The Picture Worth a Thousand Words The Periodic Table The Picture Worth a Thousand Words

The Periodic Table Function: organizational tool which classifies the over 100 elements in a systematic way; allows for easy comparison of the elements and provides substantial information by illustrating patterns Imagine reading the content of the periodic table in dictionary format

Organization of the Periodic Table The periodic table visually classifies elements by multiple traits number of protons (and electrons) groups and periods metals, nonmetals, and metalloids main-group elements (representative elements), transition metals, inner transition metals s, p, d, f blocks valence electrons and electron configurations periodic trends

Key to the Periodic Table Every element is represented by a box containing the atomic number, the element symbol, the element name, and the atomic mass of the element

Overview of Periodic Table Elements are arranged by increasing atomic number in rows known as periods and columns known as groups (or families)

Metals, metalloids, and nonmetals

Metals, Nonmetals, and Metalloids Classification as metals, nonmetals or metalloids is based on specific properties Most important to this classification is ability to conduct electricity or heat Metals are good conductors Nonmetals are poor conductors Metalloids are intermediate in nature

Metals, nonmetals, and metalloids Property Metals Metalloids Nonmetals 1 good conductors of electric current and heat ability to conduct electric current varies with temperature poor conductors of electric current and heat 2 usually solid at room temperature (except Hg) often gases at room temperature; when solids typically brittle 3 most are malleable have properties of metals and nonmetals; when have low boiling points 4 most metals are ductile conductors they do not conduct heat or   5 reactivity varies; some highly reactive, others do not react easily electricity as well as the metals reactivity varies; some highly reactive, some are extremely unreactive 6 usually have luster (shiny) when solids typically dull-looking

Properties of Metals Good conductors of electricity and heat burner on stove is made of metal handles on cookware are not metal Usually solids at room temperature (often hard) exception is Hg (mercury) – liquid Most are malleable – can be formed into sheets/foils Most are ductile – can be drawn into wires Reactivity varies (most react with acids) some are highly reactive some do not react easily Usually shiny (luster) Au (gold) Ag (silver)

Properties of Nonmetals Poor conductors of electricity and heat Often gases at room temperature Have low boiling points reason they are gases at room temperature Reactivity varies some are highly reactive – halogens some are highly unreactive – noble gases If solids they are typically dull-looking

Properties of Metalloids Metalloids have properties of metals and nonmetals; their properties tend to fall between metals and nonmetals do not conduct heat or electricity as well as the metals properties vary widely Ability to conduct electricity varies with temperature semiconductors (used in electronics) used in computer chips and solar cells

Metals Nonmetals Metalloids Shiny dull solids varied Solid liquid/gas solid Be Ca Ba C P S Si Ge Bi Os

Main Group Elements and Others main group (representative elements): the s and p blocks transition metals: the d block inner transition metals: the f block

Main Group Elements most clearly illustrate periodic patterns most abundant naturally occurring elements most important elements for living things, including most common components of organisms: O, C, N, H, P, Ca

Transition Metals Metals (with the properties of metals) Can have multiple oxidation states d orbitals allow them to behave differently, and to be less predictable Form complexes with water

Inner Transition Metals Display very similar reactivity to each other Actinides are all radioactive unstable

s, p, d, and f Blocks

Special Groups Group = column Group members typically have similar properties Four groups have special names: Alkali metals highly reactive Alkaline earth metals reactive Halogens highly reactive Noble gases highly unreactive

Special Groups A A L L K K N O I I B N M L E E E T G A R H S S T A H L

Versions of the Periodic Table Updated with new elements Updated with more accurate atomic masses Presented with labels and color-coding Different arrangements for different purposes

Versions of the Periodic Table Different versions emphasize different features They are all correct, and none is absolute

Origins of the Periodic Table Mendeleev: created one of the earliest versions of the periodic table Arranged elements by increasing mass (right or wrong?) Grouped elements with similar properties

Mendeleev’s 1872 Periodic Table Predicted undiscovered elements and their properties

Mendeleev’s Contribution Mendeleev developed Periodic Law (Modern) Periodic Law: the elements, when listed in order of their atomic numbers, fall into recurring groups, so that elements with similar properties occur at regular intervals Periodic law is observed by all members of a column having the same ending to their electron configuration

The Big Picture Periodic Table shows patterns Multiple types of patterns can be found Periodic repetition of chemical and physical properties of the elements as you move across a period (row) Or down a group (column) This trait of periodic repetition of the properties of elements when they are arranged by increasing atomic number is known as periodic law.

Electron Configurations Pattern across a period and pattern down a group

Periodic Trends Certain properties show a systematic variation across a period and down a group: metallic character valence electrons atomic radii (atomic size) ionic radii electronegativity ionization energy

Metallic Character

Valence Electrons Alkali metals Halogens

Valence Electrons Equal number of valence electrons for all members of a group For main-group elements (IA-VIIIA) the valence electrons are equal to the group number Exception: He (He only has 2 e-)

Atomic Radius As protons are added across a period, the nucleus has a larger positive charge Na +11 Al +13 S +16 Ar +18 As electrons are added across a period, they continue to fill the same energy level, n The result? Larger effective nuclear charge on the electrons from left to right across a period

Atomic Radius atomic radius increases atomic radius increases

electronic structure of ion** Ionic Radii   Na+ Mg2+ Al3+ P3- S2- Cl- no of protons 11 12 13 15 16 17 electronic structure of ion** 2,8 2,8,8 ionic radius (nm) 0.102 0.072 0.054 0.212 0.184 0.181 **How many electrons does each have? Na+ has 10, Mg2+ has 10, Al3+ has ? Cl- has 18, S2- has 18, P3- has ?

Ionic Radius The positive ions (Cations) Each ion has exactly the same electronic structure, and is equal to the structure of Ne, with n = 2. However, the number of protons for each ion is increasing. That causes a larger effective nuclear charge that pulls the electrons towards the nucleus and causes the ionic radii to decrease.

Ionic Radius The negative ions (Anions) Each ion has the same electronic structure again, but this time equal to the structure of Ar with n = 3. Compared to the cations, there is an extra layer of electrons, so the ionic radii increase.

Ionic Radii Cations become smaller than their atoms Anions become larger than their atoms

Electronegativity Electronegativity: a measure of the ability of an atom in a compound to attract electrons More electronegative elements are more able to pull electrons from another atom towards their own nucleus

Electronegativity Fluorine is the MOST electronegative element

Electronegativity Think: Are Noble Gases electronegative? What is their electron configuration? How does this impact tendency to gain or lose electrons? Noble gases DO NOT follow the trend for electronegativity

Ionization Energy Ionization energy: the energy required to remove an electron from an atom in the gaseous state Remember, atomic radius decreases across a period due to greater effective nuclear charge Electrons are closer to the nucleus and harder to remove From top to bottom within a group, the valence electrons are in a higher energy, farther shell This results in electron shielding, making it easier to remove electrons

Ionization Energy

First Ionization Energies

Ionization Energy Atoms can lose more than one electron Always requires more energy to remove a second electron (increased effective nuclear charge) Peaks change element WHY?

Subsequent Ionization Energies Each electron removed changes the peak

Electron Affinity The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion. F(g) + e- F-(g) Ho = -328.0 kJ/mol The negative sign of Ho indicates that energy is given off. This process is exothermic.

Electron Affinity Electron affinity is more or less the opposite of ionization energy Halogens: greatest electron affinity Noble gases: lowest electron affinity (requires E)

Learning Check How do electron affinity and ionization energy compare? F electron affinity F ionization energy

Learning Check How do electron affinity and ionization energy compare? F + e- F- electron affinity F F+ + e- ionization energy Electron affinity is the energy required to gain an electron. Ionization energy is the energy required to lose an electron.