HONORS
CHEMISTRY
September 18-19, 2013
How do we know what the filling
order is?

What chemistry tool might we rely on?
(Stop)
Electron Configurations and the
Periodic Table

Valence electron configurations repeat
down a group
Ground state electron configurations

Example: Li
 atomic
number = 3
 nucleus has 3 protons
 neutral atom has 3 electrons

2 electrons in 1s orbital, 1 electron in 2s orbital
2s
1s
Different ways to show electron
configuration
Energy level diagram
Box notation

1s
2s

2s
1s
Spectroscopic notation
Li 1s2 2s1
Read this “one s two”
not “one s squared”
Write the superscript 1.
Don’t leave it blank
Using the Periodic Table
The last subshell in the electron configuration is one of these
(row #) s
(row # – 1) d
(row #) p
(row # – 2) f
The f-block is inserted into to the dblock
Electron configuration of O

Atomic number of O = 8 so neutral atom has 8 e–
Electron configuration of Co

Atomic number of Co = 27 so neutral atom has 27 e–
Simplifying electron configurations
Shorthand Noble Gas
Configuration

Build on the atom’s noble gas core

He 1s2
O 1s22s22p4
O [He]2s22p4
Ar 1s22s22p63s23p6
Co 1s22s22p63s23p64s23d7
Co [Ar]4s23d7


1s

2s
  
2p
             
1s 2s
2p
3s
3p
4s
3d

Noble Gases
Far right of the periodic table
 These elements are extremely unreactive
or inert
 They rarely form compounds with other
elements

Noble Gas electron configurations

What is the electron configurations for
Neon

Abbreviated way to write configurations
 Start
Br
 Ba

with full outer shell then add on
Noble Gases
Neon- emits brilliant light when stimulated by
electricity – neon signs- 4th most abundant element
in the universe.
 Helium- light non reactive gas- used balloonsinexpensive, plentiful and harmless
 Radon- radioactive gas- can cause cancercolorless, odorless emitted from for certain rocks
underground

Why are we doing all of this?
 Properties
of atoms correlate with the
number and energy of electrons
 Electron configurations are used to
summarize the distribution of electrons
among the various orbitals
Electron configuration of ions
What is an ion?
 How many electrons does Cl1- have?

 What
is the electron configuration for the
chloride ion?

How many electrons does Ca2+ have?
 What
is the electron configuration for the
calcium ion?

What do you notice?
Why is this important
Valence electrons

Electrons in the outermost energy level
 Where
all the action occurs
The f-block is inserted into to the dblock
Find the electron configuration of
Au

Locate Au on the periodic table
Find the electron configuration of
Au

Au [Xe]

The noble gas core is Xe
Find the electron configuration of
Au

Au [Xe]6s2

The noble gas core is Xe
From Xe, go 2 spaces across the s-block in the 6th row
 6s2

Find the electron configuration of
Au

Au [Xe]6s24f14

The noble gas core is Xe
From Xe, go 2 spaces across the s-block in the 6th row
 6s2
Then detour to go 14 spaces across the f-block  4f14
 note: for the f-block, n = row – 2 = 6 – 2 = 4


Find the electron configuration of Au

Au [Xe]6s24f145d9

The noble gas core is Xe
From Xe, go 2 spaces across the s-block in the 6th row  6s2
Then detour to go 14 spaces across the f-block  4f14
 note: for the f-block, n = row – 2 = 6 – 2 = 4
Finally go 9 spaces into the d-block on the 6th row  5d9
 note: for the d-block, n = row – 1 = 6 – 1 = 5



Practice

Draw the orbital diagram for sulfur.
 What

ion does sulfur want to form and why?
Draw the orbital diagram for Potassium.
 What
ion does sulfur want to form and why?
What does this mean
 Properties
of atoms correlate with the
number and energy of electrons

Atoms like to have full outer shells.
Refer to Atomic Structure
Worksheet
Periodic Trends
Preview

4 Periodic Trends
 Atomic

Size/Radius
Ionic Size (**)
 Ionization
Energy
 Electronegativity

2 main factors affect periodic trends
 Number
of electron shells (group)
 Effective Core Charge (ECC) (period)
Term (Refer to Definition Sheet)

Effective Core Charge (ECC)

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 (label ECC on P.T.)
Term

Electron Shell
 Pattern
across the period?
 Pattern down the group?
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
from top to bottom
 More energy levels or quantum levels (or “shell”) as
you go down a group – atomic radius increases

Period Trend
 Increases
from right to left
 All electrons in the same energy level. Increased # of
protons holds them closer to nucleus.
 Decrease in Effective Core (Nuclear) Charge (ECC)

Calculate ECC for elements in period 2
Table of
Atomic
Radii
Period Trend:
Atomic
Radius
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
from top to bottom
 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
from bottom to top.
 The closer outer shell electrons are to the
nucleus the harder they are to remove.

Period Trend
 Increases
from left to right.
 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

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
from bottom to top
 As radius decreases, electrons are closer to
the nucleus (decrease in number of electron
shells)

Period Trend
 Increases
from left to right
 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
Summarize the Trends

Questions???
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 Core Charge:
 It
is 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

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
Atomic Structure Worksheet
 5-3 Worksheet
 Study for Definition Quiz on Monday

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.
The Periodic Table

Dmitri Mendeleev – credited for the first
periodic table in 1869.
 He
had put element names and a few of their
properties on cards and then arranged them in
various ways to help his students learn them
more easily.
 Arranged them so elements in the same
column have similar properties.
Reactivity: Metal/NonMetal
Trends
ELEMENT CLASSES
Periodic Song: http://www.privatehand.com/flash/elements.html
Reading the periodic table

Groups or families – vertical columns

Periods – horizontal rows
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
Questions?
White Board

Refer to Periodic Trends Review
Electron Configurations and
Periodic Trends
Write the electron configuration and draw
an orbital diagram for each element
 Order each group of elements or ions
based on given data for each property
requested on card
 Use the orbital diagrams to explain the
pattern. (does it agree with the “trend”)

Objectives
Use the periodic table to write electron
configurations
 Use the periodic table to obtain
information about the properties of
elements
