periodic table chapter 5 history organization introduction to bonding trends

Periodic TableChapter 5

HistoryOrganization

Introduction to BondingTrends

Terms• The Periodic LawStates that the physical & chemical

properties of the elements are periodic functions of their atomic numbers.

• The Periodic TableAn arrangement of the elements in order of

their atomic numbers so that elements with similar properties fall in the same columns.

• The Modern Periodic TableIs organized by electron configuration.

Periodic Table & electron configurations

Blocks of the Periodic Table

sp

d

f

Video Clip: Periodic Table by cassiopeiaproject

• https://youtu.be/5MMWpeJ5dn4

Extended Periodic Table

History

1817-Dobereiner’s TriadsGroups of 3 elements with similar properties, the atomic mass of second was an average of all three.

EX: Ca = 40Sr = 88Ba = 137

History

1863- Newland’s Law of Octaves

If elements are arranged by increasing atomic mass, their properties repeat every 8 elements.

Mendeleev and Periodicity• The first periodic table of the

elements was published by Russian chemist Dmitri Mendeleev.

• Mendeleev left empty spaces in his table and predicted elements that would fill3 of the spaces.

• By 1886, all 3 of these elements had been discovered.

Chapter 5 – Section 1: History of the Periodic Table

History1869- Mendeleev’s First Periodic

Table

He arranged all the known elements in rows by increasing mass AND he grouped elements with similar properties in columns.

Brilliant in spite of discrepancies: Co & Ni, Te & I

Brilliant because of his predictions: ekasilicon = Ge

Video Clip: Periodic Table of Elements - Chemistry: A Volatile

History - BBC 

• https://youtu.be/nsbXp64YPRQ

Mosley and the Periodic Law

• In 1911, the English scientist Henry Moseley discovered that the elements fit into patternsbetter when they were arranged according to atomic number, rather than atomic weight.

• The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

Chapter 5 – Section 1: History of the Periodic Table

History1913- Moseley discovers protons

Nuclear charge (protons) increases by one for each element!

Atomic number, not atomic mass, is the basis for organization!

Organization Groups, Periods and Blocks

• Elements in the periodic table are arranged into vertical columns, called groups or families, that share similar chemical properties.

• Elements arealso organizedhorizontally in rows, or periods.

GROUPS• 18 vertical columns (families) with

family names• same outer e- configuration.• similar chemical properties.• labeling columns has changed over

time…

Labeling Groups… then & now

PERIODS

• 7 horizontal rows• corresponds to main Energy level

that is filled• length of row determined by

sublevels that are filling

Video Clip: The (truly) Periodic Table by scienceoffice.org

• https://youtu.be/xd4-Uy2FLWc

BLOCKS

The representative elements of the Main Groups: s & p blocks

only

Group 1: Alkali Metals• Group 1 elements are called alkali metals.• Alkali metals have a silvery appearance

and are soft enough to cut with a knife.• They are extremely reactive and are not

found in nature as free elements.• They must be stored under oil or kerosene.

Group 2: Alkaline Earth Metals• Elements in group 2 are known

as the alkaline earth metals.• Group 2 metals are harder, denser

and stronger than alkali metals, and have higher melting points.

• Less reactive than group 1, but still too reactive to be found in nature as free elements.

Group 17: Halogens• Elements in group 17 are

known as the halogens.• Halogens are the most

reactive nonmetals, reacting vigorously with metals to form salts

• Most halogens exist in nature as diatomic molecules (i.e. F2, Cl2, Br2 and I2.)

Group 18: Noble Gases• Elements in group 18 are

known as noble gases.• They are completely non-

reactive and don’t form compounds under normalconditions.

• A new group was added to the periodic table in 1898 for the noble gases.

d-block: Transition Metals• Elements in the d-block are

called transition metals.• They have typical metallic

properties such as conduction of electricity and high luster.

• Less reactive than group 1 and 2 elements.• Some (i.e. platinum & gold) are so unreactive

they usually don’t form compounds.

f-block: Lanthanides & Actinides• Elements in period 6

of the f-block are called lanthanides (or rare-earth).

• Lanthanides are shiny metals similar in reactivity to alkaline earth metals.

• Elements in period 7 of the f-block are called actinides.

• Actinides are all radioactive, and many of them are known only as man-made elements.

Classification & Location of the elements

Metals: left of staircaseNonmetals: right of staircaseSemimetals: touch the staircase (except

Al)

Why are H & He sometimes separated from Periodic Table?

• Hydrogen has one valence electron so it is normally placed in group 1.

• Hydrogen has a second home at the top of group 17 because it can gain an electron as the halogens do.

• Helium has two valence electrons like group 2 but is clearly a noble gas because its outer energy level is full.

Introduction to BondingPHYSICAL PROPERTIES OF METALS

luster, conductivity, solid, malleable, ductile

3 or fewer valence e- = defines a metal

Introduction to BondingCHEMICAL PROPERTIES OF

METALS….HOW THEY FORM BONDS

Lose valence e- to nonmetals → Ionic compounds

Bond with each other → Metallic bonds

Mix with other metals → Alloys

Introduction to BondingPHYSICAL PROPERTIES OF NONMETALS

No luster, poor conductors, not malleable or ductile

5 -8 valence e- = defines a nonmetalCan be soft solids: C, P, S, Se, I2Can be liquids: Br2

Can be gases: N2, Cl2, H2, O2, F2, & the noble gases

All gases are nonmetals but not all nonmetals are gases!

Introduction to BondingCHEMICAL PROPERTIES OF NONMETALS …. HOW

THEY FORM BONDSGain valence e- from metals → Ionic

compounds (EX: salt: NaCl)

Share valence with themselves→ Diatomic Molecules (EX: N2, H2 )

Share valence with other nonmetals → Molecules (EX: H2O, CO2)

• Atomic radius – one-half the distance between the nuclei of identical atoms

that are bonded together.

Atomic RadiiChapter 5 – Section 3: Electron Configuration and Periodic Properties

Group 1

• Atomic radii tend to increase as you go down a group because electrons occupy successively higher energy levels farther away from the nucleus.

• Atomic radii tend to decrease as you go across a period because as more electrons are added they are pulled closer to the more highly charged nucleus with more protons.

Atomic Radii (continued)

Chapter 5 – Section 3: Electron Configuration and Periodic Properties

Period 2

Atomic Radius (Size)

Of the elements Mg, Cl, Na, and P, which has the largest atomic radius? Explain.Solution:Na has the largest radius.All of the elements are in the 3rd period, and atomic radii decrease across a period.

Atomic RadiiSample Problem

Chapter 5 – Section 3: Electron Configuration and Periodic Properties

• An ion is an atom of group of bonded atoms that has a positive or negative charge.

• The energy required to remove an electron froma neutral atom of an element is called the ionization energy (IE).

• Ionization energy tends to increase across each period because a higher nuclear charge more strongly attracts electrons in the same energy level.

Ionization EnergyChapter 5 – Section 3: Electron Configuration and Periodic Properties

• Ionization energy tends to decrease down each group because electrons farther from the nucleus are removed more easily.

Ionization Energy (continued)

Chapter 5 – Section 3: Electron Configuration and Periodic Properties

Consider two elements, A and B. A has an IE of 419 kJ/mol. B has an IE of 1000 kJ/mol. Which element is more likely to be in the s block? Which will be in the p block? Which is more likely to form a positive ion?Solution:Element A is most likely to be in the s-block since IE increases across the periods.

Element B would most likely lie at the end of a period in the p block.

Element A is more likely to form a positive ion since it has a much lower IE than B.

Ionization EnergySample Problem

Chapter 5 – Section 3: Electron Configuration and Periodic Properties

Ionization Energy

• Electron affinity is the energy change that occurs when an electron is acquired by a neutral atom.

• Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.

• Electronegativity applies to atoms in a compound, while electron affinity is a property of isolated atoms.

Electron Affinity and ElectronegativityChapter 5 – Section 3: Electron Configuration and Periodic Properties

• Electron affinity and electronegativity both tend to increase across periods, and decrease (or stay the same) down a group.

Electron Affinity and Electronegativity (continued)

Chapter 5 – Section 3: Electron Configuration and Periodic Properties

Electron Affinity

ELECTRONEGATIVITY

Of the elements Ga, Br, and Ca, which has the highest electronegativity? Explain .

Solution:

All of these elements are in the fourth period.

Br has the highest atomic number and is farthest to the right in the period.

Br would have the highest electronegativity since electronegativity increases across a period.

ElectronegativitySample Problem

Chapter 5 – Section 3: Electron Configuration and Periodic Properties

Periodic Trends

• Atomic Radius, and Metallic Character

• Factors affecting size:Distance of outer e- from

nucleus, shielding effect of inner e- and size of the e- cloud with each sublevel makes atoms larger down a group.

Nuclear charge across the period pulls e- cloud in and makes the atoms smaller.

Periodic Trends

• Ionic Radius

• Factors affecting size:

Metals lose valence e- the ion is smaller = previous noble gas!

Nonmetals gain valence e- ion is larger = next noble gas!

Periodic Trends

First Ionization Energy:E req’d to remove outer e-

Electron Affinity:Atom’s attraction for e-

Metals low IE low e- affinity

Nonmetals high IEhigh e- affinity

Periodic Trends-Bonded Atoms

Electronegativity:The ability of an atom to attract e- in a chemical

bond.

Fluorine is assigned highest value = 4.0

Nonmetals high EN high e- affinity

Metals low EN low e- affinity

START CH.7valence electrons

• Occupy highest principal energy level• Responsible for chemical properties =

bonding • Elements in a group have similar

properties b/c valence e- same.• Outer E level has e- in s & p sublevels

only• Outer E level full when s & p

sublevels are full = noble gas configuration

= octet

Ion formation in atoms

Ion: an atom that has a charge b/c it has lost or

gained e-

The Octet (Duet) Rule: atoms will gain, lose or share e- in order to

acquire a full set of valence e- = 8 The Noble Gas Rule: atoms attain the nearest noble gas

configuration when they become ions.

Types of Ions

CATIONS are positive ions ANIONS are negative ions

Metals form monatomic cations To name: use name of element

Na+1 Mg+2 Al+3 Si+4 (sodium ion,

magnesium ion…)

Types of Ions

Nonmetals form monatomic anions To name: use name of element + ide

N-3 O-2 F-1 (nitride ion, oxide ion, fluoride

ion)

P-3 S-2 Cl-1 (phosphide, sulfide, chloride)

Transition Metals: d block

• 1 or 2 valence e-

• Several of these metals form more than one type of ion. (MULTIPLE OXIDATION STATES)

• How? Will lose the valence s e- then the d e- one at a time.

• To name, use roman numerals

Cu+1 copper I ion Fe+2 iron II ion

Cu+2 copper II ion Fe+3 iron III ion

Lower left of p block = metals

• 3 to 5 valence e-

• Several of these metals form more than one type of ion.

• How? Will lose all the valence s & p electrons OR just the p electrons.• To name, use roman numerals

Tl+1 thallium I ion Sn+2

tin II ionTl+3 thallium III ion Sn+4

tin IV ion

Types of Ions

Polyatomic ions: groups of covalently bonded atoms that carry a charge.

NH4+1 ammonium ion

SO4-2 sulfate ion

OH-1 hydroxide ionCN-1 cyanide ionPO4

-3 phosphate ionCO3

-2 carbonate ionNO3

-1 nitrate ion