Download - Unit 3.2 Part I - Introduction to the Periodic Table Unit 3.2 Part I - Introduction to the Periodic Table

Unit 3.2Part I - Introduction to the

Periodic Table

Unit 3.2Part I - Introduction to the

Periodic Table

•History of the Periodic Table

• Until 1800’s – no clear system

• Elements grouped by similar properties or atomic mass

The Search for a Periodic Table

• In 1829, J.W. Döbereiner classified some elements into groups of three, which he called triads.

Döbereiner’s Triads

• The elements in a triad had similar chemical properties, and their physical properties varied in an orderly way according to their atomic masses.

Element

Atomic mass (g)

Density (g/mL)

Melting point (C)

Boiling point (C)

Chlorine 35.5 0.00321 -101 -34

Bromine 79.9 3.12 -7 59

Iodine 127 4.93 114 185

Döbereiner’s Triads

• The concept of triads suggested that the properties of an element are related to its atomic mass.

• Density increases with increasing atomic mass.

Element

Atomic mass (g)

Density (g/mL)

Melting point (C)

Boiling point (C)

Chlorine 35.5 0.00321 -101 -34

Bromine 79.9 3.12 -7 59

Iodine 127 4.93 114 185

Which of the Dobereiner triads shown are still listed in the same column of the modern periodic table?Triad 1 Triad 2 Triad 3Li Mn SNa Cr SeK Fe Te

•Triad 1 and triad 3

• The Russian chemist, Dmitri Mendeleev, developed a periodic table of elements.

Mendeleev’s Periodic Table

•organized the elements according to increasing atomic mass.

• Mendeleev later developed an improved version of his table with the elements arranged in horizontal rows.


• Patterns of changing properties repeated for the elements across the horizontal rows.

•Elements in vertical columns have similar properties.



• properties of the elements repeat in an orderly way from row to row of the table.

• This repeated pattern is an example of periodicity in the properties of elements.

•Periodicity is the tendency to recur at regular intervals.


***Mendeleev correctly predicted the properties of several undiscovered elements.

Why is this important?

• In order to group elements with similar properties in the same columns, Mendeleev had to leave some blank spaces in his table.

• He suggested that these spaces represented undiscovered elements.

What are two factors that contributed to the acceptance of Mendeleev’s periodic law?

• Grouping of elements with similar chemical properties

• Ability to predict properties of undiscovered elements

The Modern Periodic Table

• The atomic number of an element is equal to the number of protons in the nucleus.

• Each row (except the first) begins with a metal and ends with a noble gas.

• the basis for ordering the elements in the table is the atomic number, not atomic mass.


• In between, the properties of the elements change in an orderly progression from left to right.

• This regular cycle illustrates periodicity in the properties of the elements.


• periodic law - physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number

Use the periodic table to separate these 12 elements into 6 pairs fo elements having similair properties.

Ca, K, Ga, P, Si, Rb, B, Sr, Sn, Cl, Bi, Br

Ca K Ga P Si Cl

Sr Rb B Bi Sn Br

•Layout of the Periodic Table

Layout of the periodic table

•A group, also called a family, consists of the elements in a vertical column.

Groups are numbered 1 – 18 ORIA – VIIIA for main group elements andIB – VIIIB for transition elements

1 valence e-

2 valence e-

As you move left to right across a period the number of valence electrons increases by one

3 valence e-

4 valence e-

1 valence e-

2 valence e-

Elements in the same group have same number of valence electrons and similar properties

3 valence e-

4 valence e-

•A period consists of the elements in a horizontal row

Periods are numbered 1-7 and each new row begins a new energy level

•The elements in the middle are called transition elements

The others are main group elements

• Lithium is the first element in Group 1 and in Period 2. Check this location on the periodic table.

•4 groups have commonly used names: alkali metals in Group 1 (IA)

• the alkaline earth metals in Group 2 (IIA)

• the halogens in Group 17 (VIIA) -from the Greek words for “salt former” , compounds that halogens form with metals are salt-like.

the noble gases in Group 18 (VIIIA) – full outer shell (8 valence electrons), generally unreactive

• In the periodic table, two series of elements are placed below the main body of the table.

The elements in these two series are known as the inner transition elements.

• The first series of inner transition elements is called the lanthanides because they follow element number 57, lanthanum.

• Because of their natural abundance on Earth is less than 0.01 percent, the lanthanides are sometimes called the rare earth elements.

• The second series of inner transition elements are the actinides

• All of the actinides are radioactive, and all beyond uranium (92) are man made (synthetic).

• Elemental Funkiness – By Mark Rosengarten UF

• http://www.youtube.com/watch?v=1PSzSTilu_s

•Classification of elements

• The majority of the elements are metals (solids). They occupy the entire left side and center of the periodic table.

• Elements are classified as metals, metalloids, or nonmetals on the basis of their physical and chemical properties.

• Nonmetals occupy the upper-right-hand corner. – green, yellow, orange

Physical States and Classes of the Elements

• Metalloids are located along the staircase boundary between metals and nonmetals. - purple

Physical States and Classes of the Elements

Metals

• Metals are elements that have luster, conduct heat and electricity, and usually bend without breaking.

Click box to view movie clip.

• All metals except mercury are solids at room temperature; in fact, most have extremely high melting points.

Metals

• The periodic table shows that most of the metals (coded blue) are not main group elements.

• With the exception of tin, lead, and bismuth, metals have one, two, or three valence electrons.

Nonmetals

• Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid.

• Many are gases at room temperature

• Their melting points tend to be lower than those of metals.

• With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons.

Properties of Metals and Nonmetals

Metalloids

• Metalloids have some chemical and physical properties of metals and other properties of nonmetals. - purple

• In the periodic table, the metalloids lie along the border between metals and nonmetals.

some metalloids are semiconductors

• Silicon’s semiconducting properties made the computer revolution possible.

• A semiconductor is an element that does not conduct electricity as well as a metal, but does conduct slightly better than a nonmetal.

• Some metalloids such as silicon, germanium (Ge), and arsenic (As) are semiconductors.

Part II - Periodic Trends

• Understanding the relationship between electron configuration and position in the periodic table enables you to predict the properties of the elements and the outcome of many chemical reactions.

Periodic Properties of the Elements

•The electron structure of an atom determines many of its chemical and physical properties.

•Atomic Radius–size of atom

© 1998 LOGAL

Atomic Size • size of an atom INCREASES in any

group as you go DOWN the column because the valence electrons are in energy levels farther from the nucleus.

• “shielding effect” – electrons in energy levels closer to the nucleus “shield” the valance electrons from the positive pull of the nucleus

The shielding effect• Increases down a group because

electrons are being added to higher energy levels

• There is no shielding effect as you go across a period because electrons are being added to the same principal energy level

• size of an atom DECREASES in any period as you go to the RIGHT in any row because there is an increased nuclear (+) charge pulling e- in tighter.

•Why larger going down?

–Higher energy levels have larger orbitals

–Shielding - core e- block the attraction between the nucleus and the valence e-

•Why smaller to the right?

–Increased nuclear charge without additional shielding pulls e- in tighter

Atomic Radius

Atomic Radii of Main Group Elements

• Which atom has the larger radius?

Be or Ba

Ca or Br

Ba

Ca

Examples

For each of the following pairs, predict which atom is larger.

a. Mg, Sr

b. Sr, Sn

c. Ge, Sn

d. Ge, Br

Sr

Sr

Sn

Ge

e. Cr, W W

•Octet Rule•reactivity of atoms is based on achieving a complete octet of valence electrons(8/8)

• Everybody wants to be like a noble gas!

Ne

Atoms achieve noble gas configuration by gaining or losing their valence electrons

An ion is an atom or group of atoms that has a charge because of the loss or gain of electrons.

cation - An ion that has LOST an e- and now has a positive (+) charge

anion – an ion that has GAINED an e- and now has a negative (-) charge

Common Ion Charges

aka oxidation number

1+

2+ 3+ 3- 2- 1-

0

•GROUP #: VALENCE # WHEN FORMING IONS:

• OUT OF 8:

•Group IA 1 loses 1

•Group IIA 2 loses 2

•Group IIIA 3 loses 3

•Group IVA 4 can lose or gain

•Group VA 5 gains 3

•Group VIA 6 gains 2

•Group VIIA 7 gains 1

•Group VIIIA 8 does not form ions

Ionic Size• positive ions + (cations) acquire the

configuration of the noble gas in the preceding period.

• the outermost electrons of the ion are in a lower energy level than the valence electrons of the neutral atom.

• The electrons that are not lost by the atom experience a greater attraction to the nucleus and pull together in a tighter bundle with a smaller radius.

• all cations ions have smaller radii than their corresponding atoms.

• anions acquire the electron configuration of the noble gas at the end of its period.

• But the nuclear charge doesn’t increase with the number of electrons.

• In the case of fluorine, a nuclear charge of 9+ must hold ten electrons in the F– ion; all the electrons are held less tightly

• the radius of the anion is larger than the neutral atom.

• Ionic Radius

Cations (+)

lose e-

smaller

© 2002 Prentice-Hall, Inc.

Anions (–)

gain e-

larger

Ionic Radius

• Which particle has the larger radius?

S or S2-

Al or Al3+

S2-

Al

Examples

a. Mg, Mg2+

b. S, S2–

c. Ca2+, Ba2+

d. Cl–, I–

Mg

S2–

Ba2+

I–

e. Na+, Al3+ Na+

For each of the following pairs, predict which atom or ion is larger

First Ionization Energy Energy required to remove the 1st e- from a neutral atom.

© 1998 LOGAL

•ionization energy - the energy needed to REMOVE an electron from an atom, in kJ/mol

•group trends•(first) ionization energy decreases from top to bottom along a group

•reason: outermost electron is farther and farther from the nucleus in larger atoms, so it is more easily removed

•periodic trends

•(first) ionization energy increases from left to right in a period

•reason: ―nuclear charge(+) increases; more attraction between electrons and protons

•Successive Ionization Energies

Mg 1st I.E. 736 kJ

2nd I.E. 1,445 kJ

Core e- 3rd I.E. 7,730 kJ

Large jump in I.E. occurs when a CORE e- is removed.

Al 1st I.E. 577 kJ

2nd I.E. 1,815 kJ

3rd I.E. 2,740 kJ

Core e- 4th I.E. 11,600 kJ

•Successive Ionization Energies

Large jump in I.E. occurs when a CORE e- is removed.

• Which atom has the higher 1st I.E.?

N or Bi

Ba or Ne

N

Ne

Examples

a. Mg, Na

b. S, O

c. Ca, Ba

d. Cl, I

Mg

O

Ca

Cl

e. Na, Al Al

f. Se, Br Br

For each of the following pairs, predict which atom has the higher first ionization energy.

Periodic Trends in Electronegativity

• electronegativity— tendency of an atom to attract electrons.

• noble gases do not have electronegativity values

• chemical bonds are determined by electronegativity differences between the bonding partners

• electronegativity trends are not completely regular

• fluorine = most electronegative element with a value of 4.0 (smallest anion formed)

• cesium = least electronegative element (largest cation formed)

•electronegativity decreases from top to bottom in a group

•electronegativity increases from left to right in a period

1

2

3

4 5

6

7

•Atomic Radius

Increases to the LEFT and DOWN

Atomic Radius

1

2

3

4 5

6

7

•First Ionization Energy

Increases UP and to the RIGHT

1

2

3

4 5

6

7

•electronegativity


1

2

3

4 5

6

7

•Electron affinity


Part 3: Electron Configuration

Electrons in Atoms • Niels Bohr. • Electrons are arranged in orbits around

the nucleus• The energy level of an electron is the

region around the nucleus where the electron is likely to be moving.

modern 3-D electron-cloud model - probability model

• Heisenberg Uncertainty Principle — it is not possible to know both the exact position and velocity of an object simultaneously

Modern electron cloud model

• orbitals are areas of high probability (~95%) of finding electrons

• Electrons can change energy level, by absorbing energy. When an electron absorbs a quantum of energy, it moves up to a higher energy level.

• When the electron falls from a higher energy level to a lower energy level, energy is released, and we see light

• Energy levels have sublevels —divisions within an energy level

1) many similar energy states grouped together in a level

2) different shapes: spherical, dumbbell, cloverleaf

There are 4 sublevels s, p, d, f(s p d f stand for sharp, principal, diffuse,

fundamental)

maximum number of e- in a principal energy level = 2n 2

• n = principal quantum number• = electron energy level or ― “shell”

(period) number• n = 1, 2, 3, 4, 5, 6, 7

electron maximums in the sublevels

• s can hold 2 e-• p can hold 6 e-• d can hold 10 e-• f can hold 14 e-

• Electrons fill orbitals in a certain way

•electron configuration - a specific electron arrangement in orbitals

Electron configuration:General Rules

•Pauli Exclusion Principle

–Each orbital can hold 2 electrons

with opposite spins.

•Aufbau Principle

–Electrons fill the lowest energy orbitals first.

–“Lazy Tenant Rule”

RIGHTWRONG

•Hund’s Rule

–Within a sublevel, place one e- per orbital before pairing them.

–“Empty Bus Seat Rule”

Different sections of the periodic table correspond to the different sublevels

Groups IA & IIA = s block

Groups IIIA – VIIIA = p block

Transition = d block

Inner transition = f block

Diagonal rule -to help us remember the order in which energy level subshells fill - follow the arrows 1s

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f7s 7p

O 8e-

•Orbital Diagram

Electron Configuration =

1 s2 2s2 2p4

Example

1s2 2s2 2p4

• the sum of the superscripts = the atomic number of the element

• superscripts are NOT exponents (nothing is being squared, etc.)

1s2 2s2 2p4

• *** valence configurations will be

s OR s and p ***

Condensed (Abbreviated) Electron Configurations

• use the previous Noble Gas as the starting point in brackets, then finish the configuration

Shorthand Configuration

S 16e-

Valence ElectronsCore Electrons

S 16e- [Ne] 3s2 3p4

1s22s2 2p6 3s2 3p4

•Longhand Configuration

• Example: Indium #49

• 1) complete: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 5p1

• 2) condensed: [Kr] 5s2 5p1

• 3) valence 5s2 5p1

• I Heart Electron Configuration– by Mark Rosengarten UF

http://www.youtube.com/watch?v=Vb6kAxwSWgU

THE END !

Download - Unit 3.2 Part I - Introduction to the Periodic Table Unit 3.2 Part I - Introduction to the Periodic Table

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