Topic 3.1-3.2: Periodic TrendsEssential Ideas:

3.1: The arrangement of elements in the Periodic Table helps to predict their electron configurations3.2: Elements show trends in their physical and chemical properties across periods and down groups

Types of Elements• Metals – good conductors of heat and electricity;

mostly solid at room temperature; VERY diverse properties; Brittle or firm; Many have luster (shiny)

• Nonmetals – poor conductors of heat and electricity; many are gases at room temperature

• Metalloids/Semiconductors – possess properties of metals and nonmetals; all are solid at room temperature; semiconductors of electricity; used in electronics

The Periodic Table• Organized from left to right by atomic number

• Vertical columns of the p.t. are called Groups or Families• Contain elements with similar chemical properties

• Horizontal rows of the p.t. are called Periods

• As the elements fall into columns based on their properties, certain trends develop.

Key Words/Phrases• Principal Energy Level• Valence Electrons• Valence Electron Energy Level• Valence Electrons attracted to positive protons

in the nucleus• Nuclear Charge• Effective Nuclear Charge• Repulsion between Electrons• Full Valence Electron Shell• Mutual Repulsion• Electron Shielding

Nuclear Charge• Negatively charged electrons are attracted to

the positively charged nucleus• Related to the number of protons in the

nucleus - as proton # increases in an atom, the nuclear charge also increases (becomes more positive)

Nuclear Charge Trend• General Trend:

• Groups: L→ R Nuclear Charge Increases• Periods: Top→ Bottom- Nuclear Charge Increases

Shielding

• The more energy levels/interior electrons an atom has, the more “shielded” the outer electrons are to the nuclear charge

• Energy levels act like “Walls” to block nuclear charge from the valence electrons - makes them easier to lose (less attraction!)

•

Shielding Trend• General Trend:

• Groups: L→ R Shielding is Constant (no change to interior electrons)

• Periods: Top→ Bottom- Shielding Increases with more energy levels

Effective Nuclear Charge• Charge experienced by the valence electrons

• Less than nuclear charge due to shielding

Effective Nuclear Charge = Nuclear Charge (#p) - Interior e- (total e- - ve- )

• Energy levels act like “Walls” to block nuclear charge from the valence electrons - makes them easier to lose (less attraction!)

•

Effective Nuclear Charge Calculation

Effective Nuclear Charge• General Trend:

• Groups: L→ R Increases (more protons, same # shielding e- ) • Periods: Top→ Bottom- Constant (more protons, more

shielding electrons from more energy levels)

Explaining ENC • SHORT, CONCISE, KEY WORD-RICH ANSWERS

Explain why ENC increases across a period. ● ENC increases across a period because the nuclear

charge increases due to an increase in positive protons while shielding negative electrons remain the same, causing more attraction between the nucleus and valence electrons.

Explain why ENC remains constant down a group.● ENC remains constant down a group because there are

more principal energy levels adding more negatively charged shielding electrons. Though there are more positive protons, the shielding electrons block the attraction from the ve- to the nucleus

Atomic Radius• The atomic radius is ½ the distance from the nucleus

of one atom to the nucleus of another, identical atom when bonded. (ex. O

2)

• High ENC = Small Radius• High Nuclear Charge with Few Shielding Electrons• Strong Attraction between the nucleus and

valence e- pulls the ve- in close

Atomic Radius• General Trend:

• Groups: L→ R Decreases (Increased ENC = strong attraction) • Periods: Top→ Bottom- Increases (More energy levels w/ ENC

remaining constant)

Atomic Radius - Trend

Explaining Atomic RadiusExplain why the radius of F is smaller than Lithium. ● F and Li have the same number of principal energy

levels meaning they have the same amount of shielding. F, however, has a stronger attraction between the nucleus and the valence electrons due to its increased number of positive protons (higher nuclear charge)

Explain why Ra has a larger radius than Be.● Ra has 7 principal energy levels while Be has only 2.

Ra has more low energy level shielding electrons blocking the attraction to the nucleus.

Ionic Radius• Cations: Lose ve- ∴ lose an energy level

• Cations smaller than parent atom• Anions: gain ve- ∴ increase e- repulsion in valence

energy level pushing it out larger• Anions larger than parent atom

Ionic Radius• General Trend:

• Groups: Decrease L→ R (increasing ENC = attraction) ■ 1-14 Cations● 14 smaller than 15

■ 15-17: Anions• Periods: Top→ Bottom (Increasing =

more energy levels)

Ionic Radius - Trend

Explaining Ionic Radius Arrange Cl, Cl+1 and Cl-1 by increasing radii. Explain your reasoning.

Cl+1, Cl, Cl-1

● Cl-1 is the largest because it gained 1 electron forming a full valence energy level. The full energy level has more repulsion between the electrons, causing the radius to increase. Cl+1 has the smallest radius because it lost 1 electron, removing 1 energy level. This reduces the number of shielding electrons and increases the attraction to the nucleus. The attraction from the positive protons in the nucleus pull the ve- in close, causing a small radius.

Ionization Energy• A measure of the attraction between the nucleus and

the outer electronsM(g) → M+1(g) + e-

• The energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state

• IE increases with an increase in ENC (strong attraction = large amount of energy to remove an electron)

Ionization Energy• General Trend:

• Groups: L→ R Increases (more protons, same # shielding e- ) • Periods: Top→ Bottom- Decreases (constant ENC, more

energy levels = greater distance b/w p and e- = less attraction)

Explaining IE Explain the IE trend Across a Period. • IE increases across a period because the effective

nuclear charge increases. Positive protons are increasing while shielding negative electrons remain the same, causing more attraction between the nucleus and valence electrons.

Explain the IE trend down a Group. • IE decreases down a group because the electrons

are further away from the nucleus and electrons in the lower energy levels are blocking the attraction causing electron shielding.

Ionization Energy - Trend

Exceptions to IE• Red Dip (s2p1 electron configuration): the lone electron in

the p orbital is easier to remove than a completely full s orbital

Exceptions to IE• Blue Dip (s2p4 electron configuration): the 4th electron in the

p orbital is easier to remove because it is being repelled by the other electron in the orbital.

Explaining IE Exceptions ● What causes B to have a lower IE than Be? ○ B has a lower IE than Be because the 2p

electrons are slightly higher in energy than the 2s electrons, so the ionization energy for B is lower than for Be

● What causes O to have a lower IE than N? ○ O has a lower IE than N because the p

x , p

y , and

pz only contain one electron. The extra electron

in O causes a pair and repulsion making it easier to remove, hence giving O a lower IE than N.

Electron Affinity• Opposite of Ionization Energy

M(g) + e- → M-1(g)

• The energy change when one mole of electrons are added to one mole of gaseous atoms.

• Generally exothermic (energy released)

○ High Affinity = release large amount of energy

■ Exothermic = negative number for energy

■ Endothermic = positive number for energy

Electron Affinity• General Trend:

• Groups: L→ R Increases (more protons, same # shielding e- ) • Periods: Top→ Bottom- Decreases (constant ENC, more

energy levels = greater distance b/w p and e- = less attraction)

Exceptions to Electron Affinity• Electron Affinity decreases (+ energy, endothermic) with

increased repulsion and decreased attraction (shielding)

Exceptions to Electron AffinityExplain why Beryllium (red circle) has an endothermic/very low electron affinity? ● Be must have an electron placed into the 2p orbital

which is further from the nucleus and so experiences less attraction due to shielding from electrons in the 2s orbital.

Explain why Nitrogen (blue circle) has an endothermic/very low electron affinity? ● The added electron must occupy a p orbital that is

already singly occupied. The attraction between the electron and the atom would be reduced due to repulsion from the other p electrons.

Exceptions to Electron Affinity

Explain why Neon (yellow circle) has an endothermic/very low electron affinity? ● Neon has a full valence energy level. Any

additional electron must be placed in the next energy level which is farther from the nucleus. This added electron will experience increased shielding and decreased attraction to the nucleus.

Electronegativity• The ability of an atom in a compound to attract

electrons from other atoms in a covalent bond

• High EN = strong electron pulling power

• Low EN = weak electron pulling power

Electronegativity - Trend

Electronegativity PracticePredict what will occur if a very high electronegativity atom gets close to a very low electronegativity atom. ● An atom with high EN will have a strong attraction

to the electrons of another atom while the low EN atom will not. The strong EN atom will pull the electrons from the weak atom close to it in a polar covalent bond. A large enough difference in electronegativity will cause the strong EN atom to completely remove the e- (ionize) the weak EN atom and form an ionic bond.