hybridize—octahedron

dsp3 Hybridization of Phosphorus

PCl5—Cl is sp3 hybridized

d2sp3 hybridization of S

Combining Localized Electron and Molecular Orbital Models

Why Both? Localized electron model is

simpler. Molecular orbital model’s

delocalization of electrons describes some molecules that L.E. model cannot except by resonance structures.

Double Bonds Explanation of double bonds is

the main area in which the combined models are useful.

Sigma bonds are localized. Pi bonds (2nd bond of double

bond) are delocalized.

Sigma Bonds in Benzene

Pi Bonding in Benzene Showing Delocalized Electrons