Chapter 6 Periodic Table & Periodic Law 6.1Development of the Modern Periodic Table 6.2...

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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.

• 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.

• 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.

Key Concepts• Elements are classified as either metals, nonmetals, or

metalloids.

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

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

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)

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

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

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

# 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

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 block

d block

p block

f block

Practice

Practice problems, page 162

Problems 7-9

Section assessment, page 162

Problems 10-15

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

• 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.

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.

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

5 15 25 35 45

r (distance)

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

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

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

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

Radius

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

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

electronegativity 3.9

tivity

Trends – Electronegativity vs AN

Trend Summary (ignore EA)

Chapter 6 Periodic Table & Periodic Law 6.1Development of the Modern Periodic Table 6.2...

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