The Periodic Table

I. History of the Periodic Table

A. Dmitri Mendeleev (1860’s)

• 1st to arrange elements into an organized table

• Organized them by _____________________

• Noticed there was a periodic or repeating pattern in

the elements’ physical and chemical properties

when he placed them in specific columns

• Wasn’t perfect- resulted in gaps and some elements

were out of order

increasing atomic mass

B. Henry Moseley (1900’s)

• Revised Mendeleev’s table by suggesting that the

properties of elements were based on

_____________________

o Argued that the properties depended more on

the structure of the atom than the atomic

mass

• Much more effective – no more gaps or elements

out of order

• Modern day PT still arranged by atomic number

atomic # or # of protons

II. The Periodicity of the Table

• Periodic = ________________________

o A main reason for these patterns or behaviors in elements

is their desire _____________

Stability

• Stability means having the outermost occupied electron shell

full - ______________

o Full valence shell = ________________

o _____________

o Exception: Hydrogen and Helium are stable with 2

valence electrons– why?

• This desire to be stable is why atoms gain or lose electrons to

form ions, why elements bond, why elements have the

properties they do, etc.

repeating patterns or behaviors

to be stable

full valence shell

8 valence electrons

Eight is great!!

Because the first electron level is full at 2 electrons!

III. Arrangement of the Periodic Table

Periods

• The periodic table is arranged in several different

ways. One way is by periods and groups.

A. Periods

• ______________(run left to right)

• __ periods

• Period # = ______________________________

Horizontal rows

the # of energy levels/electron shells7

Example:

A) What period are potassium and bromine in? ____

B) Based on the period, how many principal energy

levels do potassium and bromine have? ____

• Elements in the same period

____________________________

o Properties change as you move across the table

4

4

do not have similar properties

B. Groups (or Families)

• _______________(up and down)

• ___ groups

• _____________________________________

___________________

o Why? Because elements in the ___________

have the ___________________

*groups 3-12 are exceptions*

Gro

ups

For groups 1 & 2: group

# = # of valence

electrons

For groups 13-18: last

digit of group # = # of

valence electrons

Vertical Columns

18

Elements in the same group have similar

chemical propertiessame group

same # of valence e-

Example:

A)What group are magnesium and calcium in? __

B) Based on the group, how do the chemical

properties of magnesium and calcium compare?

__________because they both have

__________________

• __________________________________________

2

similar2 valence e-

Valence electrons control chemical properties

1) Complete the chart below to observe the trends of elements in the same

group/period

2) Which sequence of atomic numbers represents elements which have

similar chemical properties?

A) 19, 23, 30, 36 C) 9, 16, 33, 50

B) 3, 12, 21, 40 D) 4, 12, 38, 88

3) Which two elements have the most similar chemical properties?

A) Aluminum and Barium C) Nickel and Phosphorous

B) Chlorine and Sulfur D) Sodium and Potassium

Practice

Element

symbol

Element

name

Electron

configura

tion

Period

#

Number of

electron

shells

Group

#

# of valence

electrons

Lewis

Dot

Diagram

K

Ca

Sr

IV. Metals, Nonmetals, Metalloids Besides classifying elements in periods and groups, they can

also be classified as metals, nonmetals, or metalloids.

A. Metals

• Make up _____ of the table

• _____ of or ________ “staircase”

o Exception = hydrogen is a nonmetal

• Alloy=_____________________________________

(ie brass = alloy of copper and zinc)

The “staircase”=the

division

mostLeft below

Mixture of metals by melting them together

Properties of Metals:1. Atoms that _______and _________________________when

bonding in order to become stable (look at the “selected

oxidation states” and # of valence e-)

2. ________________

3. ________________of heat and electricity

4. Have _____________________

5. ________ (can be hammered into thin sheets)

6. ________ (can be pulled into thin wires)

7. High densities

8. All ______________________________which is a liquid at

room temp.

*Note STP = Standard Temperature and Pressure = “normal,

room conditions”

lose e- form positive ions (cations)

Low electronegativity

Good conductors

metallic luster (shine)

Malleable

Ductile

solids at STP except Mercury (Hg)

Practice

1) Which element is malleable and can conduct electricity in the

solid phase?

A) Iodine B) Phosphorous C) Sulfur D) Tin

2) What happens to metals when they bond or form ions? (look

at “selected oxidation states”)

3) Why would metals want to lose electrons?

Lose electrons

Quickest way for their outermost shell to be

full= quickest way to become stable

Graphite is a type of carbon (a

nonmetal) used in pencils. It is brittle,

doesn’t conduct, and has a low density.

B. Nonmetals

Properties of Nonmetals:

3. _________________ of heat and electricity

5. ______ – break easily

2. _____________________

6. Low densities

• _______ of or ______ “staircase”

o Remember: _____________________

1. Atoms that ________ and _________________________ when

bonding in order to become stable (look at “selected oxidation

states”)

7. Mostly _______ and _______ @ STP—except ____

4. Lack luster - ______

Right above

hydrogen = nonmetal

gain e- form negative ions (anions)

High electronegativity

Poor conductors

dull

Brittle

gases solids Br(l)

Practice

1) Which element is brittle and not a conductor of electricity?

A) Sulfur B) Sodium C) Potassium D) Argon

2) What happens to nonmetals when they bond or form ions?

(look at “selected oxidation states”)

3) Why would nonmetals want to gain electrons?

Gain electrons

Quickest way for their outermost shell to become

full= quickest way to become stable

C. Metalloids

1. Atoms that gain or lose e- and form ions when bonding to become

stable (look at their oxidation states)

2. Have _____________________________________

o Typically metalloids can conduct electricity (semi-conductors)

and are hard like metals but are brittle like nonmetals

o Can be shiny or dull

• Sit ________ staircase (______________ and __________)

o Exceptions: Aluminum and Polonium= _______

• Also called ______________

Properties of Metalloids:

along between metals nonmetals

metal

semi-metals

properties of both metals and nonmetals

Practice

1. Which element has both metallic and nonmetallic

properties? A) Rb B) Rn C) Si D) Sr

2. Which list of elements contains 2 metalloids?

A) Si, Ge, Po, Sb C) As, Bi, Br, Kr

B) Si, P, S, Cl D) Po, C, I, Xe

• Electronegativity = measure of atom’s attraction for electrons

when bonded - ________________________

• The _____ atoms ________, the _______________________

• Electronegativities of the elements listed in Table ___

Periodic Trend

• Electronegativity ___________ as you move

__________________________

o Reason:

increasesleft to right across a period

S

desire to gain electrons

more want e- higher the electronegativity

A. Electronegativity:

V. Periodic Trends

Nuclear charge ____________________and _____________,

so more pull for e- from positive nucleus—

__________________

Proof: Pick 2 elements in the same period and

compare their electronegativity values

Example: Na and Cl

easier to gain electrons

size decreases(# of protons) increases

• Electronegativity ______________ as you move

_________________

o Reason:

_________________________ so outside electrons

more shielded from pull of the nucleus –

______________________

Proof: Pick 2 elements in the same group and

compare their electronegativity values

Example: Be and Ba

decreasesdown a group

Size increases (more shells)

harder to gain electrons

Practice

1) Electronegativities range from 0 – 4. Using Table S, which

element is the most electronegative?

2) What are the electronegativities for group 18 elements?

Why do you think they have “this” electronegativity?

3) Which element attracts electrons the most when bonding?

A) Ca C) Br

B) Sr D) I

4) Which element gives off its valence electrons the easiest?

A) Ca C) Br

B) Sr D) I

They do not have electronegativity numbers.

Fluorine

Because already stable so have no desire to

gain more electrons

5) Look back at the properties of metals and nonmetals.

Metals have a ______ electronegativity and nonmetals

have a _______ electronegativity. Using what you just

learned about electronegativity and what you know about

stability, explain these properties.

lowhigh

Metals:

Nonmetals:

don’t want to gain e- so they

have a low E.N.

do want to gain e- so they

have a high E.N.

B. Reactivity

• Metals: Reactivity ___________ as you move _____________;

___________is the most reactive metal

o Reason:

• Nonmetals: Reactivity __________ as you move _____________;

__________is the most reactive nonmetal

o Reason:

• Reactivity = ability or tendency of an element to go through a

chemical change – _______________

• The _______ an atom is ___________________, the

_____________it is (the more eager it is to change)

ability to react

to becoming stablemore reactive

closer

Periodic Trend

increases down a group

Francium

metals want to lose e- and move down a group, outermost e-

are farther from the nucleus so easier to lose

decreases down a group

Fluorine

nonmetals want to gain e- and as move down a group, the

pull from the nucleus is more shielded so harder to gain e-

VI. Specific Groups/Families

A. Group 1: ________________Alkali Metals

Brainiac – Alkali Metals

• All have __________________

• Easily lose their one electron to become ___

ions (look at oxidation state). Why?

• ____________ reactive _____________

______________

• Contains most reactive metal, Francium, but

it’s so rare so some sources say it’s

___________

Quickest way for them to become stable

1 valence electron

+1

EXTREMELY never found

alone in nature

Cesium (Cs)

• Have __ valence electrons

• Form ions with a ___charge

B. Group 2: ______________________Alkaline Earth Metals

• _______ reactive never found alone in nature

2

+2

Fairly

Practice

1. What will be the charge on a Ca ion when it bonds?

2. What does that mean in terms of electrons?

3. Why does calcium want that?

+2

By losing 2 electrons, calcium gets 8 electrons in it’s

outermost occupied shell so it becomes stable!

It loses 2 electrons

• Form ________ ions (see oxidation states)

• _____ reactive than Groups 1 and 2 therefore _________

________________________________

• Dense metals with high melting points (see table S)

C. Groups 3-12: ________________Transition Metals

• Form ________________________________(ex. Cu is

bright blue when dissolved in water)

multipleless can be

found uncombined or alone in nature

colored ions/aqueous solutions

Practice

1. Identify the metal that has multiple oxidation states.

A) K B) Ba C) Be D) Pd

2. Which compound forms a green aqueous solution?

A) RbCl B) CaCl2 C) NiCl2 D) BeCl2

3. Which set of properties is most characteristic of transition

elements?

A) Colorless ions in solutions, multiple positive oxidation states

B) Colorless ions in solutions, multiple negative oxidation states

C) Colored ions in solutions, multiple negative oxidation states

D) Colored ions in solutions, multiple positive oxidation states

• Have ___ valence electrons

• Form ions with a ____charge. Why?

• Mostly ___________ elements

D. Group 17: ________________Halogens

• Form _________________ called _________

• Contains the most reactive nonmetal, Fluorine

• All three states of matter found in group: ________________

_____

7

-1

nonmetallic

salts/compounds halides

solid (s), liquid (l),

gas (g)

• Have __valence electrons

o Exception: Helium (He) which only has ___

• Do not want to form ions or bond with anything (look at

oxidation states). Why?

E. Group 18: ________________Noble Gases

• ________________because outermost occupied shell is full =

__________________

o Exception: ___________ which only has 2 e- in outermost

shell but is still considered stable. Why?

• All other elements want to be like them –a lot of chemistry is

explained by this fact! 8 is great!

• ___________ or _____ therefore stable enough to

_____________________

Because the first shell is full with 2 e-

8

2

Most stable group octet (8 valence e-)

Helium (He)

Unreactive inert

exist alone in nature

• Has __ valence e- so grouped with alkali metals

• Very reactive

• a _________ and a _____

F. ______________ not officially part of a groupHydrogen

1

nonmetal gas

VII. Other Categories

A. Diatomics: Molecule containing __________ atoms

• What are they? (remember these)

H, O, F, Br, I, N, Cl (7-up)

B. Radioactive: atoms where nucleus is unstable and breaks

down spontaneously to form a new, more stable element

• What are they?

Atomic # greater than 83

• They pair up in order to fill valence shells and become stable

2 identical

C. Synthetic: __________ elements (all radioactive as

well, but disintegrate in milliseconds)

• What are they?

Atomic # greater than 93 and 43 (Technitium),

61 (Promethium)

man-made

D. Allotropes: 2 or more ______________________

of the _______________ in the ______________

• _________________ = ___________________

• The 2 most common examples:

1. Carbon(s) = Diamond, graphite, coal, and Buckminster

Diamond Graphite Coal Buckyball

2. Oxygen(g) = O2 (what we breath) vs. O3 – Ozone (Toxic)

different structural forms same element same phase

Different structures different properties