Chapter 4: Periodic Table of Elements

Overview

The periodic table is one of the most important tools in chemistry. It organizes all known elements in a systematic way that reveals patterns in their properties and reactivity. This chapter will guide you through the development of the periodic table, the organization of elements in the modern periodic table, and the characteristic properties of important groups and periods. Understanding the periodic table will help you predict chemical behavior and understand why elements behave the way they do.

Learning Objectives

After studying this chapter, you should be able to:

Explain the historical development of the periodic table
Understand the organization of the modern periodic table
Identify and locate different groups and periods
Analyze periodic trends in atomic properties
Describe characteristic properties of Group 1, Group 17, and Period 3 elements
Understand the properties of transition elements

4.1 Development of the Periodic Table

Early Attempts at Classification

Lavoisier's Elements (1789)

Approach: Classified elements as metals, nonmetals, and earths Limitations: Only identified 33 elements, no systematic organization

Dobereiner's Triads (1829)

Concept: Groups of three elements with similar properties Examples:

$\text{Cl, Br, I}$ (halogens)
$\text{Ca, Sr, Ba}$ (alkaline earth metals)
$\text{S, Se, Te}$ (chalcogens)

Law of Triads: Properties of middle element were average of the other two

Newlands' Law of Octaves (1864)

Concept: Elements arranged by atomic mass showed repeating properties every 8 elements

Comparison: Similar to musical notes repeating every 8 notes Limitations: Worked well for lighter elements but failed for heavier elements

Mendeleev's Periodic Table (1869)

Key Contributions

Atomic Mass Arrangement: Arranged elements by increasing atomic mass
Periodic Law: Properties are periodic functions of atomic mass
Prediction of Elements: Left blank spaces for undiscovered elements
Prediction of Properties: Predicted properties of missing elements

Achievements

Gallium (1875): Predicted eka-aluminum - discovered by French chemist
Scandium (1879): Predicted eka-boron - discovered by Swedish chemist
Germanium (1886): Predicted eka-silicon - discovered by German chemist
All predictions were remarkably accurate

Limitations

Could not explain position of isotopes
Could not explain properties of transition metals
Some elements didn't fit in proper positions

Modern Periodic Table (1913 onwards)

Moseley's Contribution (1913)

Key Discovery: X-ray spectroscopy revealed atomic numbers Law: Properties are periodic functions of atomic number, not atomic mass Result: Elements arranged by increasing atomic number

Modern Organization

Atomic Number: Number of protons in nucleus
Periods: Horizontal rows (7 periods)
Groups: Vertical columns (18 groups)
Blocks: s-block, p-block, d-block, f-block

Did You Know?

Mendeleev was so confident in his predictions that he corrected the atomic masses of certain elements, believing his periodic law was correct. When elements were discovered with properties matching his predictions, his periodic table gained scientific acceptance.

4.2 Arrangement of Elements in the Modern Periodic Table

Structure of the Periodic Table

Periods

Definition: Horizontal rows in the periodic table Number: 7 periods (1-7) Organization: Elements with same number of electron shells

Characteristics:

Period 1: 2 elements ( $\text{H, He}$ )
Periods 2-3: 8 elements each
Periods 4-5: 18 elements each
Periods 6-7: 32 elements each

Groups

Definition: Vertical columns in the periodic table Number: 18 groups (1-18) Organization: Elements with similar chemical properties and same number of valence electrons

Group Categories:

Main Group (Representative) Elements: Groups 1, 2, 13-18
Transition Elements: Groups 3-12
Inner Transition Elements: Lanthanides and Actinides

Blocks

s-block: Groups 1-2, Helium (last period)

Valence electrons: $ns^{1-2}$

p-block: Groups 13-18

Valence electrons: $ns^2np^{1-6}$
Mix of metals, metalloids, and nonmetals

d-block: Groups 3-12

Valence electrons: $(n-1)d^{1-10}ns^{0-2}$
Transition metals with variable oxidation states

f-block: Lanthanides and Actinides

Valence electrons: $(n-2)f^{1-14}$
Rare earth elements

Periodic Trends

Atomic Radius

Definition: Distance from nucleus to outermost electron

Trends:

Down a group: Increases (more electron shells)
Across a period: Decreases (increased nuclear charge, same shells)

Element	Period	Group	Atomic Radius (pm)
Li	2	1	152
Be	2	2	112
B	2	13	85
C	2	14	77
N	2	15	75
O	2	16	73
F	2	17	72
Ne	2	18	71

Ionization Energy

Definition: Energy required to remove one electron from a gaseous atom

Trends:

Down a group: Decreases (electrons farther from nucleus)
Across a period: Increases (increased nuclear charge, same shells)

Element	Period	Group	1st Ionization Energy (kJ/mol)
Li	2	1	520
Be	2	2	899
B	2	13	801
C	2	14	1086
N	2	15	1402
O	2	16	1314
F	2	17	1681
Ne	2	18	2081

Electronegativity

Definition: Ability of an atom to attract electrons in a chemical bond

Trends:

Down a group: Decreases (larger atomic size)
Across a period: Increases (increased nuclear charge)

Element	Period	Group	Electronegativity (Pauling)
Li	2	1	1.0
Be	2	2	1.5
B	2	13	2.0
C	2	14	2.5
N	2	15	3.0
O	2	16	3.5
F	2	17	4.0
Ne	2	18	-

Electron Affinity

Definition: Energy change when an atom gains an electron

Trends:

Down a group: Generally decreases
Across a period: Generally increases (with some exceptions)

Key Terms

Period: Horizontal row in periodic table
Group: Vertical column in periodic table
Atomic Radius: Size of atom
Ionization Energy: Energy to remove electron
Electronegativity: Ability to attract electrons

SPM Exam Tips

When studying periodic trends:

Remember the general patterns (radius decreases across period, increases down group)

Know the exceptions (e.g., Be > B in ionization energy)

Understand the reasons behind the trends (nuclear charge, electron shielding)

Practice identifying elements in different blocks

4.3 Elements in Group 18 (Noble Gases)

Characteristics of Noble Gases

Physical Properties

Element	Atomic Number	Electron Configuration	Boiling Point (°C)	Melting Point (°C)
Helium	2	$1s^2$	-268.9	-272.2
Neon	10	$[\text{He}]2s^22p^6$	-246.1	-248.6
Argon	18	$[\text{Ne}]3s^23p^6$	-185.7	-189.4
Krypton	36	$[\text{Ar}]4s^23d^{10}4p^6$	-153.4	-157.4
Xenon	54	$[\text{Kr}]5s^24d^{10}5p^6$	-108.1	-111.8
Radon	86	$[\text{Xe}]6s^24f^{14}5d^{10}6p^6$	-61.7	-71

Chemical Properties

Common Features:

Valence Shell: Completely filled ( $ns^2np^6$ )
Stability: High stability due to full valence shell
Reactivity: Very low (inert gases)
Ionization Energy: Highest in their respective periods

Chemical Behavior

Reactivity Order: He < Ne < Ar < Kr < Xe < Rn

Helium and Neon: Essentially inert
Argon: Forms some compounds under extreme conditions
Krypton: Forms few compounds (e.g., Kr $F_2$ )
Xenon: Forms several compounds (e.g., Xe $F_2$ , Xe $F_4$ , Xe $F_6$ )
Radon: Radioactive, forms compounds similar to xenon

Noble Gas Compounds

Compound	Formula	Formation Conditions	Properties
Xenon difluoride	Xe $F_2$	High temperature, UV light	Colorless crystalline solid
Xenon tetrafluoride	Xe $F_4$	High pressure, excess $F_2$	White crystalline solid
Xenon hexafluoride	Xe $F_6$	High pressure, excess $F_2$	Colorless crystals
Xenon trioxide	Xe $O_3$	Hydrolysis of Xe $F_6$	Yellow explosive solid

Occurrence and Uses

Element	Occurrence	Uses
Helium	Natural gas deposits	Balloons, cryogenics, welding
Neon	Atmospheric trace	Lighting signs, lasers
Argon	Atmospheric (0.93%)	Welding, lighting, inert atmosphere
Krypton	Atmospheric trace	High-intensity lamps, lasers
Xenon	Atmospheric trace	High-intensity lamps, medical imaging
Radon	Radioactive decay	Cancer treatment research

Importance of Noble Gases

Industrial Applications

Inert Atmospheres: Prevent unwanted reactions
Lighting: Neon signs, fluorescent lamps
Cryogenics: Liquid helium for superconductivity
Medical: Xenon for anesthesia, imaging

Scientific Research

Cryogenics: Superconductivity research
Plasma Physics: Noble gas plasmas
Chemical Bonding: Studying extreme compounds

Key Terms

Noble Gases: Group 18 elements with filled valence shells
Inert Gases: Elements with very low reactivity
Full Valence Shell: Complete outer electron configuration
Ionization Energy: Energy required to remove electron

Did You Know?

Helium is the only element that cannot be solidified at normal atmospheric pressure, no matter how low the temperature. It requires pressures greater than 25 atmospheres to become solid, making it unique among all elements.

4.4 Elements in Group 1 (Alkali Metals)

Characteristics of Alkali Metals

Physical Properties

Element	Atomic Number	Electron Configuration	Density (g/c $m^3$ )	Melting Point (°C)
Lithium	3	[He]2 $s^1$	0.534	180.5
Sodium	11	[Ne]3 $s^1$	0.968	97.8
Potassium	19	[Ar]4 $s^1$	0.856	63.5
Rubidium	37	[Kr]5 $s^1$	1.532	39.3
Cesium	55	[Xe]6 $s^1$	1.873	28.4
Francium	87	[Rn]7 $s^1$	1.870	~27

Chemical Properties

Common Features:

Valence Shell: n $s^1$ configuration
Reactivity: Extremely high (most reactive metals)
Ion Formation: Lose 1 electron to form M⁺ ions
Electronegativity: Lowest in their respective periods
Ionization Energy: Lowest in their respective periods

Reactivity Trend

Order of Reactivity: Li < Na < K < Rb < Cs < Fr

Down the group: Reactivity increases
Reason: Decreasing ionization energy, increasing atomic size

Chemical Reactions

Reaction with Oxygen $4\text{M} + \text{O}_2 \rightarrow 2\text{M}_2\text{O} \quad \text{(normal oxides)}$ $2\text{M} + \text{O}_2 \rightarrow \text{M}_2\text{O}_2 \quad \text{(peroxides - Na, K)}$ $\text{M} + \text{O}_2 \rightarrow \text{MO}_2 \quad \text{(superoxides - K, Rb, Cs)}$
Reaction with Water $2\text{M} + 2\text{H}_2\text{O} \rightarrow 2\text{MOH} + \text{H}_2$
- Lithium: Gentle reaction
- Sodium: Vigorous reaction
- Potassium: Very vigorous reaction
- Rubidium/Cesium: Explosive reaction
Reaction with Halogens $2\text{M} + \text{X}_2 \rightarrow 2\text{MX} \quad \text{(white ionic solids)}$

Occurrence and Uses

Element	Occurrence	Uses
Lithium	Minerals (spodumene)	Batteries, ceramics, medicine
Sodium	Salt deposits (NaCl)	Sodium vapor lamps, NaOH production
Potassium	Minerals (sylvite)	Fertilizers, soaps, glass
Rubidium	Trace minerals	Atomic clocks, electronics
Cesium	Trace minerals	Atomic clocks, photoelectric cells
Francium	Radioactive decay	Research only

Safety Precautions

Hazards

Fire Hazard: React violently with water
Storage: Stored under oil to prevent reaction with air/moisture
Handling: Use dry inert atmosphere, avoid contact with water

Emergency Procedures

Fire: Use Class D fire extinguishers (DO NOT USE WATER)
Spills: Cover with dry sand, evacuate area
Contact: Wash with copious water for skin contact

SPM Exam Tips

For Group 1 elements:

Know the reactivity trend and reasons (increasing atomic size, decreasing IE)

Understand different oxide formation (normal, peroxide, superoxide)

Memorize the reaction with water and observations

Remember storage requirements (under oil)

4.5 Elements in Group 17 (Halogens)

Characteristics of Halogens

Physical Properties

Element	Atomic Number	Electron Configuration	State at STP	Color
Fluorine	9	[He]2 $s^2$ 2 $p^5$	Gas	Pale yellow
Chlorine	17	[Ne]3 $s^2$ 3 $p^5$	Gas	Greenish-yellow
Bromine	35	[Ar]4 $s^2$ 3 $d^1$ ⁰4 $p^5$	Liquid	Reddish-brown
Iodine	53	[Kr]5 $s^2$ 4 $d^1$ ⁰5 $p^5$	Solid	Gray-black
Astatine	85	[Xe]6 $s^2$ 4 $f^1$ ⁴5 $d^1$ ⁰6 $p^5$	Solid	Black (radioactive)

Chemical Properties

Common Features:

Valence Shell: n $s^2$ n $p^5$ configuration
Reactivity: High nonmetals
Ion Formation: Gain 1 electron to form X⁻ ions
Electronegativity: Highest in their respective periods
Electron Affinity: High (tendency to gain electrons)

Reactivity Trend

Order of Reactivity: F > Cl > Br > I > At

Down the group: Reactivity decreases
Reason: Decreasing electron affinity, increasing atomic size

Chemical Reactions

Reaction with Metals $2\text{M} + 3\text{X}_2 \rightarrow 2\text{MX}_3 \quad \text{(Alkali metals)}$ $\text{M} + \text{X}_2 \rightarrow \text{MX} \quad \text{(Alkaline earth metals)}$
Reaction with Hydrogen $\text{H}_2 + \text{X}_2 \rightarrow 2\text{HX} \quad \text{(hydrogen halides)}$
- Reactivity: $F_2$ > C $l_2$ > B $r_2$ > $I_2$
Displacement Reactions $\text{X}_2 + 2\text{NaY} \rightarrow 2\text{NaX} + \text{Y}_2 \quad \text{(if X is more reactive than Y)}$

Occurrence and Uses

Element	Occurrence	Uses
Fluorine	Minerals (fluorite)	Teflon, toothpaste, uranium enrichment
Chlorine	Salt deposits (NaCl)	Bleach, disinfectant, PVC production
Bromine	Seawater, salt lakes	Flame retardants, pharmaceuticals
Iodine	Seaweed, caliche	Disinfectants, medicine, photography
Astatine	Trace amounts	Research only

Halogen Compounds

Hydrogen Halides

Compound	Formula	Properties	Uses
Hydrogen fluoride	HF	Colorless gas, acidic	Etching glass, refrigerant
Hydrogen chloride	HCl	Colorless gas, acidic	Cleaning, chemical synthesis
Hydrogen bromide	HBr	Colorless gas, acidic	Organic synthesis
Hydrogen iodide	HI	Colorless gas, acidic	Research

Important Applications

Disinfection: Chlorine and iodine for water treatment
Pharmaceuticals: Fluoride in toothpaste, iodine in antiseptics
Industrial: Bromine in flame retardants, chlorine in PVC
Research: Fluorine compounds in Teflon, pharmaceuticals

Key Terms

Halogens: Group 17 elements (F, Cl, Br, I, At)
Electron Affinity: Energy change when gaining electron
Oxidizing Agents: Substances that accept electrons
Displacement Reactions: More reactive element displaces less reactive one

Did You Know?

Fluorine is the most electronegative element, meaning it has the greatest ability to attract electrons in a chemical bond. This property makes it extremely reactive and why it's never found as a free element in nature.

4.6 Elements in Period 3

Overview of Period 3 Elements

Element	Symbol	Atomic Number	Group	Block	Electron Configuration
Sodium	Na	11	1	s-block	[Ne]3 $s^1$
Magnesium	Mg	12	2	s-block	[Ne]3 $s^2$
Aluminum	Al	13	13	p-block	[Ne]3 $s^2$ 3 $p^1$
Silicon	Si	14	14	p-block	[Ne]3 $s^2$ 3 $p^2$
Phosphorus	P	15	15	p-block	[Ne]3 $s^2$ 3 $p^3$
Sulfur	S	16	16	p-block	[Ne]3 $s^2$ 3 $p^4$
Chlorine	Cl	17	17	p-block	[Ne]3 $s^2$ 3 $p^5$
Argon	Ar	18	18	p-block	[Ne]3 $s^2$ 3 $p^6$

s-Block Elements (Na, Mg)

Sodium (Na)

Properties: Soft, silvery metal, highly reactive
Reactions: Violent with water, forms Na⁺ ions
Uses: Sodium vapor lamps, NaOH production
Compounds: NaCl (table salt), N $a_2$ C $O_3$ (soda ash)

Magnesium (Mg)

Properties: Light, strong metal, less reactive than Na
Reactions: Reacts with steam, forms M $g^2$ ⁺ ions
Uses: Alloys, flares, reducing agent
Compounds: MgO (refractory), Mg(OH)₂ (antacid)

p-Block Elements (Al to Ar)

Aluminum (Al)

Properties: Lightweight, good conductor, amphoteric
Reactions: Forms protective oxide layer
Uses: Aircraft, beverage cans, electrical wiring
Compounds: $Al_2O_3$ (alumina), $Al(OH)_3$

Silicon (Si)

Properties: Metalloid, semiconductor
Structure: Giant covalent structure
Uses: Computer chips, solar cells, glass
Compounds: $SiO_2$ (sand/quartz), $SiC$ (carborundum)

Phosphorus (P)

Allotropes: White (reactive), red (stable), black (semiconductor)
Reactions: White P ignites in air
Uses: Fertilizers, matches, detergents
Compounds: $P_4O_{10}$ , $H_3PO_4$ (phosphoric acid)

Sulfur (S)

Allotropes: $S_8$ (ring structure)
Properties: Yellow solid, brittle
Uses: Vulcanization of rubber, sulfuric acid
Compounds: $SO_2$ , $H_2SO_4$ , $H_2S$

Chlorine (Cl)

Properties: Greenish-yellow gas, strong oxidizing agent
Reactions: Forms $HCl$ with hydrogen
Uses: Disinfection, bleach, PVC production
Compounds: $NaCl$ , $HCl$ , $CaOCl_2$ (bleaching powder)

Argon (Ar)

Properties: Colorless, odorless, inert gas
Uses: Inert atmosphere, welding, lighting
Compounds: Essentially none (noble gas)

Periodic Trends in Period 3

Property	Na	Mg	Al	Si	P	S	Cl	Ar
Atomic radius	Largest	Large	Medium	Small	Small	Small	Smallest	Smallest
Ionization energy	Low	Medium	Medium	High	High	High	Highest	Highest
Electronegativity	Very low	Low	Medium	Medium	Medium	Medium	High	Very high
Metallic character	Strong	Strong	Strong	Weak	Nonmetallic	Nonmetallic	Nonmetallic	Nonmetallic
Oxide character	Basic	Basic	Amphoteric	Acidic	Acidic	Acidic	Acidic	None

Oxide Properties

Element	Oxide	Nature	pH of Solution
Sodium	$Na_2O$	Basic	High pH (alkaline)
Magnesium	$MgO$	Basic	High pH (alkaline)
Aluminum	$Al_2O_3$	Amphoteric	Neutral (can act as acid or base)
Silicon	$SiO_2$	Acidic	Low pH (neutral in water)
Phosphorus	$P_4O_{10}$	Acidic	Low pH (acidic)
Sulfur	$SO_3$	Acidic	Low pH (acidic)
Chlorine	$Cl_2O_7$	Acidic	Very low pH (strong acid)

SPM Exam Tips

For Period 3 elements:

Understand the transition from metallic to nonmetallic character

Memorize oxide properties and their pH

Know the physical states and key properties

Practice writing reactions for each element

4.7 Transition Elements

Characteristics of Transition Elements

Definition

Transition Elements: Elements in Groups 3-12 with incomplete d subshells in their ground state or common oxidation states.

General Properties

Property	Description	Example
Variable Oxidation States	Multiple stable oxidation states	Fe: +2, +3
Colored Compounds	Most compounds are colored	C $u^2$ ⁺: blue, F $e^2$ ⁺: green
Catalytic Activity	Many act as catalysts	Fe in Haber process
High Melting/Boiling Points	Strong metallic bonding	W: 3422°C
High Density	Close-packed structures	Os: 22.59 g/c $m^3$
Paramagnetism	Unpaired electrons attract magnetic field	Fe, Ni, Co

First Row Transition Elements

Element	Symbol	Atomic Number	Common Oxidation States	Color of Ion
Scandium	Sc	21	+3	Colorless
Titanium	Ti	22	+2, +3, +4	Purple (T $i^3$ ⁺), Yellow (T $i^4$ ⁺)
Vanadium	V	23	+2, +3, +4, +5	Violet ( $V^2$ ⁺), Green ( $V^3$ ⁺)
Chromium	Cr	24	+2, +3, +6	Blue (C $r^2$ ⁺), Green (C $r^3$ ⁺)
Manganese	Mn	25	+2, +3, +4, +6, +7	Pink (M $n^2$ ⁺), Purple (Mn $O_4$ ⁻)
Iron	Fe	26	+2, +3	Green (F $e^2$ ⁺), Yellow (F $e^3$ ⁺)
Cobalt	Co	27	+2, +3	Pink (C $o^2$ ⁺), Blue (C $o^3$ ⁺)
Nickel	Ni	28	+2, +3	Green (N $i^2$ ⁺)
Copper	Cu	29	+1, +2	Blue (C $u^2$ ⁺)
Zinc	Zn	30	+2	Colorless

Important Transition Elements

Iron (Fe)

Properties:

Most common transition metal
Excellent electrical and thermal conductivity
Forms magnetic compounds
Corrodes easily in moist air

Compounds:

$FeO$ : Iron(II) oxide - basic
$Fe_2O_3$ : Iron(III) oxide - basic
$FeCl_2$ : Iron(II) chloride - green
$FeCl_3$ : Iron(III) chloride - brown
$FeSO_4·7H_2O$ : Green vitriol - light green crystals

Uses:

Steel production (alloy with carbon)
Hemoglobin in blood
Magnets
Catalyst in Haber process

Copper (Cu)

Properties:

Excellent conductor of electricity and heat
Resistant to corrosion
Reddish-brown metal
Low reactivity

Compounds:

$CuO$ : Copper(II) oxide - black
$Cu_2O$ : Copper(I) oxide - red
$CuSO_4·5H_2O$ : Blue vitriol - blue crystals
$Cu(OH)_2$ : Copper(II) hydroxide - blue gelatinous precipitate

Uses:

Electrical wiring
Plumbing
Coinage
Catalysis
Antimicrobial applications

Zinc (Zn)

Properties:

Bluish-white metal
Low melting point
Galvanization protection for iron

Compounds:

$ZnO$ : Zinc oxide - white, used in creams and paints
$ZnS$ : Zinc sulfide - used in phosphors
$ZnCO_3$ : Smithsonite - zinc ore

Uses:

Galvanizing steel
Batteries (Zn-C, Zn-Ag)
Alloys (brass, bronze)
Die-casting

Catalytic Properties

Industrial Catalysts

Catalyst	Reaction	Uses
Iron ( $Fe$ )	$N_2 + 3H_2 \rightleftharpoons 2NH_3$	Haber process (ammonia)
Vanadium(V) oxide	$2SO_2 + O_2 \rightleftharpoons 2SO_3$	Contact process (sulfuric acid)
Platinum ( $Pt$ )	$4NH_3 + 5O_2 \rightarrow 4NO + 6H_2O$	Ostwald process (nitric acid)
Nickel ( $Ni$ )	C=C + $H_2 \rightarrow$ C-C	Hydrogenation of oils
Manganese(IV) oxide	$2KClO_3 \rightarrow 2KCl + 3O_2$	Laboratory oxygen preparation

Biological Catalysts

Many transition metals are essential in biological systems:

Iron: Hemoglobin (oxygen transport)
Copper: Enzymes involved in respiration
Zinc: Carbonic anhydrase (C $O_2$ transport)
Magnesium: Chlorophyll (photosynthesis)

Key Terms

Transition Elements: Elements with incomplete d subshells
Variable Oxidation States: Multiple possible charge states
Catalyst: Substance that increases reaction rate
Paramagnetism: Attraction to magnetic fields

SPM Exam Tips

For transition elements:

Remember the common oxidation states for each element

Know the characteristic colors of their compounds

Understand catalytic applications in industry

Practice identifying which elements are transition metals

Summary

This chapter has explored the periodic table and its organization:

Historical Development: From Dobereiner to Mendeleev to modern table
Modern Organization: Periods, groups, blocks, and their significance
Periodic Trends: Atomic radius, ionization energy, electronegativity
Group 18: Noble gases with filled valence shells
Group 1: Alkali metals with high reactivity
Group 17: Halogens with high electronegativity
Period 3: Transition from metallic to nonmetallic character
Transition Elements: Variable oxidation states and catalytic properties

Understanding the periodic table is essential for predicting chemical behavior and explaining why elements have specific properties.

Practice Tips for SPM Students

Create a periodic table trend summary chart

Practice identifying elements in different groups and periods

Work through periodic trend calculations

Memorize key properties of important groups

Review laboratory applications of transition metals