Periodic Trends

3/11/16 SLO

Students will be able to describe trends

among elements for atomic size,

ionization energy, ionic size, and how

ions are formed.

How can we further use the

Periodic Table to make

predictions about the elements?

Recognize trends in the elements as we

move across a period or down a group.

We can determine the following characteristics

of an element using the periodic table:

Atomic Radius

Ionization Energy

Electronegativity

2 Trends to explain all other trends

Electron Shielding- the reduction of the

attractive force of the nucleus for the outer

electrons

caused by electrons in energy levels between

the nucleus and the outer electrons

Nuclear charge- the number of protons in

the nucleus.

More protons = increased nuclear charge and

increased attraction between the nucleus and

electrons.

Atomic Radius

Atomic Radius- ½ of the distance

between the nuclei of two atoms of the

same element in a diatomic molecule

Atomic Radius

Atomic Radius

Independent Practice

Using your periodic table and the atomic

radius chart, determine which of the elements

in each pair has a larger atomic radius:

1. Cesium (Cs) and Potassium (K)

2. Calcium (Ca) and Gold (Au)

3. Rubidium (Rb) and Strontium (Sr)

4. Oxygen (O) and Sulfur (S)

5. Xenon (Xe) and Neon (Ne)

6. Aluminum (Al) and Tin (Sn)

7. Helium (He) and Fluorine (F)

8. Boron (B) and Bromine (Br)

Why do elements have different

reactivities? It all depends on valence electrons

Almost every atom is stable (not reactive)

if it has 8 valence electrons

Exceptions: those that only need 2 electrons

to fill the outer electron shell: H, He, Li, Be, B

If an atom has less than 8 valence

electrons, it will gain, lose, or share

electrons to become stable

Ions= When atoms gain or lose electrons,

they become charged

Becoming Stable

Atoms that have 1-3 valence electrons will

LOSE (or share) electrons to become

stable

When electrons are lost, this causes a charge

More protons than electrons results in a

positively charged ion called a cation

Example: potassium has 1 valence electron.

It loses an electron to become K+

Calcium has 2 valence electrons. It loses 2

electrons to become Ca2+

Becoming Stable

Atoms that have 5-7 valence electrons will

GAIN (or share) electrons to become

stable

When electrons are gained, this causes a

charge

More electrons than protons results in a

negatively charged ion called an anion

Example: Chlorine has 7 valence electrons. It

gains 1 electron to be stable. Cl-

Oxygen has 6 valence electrons. It gains 2

electrons to be stable O2-

Forming cations

Forming anions

In addition to Atomic Radius, we

have an Ionic Radius

Where do ions or ionic radius come from?

Periodic Trends

Trend for atomic size (atomic radius)-

Down a group, size increases

Occurs because # of E levels increases &

Electrons shielding reduces amount of attraction

between nucleus and outer electrons

Across a period, size decreases

# of protons increases (nuclear charge increases),

pulling electrons closer

Electron shielding doesn’t change because

electrons are added to the same energy level

Atomic Radius Decreases

Atomic Radius Increases

Ionization Energy Ion- atom that gains or loses electrons

Ionization Energy- energy required to remove an electron. Energy must be added to overcome the attraction of the

positive charge of the nucleus

X(g) X+(g) + e- 1st ionization

X+(g) X2+(g) + e- 2nd ionization

Easiest to remove 2 electrons from 2A Because there are 2 valence

Easiest to remove 3 electrons from 3A Because there are 3 valence

Outer shell electrons are easier to remove than other electrons!

Ionization Energy

Periodic Trends

Ionization energy

Down a group- decreases because electrons

are held more loosely due to increased

electron shielding

Across a period- increases because electrons

are held more tightly due to increased nuclear

charge (increased # of protons in the nucleus)

Ionization Energy Increases

Ionization Energy Decreases

Periodic Trends

Metals form positive ions- Cations

More likely to lose electrons (lower ionization

energy)

Nonmetals form negative ions- Anions

More likely to gain electrons (higher ionization

energy)

Periodic Trends

Ionic Radii Trends

Cations- smaller than neutral atom because

fewer electrons result in greater attraction by

nuclei

Anions- larger than neutral atom because

more electrons result in less attraction by

nuclei

Across a period- size decreases

Down a group – size increases

Atom versus Ion

Ion Size Decreases

Ion Size Increases

CLASSWORK!

Reading Assignment: Sec 6.3

Written Assignment: pg. 182, #18-21, 24, 25

Periodic Trends Part 2

Do Now

Complete the half sheet of paper relating

to the Periodic Trends that we have

already discussed in class.

Periodic Trends Part 2

SLO

Students will be able to describe trends

in electronegativity and electron affinity

on the periodic table.

Homework Check!

18. Atomic size generally increases as you move

down a group, and decreases from left to right

across a period.

19. Ions form when electrons are transferred

between atoms.

20. First ionization energy generally decreases as

you move down a group and increases from left

to right across a period.

21. Anions are larger and cations are smaller than

the atoms from which they form.

Vocabulary Review

Ion

Electronic Shielding

Alkaline Earth Metals

Nuclear Charge Metal Period Nonmetal Metalloid

Cation

Ionization Energy

Anion

Transition Metals

Inner Transition Metals

Periodic Law

Noble Gases

Group

Representative Elements

Alkali Metals

Halogens

Periodic Trends

Electronegativity- tendency for the atoms

of the element to attract electrons when

the atoms are part of a compound

Fluorine (F) is most electronegative

Noble gases- no electronegativity values-

don’t form compounds

Periodic Trends

Electronegativity Trends-

Down a group – decreases- since electron shielding results in less attraction for electrons by the nucleus

Across a period- increases- since there is a higher atomic number and consistent electron shielding result in more attraction for electrons

Electronegativity allows you to predict bond type: covalent (includes polar vs. nonpolar) and ionic

Electronegativity Increases

Electronegativity Decreases

Electron Affinity

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

Example: F(g) + e- F-(g)

Ho (ENERGY) = -328.0 kJ/mol

Think of it like electronegativity without the

need to bond… It still has to do with attraction

for electrons.

Trends in Electron Affinity

Down a group, it decreases because

electron shielding blocks some of the

attraction from the nucleus

Across a period, it increases because

nuclear charge increases, attracting

electrons more strongly.

Electron Affinity Increases

Electron Affinity Decreases

Periodic Trends

Knowledge of trends in electron shielding

and nuclear charge explain all other trends

http://www.teachersdomain.org/resource/ls

ps07.sci.phys.matter.graphperiodic/

1. Which of the following sequences is

correct for atomic size?

Mg > Al > S

Li > Na > K

F > N > B

F > Cl > Br

6.3 Section Quiz

6.3 Section Quiz

2. Metals tend to

gain electrons to form cations.

gain electrons to form anions.

lose electrons to form anions.

lose electrons to form cations.

6.3 Section Quiz

3. Which of the following is the most

electronegative?

Cl

Se

Na

I

CLASSWORK!

Written Assignment: pg. 186, #38, 40, 41,

43, 44, 45

Summary of Trends

6.3

Atomic Size Increases

Incre

ases

Decreases

Decre

ases

Size of cations Shielding Nuclear Charge Electronegativity Ionization energy Size of anions Ionic size Constant

Periodic Trends Part 2

Do Now

Arrange the following elements in order of

increasing atomic radius:

Radon (Rn) Xenon (Xe)

Nickel (Ni) Cobalt (Co)

Sodium (Na) Potassium (K)

Tellurium (Te) Antimony (Sb)