Chapter 11 & 12

Chapter 11 & 12

2

John Newlands

• In 1864, noticed when the

elements were arranged in

order of

increasing

atomic mass, their properties repeated every eight elements.

– THE LAW OF OCTAVES

4

Meyer & Mendeleev

• In 1869, published almost identical versions with the elements in order of increasing atomic mass and in columns with similar properties.

5

Mendeleev

• Mendeleev is given more credit than Meyer BECAUSE:

– He published his table first

– He better demonstrated his table

•Suggested some of the previously measured masses were incorrect

•Left blanks for not yet discovered elements

6

Mosley

• In 1913, using X-rays, he

discovered a unique number of protons in the nuclei of atoms for each element.

• Today the elements are arranged in order of increasing atomic

number

• There is a periodic repetition of chemical and physical properties of the elements when they are

arranged in order of increasing atomic number

Periodic Law

Arrangement of the Periodic Table

• Groups/Families

– 18 vertical columns (↑↓) – Two Labeling Systems

1. Number-and-letter system

- 1A through 8A columns (representative elements)

- 1B through 8B short columns (transition elements)

2. Number system, Group 1 to Group18

• Periods

– 7 horizontal rows (↔)

8

PERIODS

GROUPS/FAMILIES

Arrangements of the Periodic Table

10

Metals

Non-Metal s

Metals

Metalloi d

Non-Metal

s

Create a

Legend

Metals

• Shiny

• Solid at room temperature

• Good conductors of heat and electricity

• Malleable

• Ductile ^{Group 1}

Alkali Metals Group 2

Alkaline Earth Metals Groups 3-12

Transition Metals

Lanthanide & Actinide Groups

Inner Transition Metals or Rare Earth Metals

12

Valence Electrons

Group # Group Name # Valence e^- 1 (1A) Alkali Metals 1

2 (2A) Alkaline Earth Metals 2

3-12 Transition Metals varies 13 (3A) Boron Group 3

14 (4A) Carbon Group 4

15 (5A) Nitrogen Group 5 16 (6A) Oxygen Group 6

17 (7A) Halogens 7

18 (8A) Noble Gases 8

13

Nonmetals & Metalloids

• Nonmetals

– Dull

– Generally gases or brittle solids at room temperature

– Poor conductors of heat and electricity

• Metalloids

– Elements with physical and chemical properties of both metals and

nonmetals

– Rest on the “stair-step”

B

Si

As

Te

At Ge

Sb

Po

←Metals

Nonmetals →

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Element Placement

Why are elements put into groups/families together?

Because they have similar chemical properties

Why do elements have similar chemical properties?

Because they have the same number of valence electrons

Group 1 – Alkali Metals Period 2 Lithium 1s²2s¹ [He]2s¹ Period 3 Sodium 1s²2s²2p⁶3s¹ [Ne]3s¹

Period 4 Potassium 1s²2s²2p⁶3s²3p⁶4s¹ [Ar]4s¹

ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON

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Dot Diagrams for Representative Elements

Representative Elements

• s-block elements

– Groups 1&2, hydrogen & helium

– Two columns Maximum 2 electrons fills the 1 s orbital.

– Valence electrons occupy outermost s sublevels only.

• p-block elements

– Groups 13-18 (except helium)

– Six columns Maximum 6 electrons fills the 3 p orbitals.

– Valence electrons include a full outermost s

sublevel and a filled or partially filled p sublevel.

Period number is equal to the principle energy level where the valence electrons are located

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Transition Elements

• d-block elements –Groups 3-12

–Ten columns Maximum 10

electrons fills the 5 d orbitals.

–Valence electrons include a full outermost s sublevel and a filled or partially filled d

sublevel.

Inner Transition metals or Rare Earth Metals

• f-block elements

– Fourteen columns Maximum 14 electrons fills the 7 f orbitals.

– Lanthanide & Actinide Groups – Full or partially full outermost s

sublevel, and full or partially full

outermost f sublevel.

20

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Periodic Trends

• Atomic Radii:

• Ionic Radii:

• Ionization Energy:

• Electron Affinity:

• Electronegativity:

attract electrons toward itself from a covalent chemical bond.

Atomic Radius

• Half the distance between two nuclei of identical atoms that are chemically

bonded together

• Down the group

– atomic radius increases, because…

• Across the period

– atomic radius decreases, becauses….

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Practice Atomic Radius

• Which has the larger atomic radii of the following?

B or Al Na or Mg F or Cl

• Which has the smaller atomic radii of the following?

H or He K or Cs N or Ne

• Circle the one with the largest atomic radius and underline the one with the smallest.

C, Si, Ge V, Cr, W N, Mg, Ca

Atomic Radius Decreases

Atomic Radius Increases

B Al Na Mg

F Cl

H He

K

Cs

N Ne

C Si Ge V Cr

W

N Mg

Ca

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Ionization Energy

• The amount of energy required to remove an electron from the atom

(how tightly an atom holds on to its electrons)

• A general term for the energy required to remove an electron from an orbital in an

atom. Think of it also as the energy required to make a cation.

• Down a group

– ionization energy decreases, because…

• Across a period

– ionization energy increases, because…

Practice Ionization Energy

• Which has the greater ionization energy?

Ne or Ar N or O Sc or Ti

• Which has the smaller ionization energy?

Al, Si, P K, Rb, Sr Be, Mg, Ca

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Ionization Energy Increases

Ionization Energy Decreases

Ne Ar N O

Sc Ti

Al Si P K

Rb Sr Be Mg Ca

Electron Affinity

• The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively

charged ion.

• The attraction to additional electrons

• A fluorine atom in the gas phase, for example,

gives off energy when it gains an electron to form a fluoride ion.

• Down the group

– electron affinity decreases, because

• Across the period

– electron affinity increases, because

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Practice Electron Affinity

• Which has the smaller electron affinity?

Te or Xe Ag or Au Zn or Br

• Which has the greater electron affinity?

F, Cl, Br B, C, Si Ni, Cu, Au

Electron Affinity Increases

Electron Affinity Decreases

Te Xe Ag

Au

Zn Br

F Cl Br B C

Si Ni Cu

Au

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Ionic Radius

• Octet Rule

– Atoms tend to gain, lose, or share electrons in order to achieve a full outer energy level (typically 8 are needed) – Groups 1 A – 3 A, loses Valence e^-( 1 to 3 e^-).

– Group 4 A, share Valence e^-.

– Groups 5 A – 7 A, gain electrons ( 3 to 1 e^-).

• Noble Gases, Outer most shell are full. These elements don’t gain nor lose e^-, Non-reactive.

• Ion

– An atom that has an overall charge due to the gaining or losing of electrons

– Cation, positive charges (Groups 1A – 4A) – Anion, negative charges (Groups 5A – 7A)

Ionic Radius Comparisons

• Metals have LOW ionization energy and electron affinity

– They lose electrons to form positively charged ions

– Positive charged ions are smaller than the original atom

• Nonmetals have HIGH ionization energy and electron affinity

– They gain electrons to form negatively charged ions

– Negatively charged ions are larger than the original atom

Atom size decreases

Atom size increases

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Ionic Radius Practice

• Which is the smaller of the two?

Lithium ion or Lithium atom Chlorine ion or Chlorine atom

• Underline the following that will form a positively charged ion and circle the ones that will form a negatively charged ion.

Mg F Al Cu

Br N S K

• How will the radius of each of the above change when an ion is formed?

Mg F Al Cu

Br N S K

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Li

Mg Cl

Br F N

Al S

K Cu

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Electronegativity

• The attraction an atom has for electrons in a covalent bond.

• The ability of an atom to attract electrons in a chemical bond.

• Down the group

– Electronegativity values decrease, because

• Across the period

– Electronegativity values increase, because

*

Noble gases are the exception to

this rule.

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You must subtract the values of

electronegativity to determine it the

bond is covalent, polar covalent or ionic

Electronegativity of 0.0 to 0.3 is a non-polar covalent bond.

Electronegativity of 0.4 to 1.9 is a polar

covalent bond.

Electronegativity of 2.0 to 4.0 is an ionic bond.

Electronegativity

Covalent shares the electrons equally

Polar is slightly

negative on one side

Ionic has electrons captured by one atom.

Electronegativity Practice

• Which has the greater electronegativity value?

B or N Si or Sn Cr or W

• Which has the smaller electronegativity value?

Rb, Sr, Y Ge, In, Sn As, Se, S

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Electronegativity Increases

Electronegativity Decreases

B N

Si Si

Sn Cr

W Rb Sr Y

Ge In

As Se S

How can you remember all of these trends?

• Remember the cartoon characters Mighty Mouse

and Foghorn Leghorn?

• Put Mighty Mouse in the top right hand corner of the periodic table and Foghorn Leghorn in the bottom left; now draw the trends and it should make sense. Mighty Mouse, aka fluorine, is strong

(electronegative and electron affinity) and

small (atomic radii), while Foghorn is bigger (larger radius) and slower (less electron

affinity), etc.

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CUMULATIVE REVIEW, Due at the End of Class

1) Between these elements Ga, In, & Tl:

– Which has the highest ionization energy?

– Which has the smallest atomic radius?

2) Which is the smallest radius: (a) an atom of sodium, (b) an ion of sodium, or © an atom of potassium?

3) Between these elements (a) zinc, (b) arsenic, or (b) bromine:

- Which has the greatest electron affinity?

- Which has the lowest ionization energy?

Name: ______________ Date: ________Block: ________