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Chapter 6Periodic Table & Periodic Law
6.1 Development of the Modern Periodic Table
6.2 Classification of the Elements
6.3 Periodic Trends
Section 6.1 Development of the Modern Periodic Table
• Trace the development of the periodic table from the law of octaves through the current table ordered by atomic number, including the scientists who contributed to each stage of development.
• State the periodic law.
• Identify key features of the periodic table.
• Explain the common feature of elements within a group.
The periodic table evolved over time as scientists discovered more useful ways to compare and organize the elements.
Section 6.1 Development of the Modern Periodic Table
• Identify the portion of the periodic table the terms “representative elements”, “transition metals”, “inner transition metals”, “alkali metals”, “alkaline earth metals”, metalloids, halogens, “noble gases”, “lanthanide series”, “actinide series”, and “transuranium elements” refer to and be able to give examples of some characteristics of the elements found in these regions.
• State the number of naturally occurring elements on Earth, the total number of elements that are currently formally recognized as existing, and the names of the 2 most recently recognized elements.
Section 6.1 Development of the Modern Periodic Table
• State the typical characteristics of metals, nonmetals and metalloids and be able to give an example of an element in each of these categories
• Identify the states and colors of the 4 halogens at room temperature and describe the trend in reactivity among the halogens.
• Identify what is different about copper, gold, and mercury compared with other transition metals.
Key Concepts• The elements were first organized by increasing atomic
mass, which led to inconsistencies. Later, they were organized by increasing atomic number.
• The periodic law states that when the elements are arranged by increasing atomic number, there is a periodic repetition of their chemical and physical properties.
• The periodic table organizes the elements into periods (rows) and groups (columns); elements with similar properties are in the same group.
Section 6.1 Development of the Modern Periodic Table
Key Concepts• Elements are classified as either metals, nonmetals, or
metalloids.
Section 6.1 Development of the Modern Periodic Table
History of Development
John Newlands (~ mid 1860s)
• Octave rule for elements ordered by atomic mass
History of Development
Lothar Meyer, Dimitri Mendeleev• Both demonstrated relationships
between elemental properties and atomic mass
• Mendeleev given more credit for idea• Published first table in 1872
Dimitri Mendeleev
First Periodic Table • Ordered by atomic mass
Predicted existence/properties of undiscovered elements
• Scandium (Sc)• Gallium (Ga)• Germanium (Ge)
Prediction of Germanium PropertiesProperty Predicted
Eka-Silicon(1871)
ObservedGermanium
(1886)Atomic Mass 72 72.6
Density, g/cm3 5.5 5.47
Color Dirty gray Grayish white
Dens. Oxide EsO2: 4.7 GeO2: 4.703
BP of chloride EsCl4: < 100 C
GeCl4: 86 C
Dens. Chloride EsCl4: 1.9 GeCl4: 1.887
Henry Moseley (~1913)Known that some elements in wrong order
Used term atomic number (AN) to indicate amount of charge in nucleus (determined this charge from positions of spectral lines)
6 years prior to proton discovery
Arrangement by AN fixed periodic table problems
Contributions to Classification of the Elements (Table 6.2)
positive charge
Periodic Law
There is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number
Elements – Periodic TableOrdered by atomic number (number of protons in nucleus)
Columns (vertical) = groups or families Current IUPAC numbering system is 1-18
Elements in same group tend to have similar chemical and physical properties
Rows (horizontal) = periods (currently 7)
Table “periodic” because pattern of variation of chemical and physical properties repeats in each period
Elements - NewestElements with atomic numbers 114 and 116 officially named 5/2012 as Flerovium (Fl) and Livermorium (Lv), respectively. Copernicium (Cn), atomic number 112 officially named 2/2010
Elements 113, 115, 117 (newest – April 2010), 118 have claimed to have been made, but evidence not yet convincing enough for official recognition by IUPAC
Official current total = 114 elements
Periodic Table of ElementsEach box shows atomic number and the element’s symbol
Newest elements Fl & Lv; remainder claimed to have been made but have not been officially
recognized as existing
Periodic Table – Fig. 6.5
Representative Elements
Modern Periodic Table
Representative Elements• Groups 1-2, 13-18• Wide range of chemical & physical
properties
Transition Metals• Groups 3 to 12
General Classifications
Metals
Nonmetals
Metalloids
ClassificationsMetals• Shiny when smooth
& clean• Solid at room
temperature• Good conductors of
heat & electricity• Ductile• Malleable
Metal ClassificationsAlkali Metals – Group 1Li, Na, K, Rb, Cs, FrVery reactive, soft, mostly exist as compounds with other elements
Metal Classifications
Alkaline Earth Metals – Group 2
Be, Mg, Ca, Sr, Ba, Ra
Also reactive, but not as much as the alkali metals
Ca & Mg important nutrients (as ions, not as elements)
Mg alloys useful as light-weight materials (bikes, laptop cases)
Transition and Inner Transition Metals
Inner transition metals: Lanthanide & Actinide Series
Metal Classifications
Transition Metals / Inner Transition MetalsGroups 3 to 12Inner transition metals consist of two parts
• Lanthanide series - Ce through Lu• Actinide series – Th through Lr• All elements past U (AN 92) are
synthetic (man made) Called transuranim elements
Some Transition Metals
Cr
Only transition metals not having a silver/gray color
An Atypical Transition Metal - Mercury
Only transition metal that is liquid at room temperature
Metalloids
Semi-Metals / Metalloids
The “staircase” – dividing line• Steps start between boron and aluminum• Elements on either side of the dividing line
are metalloids except for Al• Some texts do not include Po• Most do not include At – highly radioactive,
estimated that total amount in Earth’s crust <30 g at any time – hard to study
Silicon & germanium most important
Nonmetals
Nonmetals
Gases or brittle, dull-looking solids
Poor conductors of heat and electricity
Group 16 Nonmetals
O, S, Se, Te
Sulfur Selenium
Group 17 (Halogens): F, Cl, Br, I All diatomic (F2, etc)
States @ RT: F2(g), Cl2(g), Br2(l), I2(s)
Colors: colorless, pale greeen, dark red-brown, very dark violet (almost black)All reactive, but F2 most, I2 least
Group 18 - Noble Gases
He, Ne, Ar, Kr, Rn
Unreactive
Used in lasers, light bulbs, signs and certain types of welding (TIG, argon)
Chapter 6Periodic Table & Periodic Law
6.1 Development of the Modern Periodic Table
6.2 Classification of the Elements
6.3 Periodic Trends
Section 6.2 Classification of the Elements
• Explain why elements in the same group have similar properties.
• Identify the four blocks of the periodic table based on their electron configuration.
• Identify the two exceptions in normal orbital filling order in the period 4 transition metals.
Elements are organized into different blocks in the periodic table according to their electron configurations.
Key Concepts
• The periodic table has four blocks (s, p, d, f).
• Elements within a group have similar chemical properties.
• The group number for elements in groups 1 and 2 equals the element’s number of valence electrons.
• The energy level of an atom’s valence electrons equals its period number.
Section 6.2 Classification of the Elements
Periodic Table OrganizationBasic organization of periodic table by Mendeleev (~1872) was by recurring trends in elemental properties
Electron not discovered until 1897
Schrodinger model atom was ~1927
Agreement between property based table and electron configuration based table demonstrates influence of electron configuration on properties
Organizing Elements by Electron Configuration
Group 1 – see table 6.3
Period
Element
Electron Configuration
N. Gas Config
1 H 1s1 1s1
2 Li 1s22s1 [He]2s1
3 Na 1s22s22p63s1 [Ne]3s1
4 K 1s22s22p63s23p64s1 [Ar]4s1
Organizing Elements by Electron Configuration
Group 1 – see table 6.3
Period
Element N. Gas Configuration
# Valence Electrons
1 H 1s1 1
2 Li [He]2s1 1
3 Na [Ne]3s1 1
4 K [Ar]4s1 1
Organizing Elements by Electron Configuration
Elements in a group have similar chemical properties because they have the same number of valence electrons
• Group 1 ns1 1• Group 2 ns2 2• Group 13 ns2np1 3• Group 14 ns2np2 4
Organizing Elements by Electron Configuration
# of valence electrons = # of electrons in highest principal energy level= group number (groups 1 and 2)= group number – 10 (groups 13 to 18)
He (18) exception (2 valence electrons)
For transition metals, # valence electrons can vary and may not be simply related to the group number
Organizing Elements by Electron Configuration
By definition, energy level of an element’s valence electrons = period of element
Ga [Ar]3d104s24p1 period 4
s-, p-, d-, and f- Block Elements
Block: section of the periodic table that corresponds to the energy sublevel being filled with valence electrons
• s, p, d, and f sublevels
s-block elements
Groups 1 and 2
• 1 s1
• 2 s2
Can only have 2 s-block groups because s sublevel only holds 2 electrons
p-block elements
Groups 13 to 18 (except He)• s sublevel is already filled (s2)• 13 p1
• 14 p2
• 15 p3
• 16 p4
• 17 p5
• 18 p6 completely filled s & p - very stable
Representative Elements
The representative elements are thes- and p- block elements
d-block elements
Same as transition metals
In period n• ns2 (filled s sublevel)• Partially filled or filled d orbitals of level
(n-1)• d sublevel can hold 10 electrons; d-
block spans 10 groups on periodic table
d-block elements
Period 4 - Filling the n-1 d sublevel
• Sc [Ar]3d14s2
• Ti [Ar]3d24s2
Filling is more or less regular but exceptions occur
• Cr and Cu irregular because of stability of filled and half-filled d sublevel
Period 4, d Block Exceptions
Aufbau diagram works to vanadium, AN 23
Half-filled and fully-filled set of d orbitals have extra energy stability, so chromium is
Cr [Ar]3d54s1 (1/2 filled d)
Not [Ar]3d44s2
Next exception is copper:
Cu [Ar]3d104s1 (filled d)
Not [Ar]3d94s2
f-block elementsf sublevel can hold 14 electrons; f-block spans 14 columns of periodic table
Electrons don’t fill orbitals in a predictable manner
f-block Inner Transition MetalsFor period n (n = 6 or 7)• Filled ns (ns2); Filled or partially filled (n-2) f (4f
& 5 f)
• First member of lanthanides (La, AN 57) & actinides (Ac, AN 89) not in Aufbau order – have d1 configuration, expect f1
• f sublevel filled one element prior to last member of either row
Aufbau Order & Periodic TableCan “read” Aufbau order (page 160) directly from periodic table and knowledge of how blocks fill
Move left to right
Keep period number for representative elements (groups 1-2, 13-18)
(Period number -1) for d block
(Period number -2) for f block (start after La or Ac to get closer to actual order)
s, p, d, and f blocks
s, p, d, and f blocks
s block
d block
p block
f block
Practice
Practice problems, page 162
Problems 7-9
Section assessment, page 162
Problems 10-15
Chapter 6Periodic Table & Periodic Law
6.1 Development of the Modern Periodic Table
6.2 Classification of the Elements
6.3 Periodic Trends
Section 6.3 Periodic Trends
• Compare period and group trends of several properties.
• Explain the meaning of electronegativity
• Relate period and group trends in atomic radii, ionic radii and electronegativity to electron configuration.
• Describe the roles of electron-electron repulsion, electron-nucleus attraction, shielding (effective nuclear charge), and the added stability of favored electron configurations (octet, half filled and filled sublevels) in determining periodic property trends.
Trends among elements in the periodic table include their size and their ability to lose or attract electrons
Section 6.3 Periodic Trends
• Use Coulomb’s law to explain shielding and its impact on atomic radii.
• Predict and explain the change in size that occurs when an atom forms either a cation or an anion.
• Predict and explain the relative changes in ionization energy for the first, second, third, etc. ionizations of a given atom.
• Explain how departures from overall first ionization energy trends for some elements can occur in terms of the specific electron configurations of these elements.
Trends among elements in the periodic table include their size and their ability to lose or attract electrons
Key Concepts
• Atomic and ionic radii decrease from left to right across a period, and increase as you move down a group.
• Core electrons are effective at shielding valence electrons whereas other valence electrons are not effective shielders.
• Ionization energies generally increase from left to right across a period, and decrease as you move down a group.
• The octet rule states that atoms gain, lose, or share electrons to acquire a full set of eight valence electrons, which is a particularly stable electron configuration. Filled and half-filled sublevels (particularly d) also impart extra stability.
• Electronegativity generally increases from left to right across a period, and decreases as you move down a group.
Section 6.3 Periodic Trends
Key Concepts
• Electronegativity is a property related to bonding and therefore is a measure of an interaction of the atom with another atom.
• Disruptions of filled energy levels and of filled sublevels require more energy than removing electrons from other electron configurations.
Section 6.3 Periodic Trends
Periodic Trends - Principles
Negative electrons are attracted to the positive nucleus
Coulomb’s Law: F (+q) (-q) / r2
F = attractive force
+q = charge on nucleus
-q = electron charge
r = distance between charge centers
Force vs distance for 1/r2
0
0
0
0
0
0
0
0
0
5 15 25 35 45
r (distance)
Fo
rce
Periodic Trends - Principles
Coulomb’s Law: F (+q) (-q) / r2
The greater the nuclear charge (+q), the more strongly an electron is attracted
The closer an electron is to nucleus (smaller r), the more strongly it is attracted
Coulomb’s Law: F (-q)(-q)/r2
Force repulsive if charge same sign
Electrons repelled by other electrons in an atom
The further away two electrons are from each other, the weaker the repulsive force between them
Periodic Trends - Principles
Electrostatic Forces in Atom
If other electrons are between a valence electron and nucleus, valence electron will be less attracted to nucleus
Periodic Trends - Shielding
+3 F1 (+3)(-1) / r2
+3 F2 (< +3)(-1) / r2
Valence electron Inner (core) electron
r
Both forces are attractive, but F2 is smaller than F1 due to shielding
If other electrons are between a valence electron and nucleus, valence electron will be less attracted to nucleus – nuclear charge is shielded
Effective nuclear charge less than full nuclear charge due to shielding of charge by negative core electrons
Other valence electrons not effective at shielding a valence electron
Periodic Trends - Shielding
Shielding Effect – Mg: [Ne]3s2
RED – attractive forces BLUE – repulsive forces
Nucleus
ValanceElectron
CoreElectronCloud
Represent-ativeCore
Electron
Filled principal energy levels are very stable (noble gas configuration)
Atoms prefer to add/subtract/share valence electrons to completely fill a principal energy level if possible (octet rule)
Completely filled sublevels (s, p, d) also have extra stability
Periodic Trends - Principles
Atomic RadiusSizes mostly obtained from crystalline form of element or from diatomic moleculeIn general, are averages of internuclear separations observed in a variety of substancesTypically expressed in pm = picometer = 10-12 mMay also see in nm = nanometer = 10-9 m
Atomic Radius
Sodium in crystal
Radius
372 pm
186 pm
Bonded sodium atoms
Atomic RadiusHydrogen in gaseous diatomic molecule
Bonded hydrogen
atoms
Radius
74 pm
37 pm
Atomic Radius Trends: Figure 6.11
Atomic Radii Trends: Figure 6.12
Group trend – more shielding, more distant orbitals with increasing nPeriod trend – increasing effective nuclear charge
Trends – Atomic RadiusPeriod trend (left to right) – dominated by increasing effective nuclear charge with no change in n and with little additional shielding provided by valence electrons
Group trend (top to bottom) – dominated by increasing n (orbitals more distant) and shielding of nuclear charge by core electrons (from filled energy levels) – effective nuclear charge decreases
1s, 2s, 3s Orbitals – Distance From Nucleus
1s 2s 3s
Node Nodes
1s
2s3s
Trends – Ionic RadiusIon is charged atom - electrons have been added or removed
• Positive ion (+ ion) formed if electrons removed – called cation
• Negative ion ( ion) formed if electrons added – called anion
Trends – Ionic Radius: CationsWhen atoms lose electrons and form positively charged ions (cations), they always become smaller for 2 reasons:
1. The loss of a valence electron can leave an empty outer orbital resulting in a small radius
2. Electrostatic repulsion decreases allowing the electrons to be pulled closer to nucleus
Trends –Ionic RadiusIonization of elemental sodium
Na Na+ + e-
Sodium atom Sodium ion
[Ne]3s1 [Ne]
Trends – Ionic Radius: AnionsWhen atoms gain electrons and form negatively charged ions (anions), they always become larger because electrostatic repulsion increases, causing the electrons to spread apart
If added electron placed in previously unoccupied energy level, then average distance from nucleus will be greater for that electron
Trends –Ionic Radius
Ionization of atomic chlorine
Cl + e- Cl-
Chlorine atom Chlorine ion
[Ne]3s23p5 [Ne]3s23p6 or [Ar]
Ionic Radii in Crystal
Ionic Radius – Magnesium oxide
Trends – Ionic Radius
Formation of + ions always results in decrease in radius
Formation of - ions always results in increase in radius
Electron configuration of ions follows same filling order as for neutral atoms
Within a set of + ion or – ions, radius trends for groups/periods are the same as for neutral atoms
Ionic Radius TrendsSee figure 6.14
Ionic Radius Trends – Fig. 6.15
Within a category (+ or – ions), trend is same as trend for atomic radius
Comparison of Atomic & Ionic Radii
Ionization Energy (IE)Energy needed to remove electron from neutral gaseous atom to form + ion
X(g) X+(g) + e- kJ/mol • Ionization potential is per atom value in eV (1 eV=1.602 x 10-19 J)
Must overcome electrostatic attraction to remove electron
Low IE = easy valence electron loss – element will readily form a cation
2nd Ionization EnergyEnergy needed to remove electron from ion with single positive charge
X+ (g) X2+(g) + e- kJ/mol
Must overcome much stronger electrostatic attraction to remove second electron compared to the first electron
Trends – Ionization EnergyGroup 1 lowest values and Group 18 highest values within a period
Group 1 likely to form M+1 but unlikely to form M+2
• Difficult to remove electron from noble gas configuration (filled p sublevel)
Trends – Ionization Energy (IE)
Period and group trends inverse of radii trends – increased nuclear charge pulls electron in tighter
Shape of trend for any given period is irregular due to stability of filled / half filled sublevels
Trends – 1st Ionization Energy Periods 1-5 Fig. 6.16
Trends – 1st Ionization Energy Fig. 6.17
Trend opposite that for atomic radiusReasons for trend same as for atomic radius
Trends – Ionization Energy (IE)2nd IE always > 1st
3rd IE always > 2nd etc.• Removing electron from more + ion
Once valence electrons removed, energy always takes big jump
• Must remove electron from filled level
Period 2 Successive Ionization Energies (Table 6.5)
Trend Exceptions – (IE)1st IE (values in kJ/mol)
Be (1s22s2) 900 vs B (1s22s22p1) 800
• B has inner 2s2 configuration which effectively shields 2p1 electron
N (1s22s22p3) 1400 vs O (1s22s22p4) 1310
• Half-filled p sublevel has extra stability
Trend Exceptions – IE
2nd IE
B (1s22s22p1) 2430 vs C (1s22s22p2) 2350 • 2s2 effectively shields
O (1s22s22p4) 3390 vs F (1s22s22p5) 3370 • Disrupts/creates 2p3 configuration
Trends - Electronegativity
Relative ability to attract electrons in a chemical bond
Max 3.98 (F) to min 0.7 (Fr)
Elements with high EN tend to form negative ions
• F-, Cl-, O2-
Noble gases not tabulated• Very few compounds to get info from
Trends - Electronegativity
Increasing Electronegativity
electronegativity < 1.0
1.0 electronegativity < 2.0
2.0 electronegativity < 3.0
3.0 electronegativity < 3.9
electronegativity 3.9
Dec
reas
ing
Ele
ctro
nega
tivity
Trends – Electronegativity vs AN
Trend Summary (ignore EA)