Unit 3 Periodic table all slides

Periodic Table History

1. Mendeleev ( 1800’s ): Father of the Periodic
Table
- 1st scientist to arrange elements

a. This arrangement was based on atomic mass.
b. The empty spaces in the Periodic Table proved to
be undiscovered elements.

2. Moseley ( 1850 ): Father of the Modern
Periodic Table

arranged the Periodic Table according to

atomic number

Mendeleev left gaps in his periodic table. He correctly predicted the
existence and properties of these unknown elements.

Periodic Law

• when the elements are arranged in order of

increasing atomic number, there is a periodic
repetition of their physical and chemical
properties .

As you go down and to the left, elements take
on more metallic properties.

Arrangement

•Groups or

families = columns

• 8 main columns

• all the elements in

a group have similar
but not identical
properties

Arrangement

•Periods = rows

• 7 rows
•the properties

of elements
change
greatly across
any given row

Review…

• Element – pure substances composed of only

one type of atom

• Compounds – two or more different elements

bonded together in a fixed proportion

• Compounds are always molecules, but molecules are

not always compounds

Metals - elements located to the

left of the zig-zag line

Nonmetals - elements located to

the right of the zig-zag line

(except H too)

Metalloids - elements touching the zig-zag

line ( B, Si, Ge, As, Sb, Te)

exceptions: (Al is a METAL)

Metal Properties

Physical properties

Chemical Properties

• good conductors of heat and

electricity

• lustrous - reflect light, shine

when they are polished

• flexible

• malleable- can be rolled

or hammered into sheets

• ductile - can be drawn

into wires "

• are solids at room

temperature except for
mercury

• have low ionization energy

(energy needed to remove
electrons)

• have low electron affinity

(attraction for electrons)

• tend to lose electrons easily

• form positiveions

(cations) when combining
with other atoms

• form ionic bonds

Nonmetal Properties
(opposite of metals)

Physical properties

Chemical Properties

• Usually exist as gases or

solids at room temperature

• brittle - not ductile or

malleable

• solids are dull - do not

reflect light even when
polished

• poor conductors of heat

and electricity (insulators)

• tend togainelectrons

• form negativeions

(anions) when combining
with other atoms

• have high electron

affinities (electron
attraction)

• have high ionization

energy

• produce covalent bonds by

sharing electrons with other
nonmetals

Properties of Metalloids

Physical Properties
Chemical
Properties

-- Solid
-- brittle
-- metallic luster
-- semi-conductors

-- intermediate in
electronegativity
(tendency to gain or
lose electrons)
-- will react with either
metals or nonmetals

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“Happy” Elements

•All elements want to be
“happy”
•https://youtu.be/C7dPqrmDWxs

•An atom wants a stable
outer energy level to
make them “happy”

How do they become happy?

• Atoms will gain or lose

electrons in order to become
stable.

• Usually the most stable

configuration involves having
8 valence e-

• 8 electrons = full octet

Stable Elements…

•Noble Gases do not need to transfer or share
any electrons because they already have 8
valence electrons and are stable

Octet Rule

•Atoms will form bonds with other atoms in

order to have eight valence electrons.

• Can share or transfer electrons with other

elements

• Elements will not just lose or gain electrons. They will

give to another element or take from another element

Oxidation Numbers

•The oxidation
number indicates
how many electrons
are going to be lost
or gained during
bonding.

• Range from -4 to +4

Predicting Oxidation Numbers

1. First, determine how many valence electrons the neutral atom

has.

2. Determine if the atom will lose or gain electrons.

• Ask yourself, which would be faster?
○ A) Losing down to zero
○ B) Gaining up to eight

3.Count the number of electrons gained or lost.

○ If you lose electrons, the charge will be (+)
○ If you gain extra electrons, the charge will be (-)

Practice

• What would the oxidation number be for the

following atoms?

A) Lithium

B) Sulfur

C) Aluminum

D) Iodine

E) Phosphorous

F) Carbon

G) Neon

Practice

• What would the oxidation number be for the

following atoms?

A) Lithium

B) Sulfur

C) Aluminum

D) Iodine

E) Phosphorous

F) Carbon

G) Neon0

1+

2-

3+

1-

3-

4±

Oxidation #s on the PT

Periods and Families

• Periods

• Do not have similar properties,

except they have the same
number of energy levels

• Families

• Similar properties

Lewis Dot Diagrams of Elements in
Periods 1 – 4 (Electron Dot Diagrams)
What do you notice?

Group 1 – Alkali Metals

• Examples: Li, Na, K, Rb

• Properties:

• Form+1 ions (lose one

valence e-)

• Usually react with

Halogen Family

• Low in electronegativity

• Most reactive metals

• H is not a member

Group 2 – Alkaline Earth Metals

• Examples: Be, Mg, Ca

• Properties:

• Found only in

compounds

• Form+2 ions (lose 2

valence e-)

• Second most reactive

metals

Group 3-11 – Transition Metals

• Examples: W, Cu, Ag
• Properties:

•Have multiple oxidation numbers

•Can form more than one binary
compound

•CuCl or CuCl2

•Unpredictable properties

•Usually good conductors, used for
wiring, jewelry, metalwork

•Have an incomplete d-orbital

•Most form colored ions in
solution

Group 12 - Zinc Family

• Examples: Zn, Cd, Hg
• Properties:

• Hg is liquid at room

temperature

• Stable electron

configurations

• Form +2 ions
• Zn is used in alloys

(mixtures of metals) with
copper to form brass and is
used to coat iron
(galvanizing)

Group 13 - Boron Family

• Examples: B, Al, Ga

• Properties:

• Form +3 ions

• All are metals except

boron (a metalloid)

• Aluminum is used for

coatings but is rarely
found outside of
compounds in nature

Group 14 - Carbon Family

• Examples: C, Si, Ge

• Properties:

• Have 4 valence e-

• Intermediate in

electronegativity (carbon
tends to form covalent
bonds)

• Silicon is a common

semiconductor material

Group 15 – Nitrogen Family

• Examples: N, P, As, Sb,

Bi

• Properties:

• Form -3 ions

• N and P are nonmetals

• As and Sb are metalloids

• Bi is a metal

Group 16 - Oxygen Family

(Chalcogens)

• Examples: O, S, Se, Te, Po
• Properties:

• Consists of three nonmetals

(O, S, Se), one metalloid (Te),
and one metal (Po)

• Form -2 ions (gain 2 e-)
• Tend to form covalent

compounds with other
elements

• Tend to exist as allotropes

(several different forms)

•Example: O2 and O3
(ozone)

Group 17 - Halogens

• Examples: F, Cl, Br, I

• Properties:

• Most active nonmetals

• Form -1 ions

• Tend to want to gain an

electron to form an ionic
compounds with metals
called metal halides

Group 18 – Noble Gases

• Examples: Kr, Ne, Ar

• Properties:

• Have a filled outer ring of

valence e- and therefore
tend not to form bonds

Bottom 2 Rows – Rare Earth

Metals

• 1st row – Lanthanoid

series

• 2nd row – Actinoid series
• Properties:

• Some are radioactive

• Usually silver,

silvery-white, or gray
metals

• Conduct electricity

Tuesday 10/3: Unit 3: PT trends

► Quiz tomorrow on PT and

trends

► Mon 10/9 Unit 3 test
► Thurs 10/12 -Benchmark 1

on Units 1-3 (end of Q1)

► Fri 10/13- NO SCHOOL- TWD
► Extra help is available

every day- use sign up
sheet on GC

►valence electrons and oxidation
states

►nuclear pull

►trends and Review of graphs

►Questions on hw review

►

Announcements:
Agenda:

Bellwork:

Turn in graphs on GC if you haven’t done so

Think about it: When you look at your graphs, what trends do
you see across periods 2&3, and down groups 1 & 17?

Trends graphs- When you look at your graphs, what trends
do you see across periods 2&3, and down groups 1 & 17?

Trends graphs- When you look at your graphs, what trends
do you see across periods 2&3, and down groups 1 & 17?

Why?

Each electron is
simultaneously attracted
to the nucleus as it is
repelled by the other
electrons.

The inner electrons
between the nucleus and
the electron of interest
make up the shield,
blocking the nucleus
from “pulling” on the
outer electrons.

Nuclear Pull

(aka Electron Affinity)

Nuclear pull aka electron affinity
sodium vs fluorine
►

electron affinity

►

atomic radius

►

ionization energy

►

electronegativity

Page 20

1)

Put on safely goggles

2)

Add a small amount of each metal the the appropriately
labeled test tube

3)

record your observations of the metal before the
reaction

4)

Fill test tube 1/3 full with water

5)

observe the metal’s reaction with water (what are signs
of a reaction?) and record

6)

What happens when I add the phenolphthalein?

7)

Dispose of the water into the waste beaker and dump
the metal into the waste beaker

Periodic trends

• We look at trends in 2 ways:

► down a group

► across a period

• * Trends are always opposite of

each other!

• Atomic Radius

► ½ distance between 2 adjacent nuclei

• Increases down a group
• Decreases across a period

Why?

# PELs
increase

Increasing nuclear charge

Ionic Radius

• Do you find ionic radius on Ref Table S?
• Rule: All positive ions will be smaller than their

"parent" atoms because they formed by losing e-

• Rule: All negative ions will be larger than

their "parent" atoms because they formed
by gaining e-.

More Periodic trends…

Which Group 1 metal is the most

reactive?

increases

decreases

• What is the trend for reactivity for the metals in

group 1?

• What is the trend for reactivity of metals/ metallic

character of elements across a period?

What is the most reactive

metal on the periodic table?

decreases

increases

Reactivity of non-metals/

non-metallic character

Reactivity of non-metals/

non-metallic character

► Which element is the most reactive

non-metal on the periodic table?

► Fluorine

e-

First Ionization Energy

Can be looked up on Reference

Table S

► Unit?
► KJ/mol

What is it?

► Ionization energy is the amount of energy required to remove an electron

from the outer shell (PEL) of an atom

► The first ionization energy is the amount of energy required to remove the

first electron from an atom

What are the trends for first

IE?

•First IE increases across
a period
•First IE decreases down
a group

What are the trends for first IE?

decreases

increases

Why?

# PELs
increase

Increasing nuclear charge

Explaining the trend…

► Increasing # PELS
► valence e- are farther from

nucleus, therefore they are easier
to remove

Explaining the trend…

► Increasing nuclear charge
► electrons are more difficult to

remove

Electronegativity (aka Electrolove)

► The ability of an atom to attract electrons
► introduced by Linus Pauling in 1932 who defined the new atomic property

as:

► "The power of an atom in a molecule to attract electrons to itself."

► Which element on table is best at gaining

e-?

► Fluorine

► It is __________.

e-

4.0

What are the trends for electronegativity?

decreases

increases

► Note: Noble gases do not have electronegativity values.
► Why not?

Effective Nuclear Charge

The shield between the
nucleus and the
electron is the primary
influence on the
effective nuclear
charge.

Effective nuclear
charge: charge felt on
valence electrons

Zeff = Z – S

Zeff = Effective Nuclear Charge
Z = # of protons
S = Number of inner “core” electrons

You will not need to do this calculation, but it is a VERY

good guide to figuring out nuclear pull!

Nuclear Pull:

Learning Check:

In the following pairs, pick the atom that has the
greater nuclear pull:

B or F

Mg or Si

S or Cl

From left to right in a period, electron affinity
______________.

From top to bottom in a group, electron affinity
______________.

Learning Check:

In the following pairs, pick the atom that has the
greater nuclear pull:

F

Si

Cl

From left to right in a period, electron affinity
increases.

From top to bottom in a group, electron affinity
decreases.

Atomic Radius

Definition: The distance from the nucleus to the outermost
electron. In the electron cloud. This measures the size of atoms!

• Measurement usually taken when

two atoms are bonded.

• Units of pm = picometers

(that’s a trillionth of a meter!)

Learning Check: Atomic Radius

In the following pairs, pick the atom that has the greater atomic
radius:

B or F

Mg or Si

S or Cl

From left to right in a period, atomic radius ______________.

From top to bottom in a group, atomic radius ______________.

Learning Check: Atomic Radius

In the following pairs, pick the atom that has the greater
atomic radius:

B

Mg

S

From left to right in a period, atomic radius decreases.

From top to bottom in a group, atomic radius increases.

Ionization Energy

Across a Period (left to right)

WHAT: Increases

WHY: Nuclear Pull

Down a Group (top to bottom)

WHAT: Decreases

WHY: More Energy Levels

Learning Check: Ionization Energy

In the following pairs, pick the atom that has the greater ionization
energy:

B or F

Mg or Si

S or Cl

From left to right in a period, ionization energy ______________.

From top to bottom in a group, ionization energy ______________.

Learning Check: Ionization Energy

In the following pairs, pick the atom that has the greater ionization
energy:

F

Si

Cl

From left to right in a period, ionization energy increases.

From top to bottom in a group, ionization energy decreases.

Electronegativity

Definition: The ability of an atom to
attract electrons from ANOTHER
atom into a bond.

Measured on the Pauling Scale

Values range from 0.0 to 4.0 and have
NO UNITS.

Electronegativity

Across a Period (left to right)

WHAT: Increases
WHY: Nuclear Pull

Down a Group (top to bottom)

WHAT: Decreases
WHY: More Energy Levels

Learning Check: Electronegativity

In the following pairs, pick the atom that has the greater
electronegativity:

B or F

Mg or Si

S or Cl

From left to right in a period, electronegativity ______________.

From top to bottom in a group, electronegativity ______________.

Learning Check: Electronegativity

In the following pairs, pick the atom that has the greater
electronegativity:

F

Si

Cl

From left to right in a period, electronegativity increases.

From top to bottom in a group, electronegativity decreases.

Going across the table from left to right within a row
or period
• the number of protons increases
• the pull on the electrons (electronegativity) increases

(Table S)

• the covalent atomic radiusdecreases
• metallic properties decrease
• Ionization energy increases ( energy needed to

remove the most loosely held electron) (Table S)

• There is a gradual change from positive oxidation

statestonegativeoxidation states

• the number of valence electrons increases

Going down the table within a group or family

• the number of protons increases

• the number of shells increases

• the atomic radius increases

• the pull on the electrons decreases

• metallic properties increase

• the number of valence electrons remains the same

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Unit 3 quiz, 20 MC on formative
Put away everything except for your
school chromebook
Shift to the CORNER of your table

When you are finished, find the assignment on
GC

Reactivity of metals lab…page 15
(turn in on GC)

You will add each of
the following metals
to water and observe
and record….

1) calcium
2) iron
3) aluminum
4) magnesium

I will demonstrate the
following metals
reacting with water,
you will observe and
record….

1) lithium
2) sodium
3) potassium (video)
4) francium (video)