Download Periodic Trends

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?
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Recognize trends in the elements as we
move across a period or down a group.
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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
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Electron Shielding- the reduction of the
attractive force of the nucleus for the outer
electrons
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caused by electrons in energy levels between
the nucleus and the outer electrons
Nuclear charge- the number of protons in
the nucleus.
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More protons = increased nuclear charge and
increased attraction between the nucleus and
electrons.
Atomic Radius
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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
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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
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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
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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 Increases
Atomic Radius Decreases
Ionization Energy
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Ion- atom that gains or loses electrons
Ionization Energy- energy required to remove an
electron.
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Easiest to remove 2 electrons from 2A
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Because there are 2 valence
Easiest to remove 3 electrons from 3A
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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
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 Decreases
Ionization Energy Increases
Periodic Trends
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Metals form positive ions- Cations
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More likely to lose electrons (lower ionization
energy)
Nonmetals form negative ions- Anions

More likely to gain electrons (higher ionization
energy)
Periodic Trends
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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 Increases
Ion Size Decreases
CLASSWORK!

Reading Assignment: Sec 6.3
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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
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Ion
Electronic Shielding
Alkaline Earth
Metals
Nuclear Charge
Metal
Period
Nonmetal
Metalloid
Cation
Ionization Energy
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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

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Fluorine (F) is most electronegative
Noble gases- no electronegativity valuesdon’t form compounds
Periodic Trends

Electronegativity TrendsDown 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
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Electronegativity allows you to predict
bond type: covalent (includes polar vs.
nonpolar) and ionic
Electronegativity Decreases
Electronegativity Increases
Electron Affinity
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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.
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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.
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Electron Affinity Decreases
Electron Affinity Increases
Periodic Trends
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Knowledge of trends in electron shielding
and nuclear charge explain all other trends
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http://www.teachersdomain.org/resource/ls
ps07.sci.phys.matter.graphperiodic/
6.3 Section Quiz
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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
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2.
 gain
Metals tend to
electrons to form cations.
 gain electrons to form anions.
 lose electrons to form anions.
 lose electrons to form cations.
6.3 Section Quiz
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3.
Which of the following is the most
electronegative?
 Cl
 Se
 Na
I
CLASSWORK!
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Written Assignment: pg. 186, #38, 40, 41,
43, 44, 45
6.3
Summary of Trends
Increases
Increases
Decreases
Ionization
Ionicof
size
energy
Size
Size
Electronegativity
Atomic
Nuclear
Shielding
of
anions
cations
Size
Charge
DecreasesConstant
Periodic Trends Part 2
Do Now
Arrange the following elements in order of
increasing atomic radius:
Radon (Rn)
Nickel (Ni)
Sodium (Na)
Tellurium (Te)
Xenon (Xe)
Cobalt (Co)
Potassium (K)
Antimony (Sb)