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