Friday 12/04/15
Objectives
Understand the general trends in atomic
properties in the periodic table
 Understand the nature of bonds and their
relationship to electronegativity

Reading the periodic table

Groups or families – vertical columns

Periods – horizontal rows
Term

Effective Nuclear Charge (ENC)

1) The net charge that pulls on the valence electrons in an
atom. The greater the effective core charge, the greater
the pull. It is determined by subtracting the number of
core electrons from the number of protons in the
nucleus
 For example: Magnesium
Periodic Trends

Atomic radius


The distance from the
center of an atoms
nucleus to it’s outermost
electron
Measure of atomic size
Periodic Trends
Graph the first 20 elements.
What is the trend down a group? Across a Period?

Atomic radius
Periodic Trends
Atomic Radius

Group Trend
 Increases
down a group
 More energy levels or quantum levels (or “shell”) as
you go down a group – atomic radius increases

Period Trend
 Decreases
across a period
 All electrons in the same energy level. Increased # of
protons holds them closer to nucleus.
 Increase in Effective Core (Nuclear) Charge (ECC)

Calculate ECC for elements in period 2
Table of
Atomic
Radii
Period Trend:
Atomic
Radius
Periodic Trends Worksheet
 Work
with a table partner to answer
questions:
1,5a,
6, 8,
Periodic Trends


Ionic Size



Size of an atom when
electrons are added or
removed.
Electrons removed
atom becomes smaller.
Electrons added atoms
become larger
Why?

Electron-Electron
Repulsion
Ionic Size
Cations
 Positively charged ions formed when
an atom of a metal loses one or
more electrons
 Smaller than the corresponding
atom
 Negatively charged ions formed
when nonmetallic atoms gain one
Anions
or more electrons
 Larger than the corresponding
atom
Periodic Trends
Graph the first 20 elements.
What is the trend down a group? Across a Period?
 Ionic Size (label P.T.)
Table of
Ion
Sizes
Ionic Size

Group Trend
 Increases
down a group
 More energy levels as you go down a group – ionic
size increases

Period Trend
 Decreases
as atoms lose more electrons
 Increases dramatically as atoms start gaining
electrons, decreases as atoms gain fewer electrons.
Periodic Trends

Ionization Energy


Energy needed to remove
one of the electrons on
an atom’s outer shell.
How strongly does an atom
hold it’s outermost electron.
Periodic Trends
Graph the first 20 elements.
What is the trend down a group? Across a Period?
 Ionization Energy
Ionization Energy

Group Trends
 Increases
up a group.
 The closer outer shell electrons are to the
nucleus the harder they are to remove.

Period Trend
 Increases
across a period.
 The more electrons in the outer shell the
harder it is to remove one.
 Increase in Effective Core Charge (ECC)
Periodic Trend:
Ionization Energy
Periodic Trends Worksheet
 Work
with a table partner to answer
questions:
3,
5b, 7, 9
Periodic Trends

Electronegativity



Is a measure of the level of
attraction (pull) an atom
exerts on the electrons of
another atom.
Ability of an atom to attract
electrons
Which elements want to gain
electrons the most?
Periodic Trends
Graph the first 20 elements.
What is the trend down a group? Across a Period?
 Electronegativity
Periodic Table of Electronegativities
Electronegativity

Group Trend
 Increases
up a group
 As radius decreases, electrons are closer to
the nucleus (decrease in number of electron
shells)

Period Trend
 Increases
across a period
 The more electrons in the outer shell (up to 7)
the more the atom wants to attract electrons

Exception: Trend does not apply to Noble Gases
 Increase
in Effective Core Charge (ECC)
Periodic Trend:
Electronegativity
Periodic Trends Worksheet
 Work
with a table partner to answer
questions:
2,
4, 5c, 10, 11
Summary of
Periodic Trends
Practice
1.
Se and Br
1.
2.
2.
P, S, Se
1.
2.
3.
Largest atom
Highest Ionization Energy
Cl, Cl1-, Br, Br11.
4.
Smallest atom
Lowest Ionization Energy
Largest ionic size
Mg, Mg2+, Na, Na1+
1.
Smallest ionic size
Atomic Properties Definitions
For Quiz – Monday

Effective Nuclear Charge:
 It
is the net charge that pulls on the valence electrons
in an atom.
 The greater the effective nuclear charge, the greater
the pull.
 It is determined by subtracting the number of core
electrons from the number of protons in the nucleus

Valence Electrons
 Are
found in the outermost, valence, electron shell
(Bohr model) of the atom

Core electrons
 occupy
all of the inner electron, core, shells
Atomic Properties Definitions

Ionization Energy:



Atomic size




Energy needed to remove an electron from an atom or molecule.
The higher the effective core charge and lower the number of
electrons shells, the greater the ionization energy
How big (e.g., radius) an atom is
Atomic radius is measured from the center of the nucleus to the
valence electron shell.
The higher the effective core charge and lower the number of
electron shells, the smaller the atom.
Electronegativity


Measure of the level of attraction (pull) an atom exerts on the
electrons of another atom.
The higher the effective core charge and lower the number of
electron shells, the greater the electronegativity
Homework

Read pages:
 327-331

Answer questions: Pg 336 69-78
 343-345


Answer questions: Pgs 374-375, 7-20
Due 12/09
Periodic Table
Objective: Students know how to
use the periodic table to identify
alkali metals, alkaline earth metals,
transition metals, metals, semimetals
(metalloids), nonmetals, halogens and
noble gases.
Alkali Metals
 All
alkali metals have 1
valence electron
 They are very reactive
 Reactivity of these elements
increases down the group
 Alkali metals:
Potassium, K
reacts with
water and
http://video.google.com/videopl must be
stored in
ay?docid=kerosene
2134266654801392897#
Alkaline Earth Metals
All alkaline earth metals have 2 valence
electrons
 Alkaline earth metals are less reactive than
alkali metals
 The word “alkaline” means “basic”
 common bases include salts of the metals
 Ca(OH)2
 Mg(OH)2

Properties of Metals
 Metals are good
conductors of heat and
electricity
 Metals are malleable
 Metals are ductile
 Metals have high tensile
strength
 Metals have luster
Transition
Metals
Copper, Cu, is a relatively soft
metal, and a very good
electrical conductor.
Mercury, Hg, is the only
metal that exists as a liquid
at room temperature
Properties of
Metalloids
 They have properties of both
metals and nonmetals.
Metalloids are more brittle
than metals, less brittle than
most nonmetallic solids
 Metalloids are
semiconductors of electricity
 Some metalloids possess
metallic luster
Silicon, Si – A Metalloid
 Silicon has metallic luster
 Silicon is brittle like a
nonmetal
 Silicon is a semiconductor of
electricity
Other metalloids include:
 Boron, B
 Germanium, Ge
 Arsenic, As
 Antimony, Sb
 Tellurium, Te
Nonmetals
 Nonmetals are poor
conductors of heat and
electricity
 Nonmetals tend to be
brittle
 Many nonmetals are gases
at room temperature
Carbon, the graphite in “pencil
lead” is a great example of a
nonmetallic element.
Examples of Nonmetals
Sulfur, S, was
once known as
“brimstone”
Graphite is not the only pure
form of carbon, C. Diamond is
also carbon; the color comes
from impurities caught within
the crystal structure
Microspheres of
phosphorus, P, a
reactive
nonmetal
Halogens
 Halogens all have 7 valence
electrons
Halogens in their pure form are
diatomic molecules (F2, Cl2, Br2, and
I2)
Chlorine is a yellow-green
poisonous gas
Noble Gases
Noble gases have 8 valence electrons
(except helium, which has only 2)
•they are chemically unreactive
• Colorless, odorless and unreactive; they
were among the last of the natural
elements to be discovered