Transcript BOOKLETColoring-the-Periodic-Table-Familiesx

 Quick overview of Periodic Table
use 1st (9:00)
 Periodic Table (29:00, start at 19:00)
Early Tries at Organizing Elements
 Up to the 1800’ s scientists were busy trying to discover
new elements
 Not much thought was given to how they should be
organized, so elements were just listed alphabetically.
 Some scientists thought the elements might be
classified by color.
 Some thought by taste but then found that many were
poisonous.
 Because scientists liked to measure things, they were
able to measure the atomic weight.
 ALONG COMES MENDELEEV
 He placed the name of an element on a piece of paper
along with its known properties.
 Then he began to arrange the cards, keeping in mind
that atomic mass and properties were very important.
 By 1871 he had all elements arranged and came up with
a law.
If elements are arranged
according to their atomic mass a
pattern can be seen in which
similar properties occur
regularly.
 Sometimes when arranging he would come
to a place in which no element would fit. So
he decided to leave it blank thinking that
this might be a place for an element that
had yet to be discovered.
 He could also predict the properties of the
new element by looking at the properties of
the known elements that were above and
below the unknown element.
Families on the Periodic Table
 Elements on the periodic table can be grouped into
families based on their chemical properties.
 Each family has a specific name to differentiate it
from the other families in the periodic table.
 Elements in each family
react differently with
other elements.
 The elements in a group have the same number of
electrons in their outer orbit.
 They tend to form ions by gaining or losing the same
number of electrons
Important facts
 Arrangements of elements in order of increasing
atomic number such that the elements show related
chemical properties.
 The periodic table arranges elements by:
 -vertical columns called groups- group number tells
the number of electrons in the outermost shell,
(ie. the number of valence electrons). There are 18
groups.
 -horizontal rows called periods- period tells the
number of shells or energy level. There are 7 periods.
ALKALI METALS
Group 1
 Hydrogen is not a member, it is a
non-metal
 1 electron in the outer shell
 Soft and silvery metals
 Extremely reactive, esp. with water
because outer orbital shell has only
one electron (unstable arrangement)
 Found in nature only as compounds
 Conduct electricity
HYDROGEN
 Hydrogen is a unique
element.
 Like the alkali metals
it has only one
electron in its outer
orbit.
 It has little else in
common with the
alkali metals: it is a
colourless , odourless,
tasteless highly
flammable gas.
ALKALINE EARTH METALS
Group 2
 2 electrons in the
outer shell
 White and
malleable
 Reactive, but less
than Alkali metals
 Conduct electricity
TRANSITION METALS
Groups in the middle
 Good conductors of heat
and electricity.
 Some are used for
jewelry.
 The transition metals are
able to put up to 32
electrons in their second
to last shell.
 Can bond with many
elements in a variety of
shapes.
BORON FAMILY
Group 3
 3 electrons in the
outer shell
 Most are metals
 Boron is a
metalloid
CARBON FAMILY
Group 4
 4 electrons in the
outer shell
 Contains metals,
metalloids, and a
non-metal Carbon
(C)
NITROGEN FAMILY
Group 5
 5 electrons in the
outer shell
 Can share electrons
to form compounds
 Contains metals,
metalloids, and
non-metals
OXYGEN FAMILY
Group 6
 6 electrons in the
outer shell
 Contains metals,
metalloids, and
non-metals
 Reactive
HALOGENS
Group 7
 7 electrons in the outer
shell
 All are non-metals
 Most reactive
 Are often bonded with
elements from Group 1
Noble Gases
Group 8
 Exist as gases
 Non-metals
 8 electrons in the outer shell
= Full
 Helium (He) has only 2
electrons in the outer shell =
Full
 Not reactive with other
elements,(inert) they almost
never form chemical
compounds with other
elements
 Con’t
 The lack of reactivity of noble gases is explained by
their electronic structure.
 Their outer orbits are filled.
 They do not react with other atoms because they
already have a stable arrangement of electrons.
The 3 Sections:
Metals, Non-metals, and Metaloids ( the elements: chemistry
lesson 25:00, stop at 19:00)
Metals
Non-Metals
Metalloids
Characteristics of Metals
 - Metals are good conductor of heat and electricity.
 - Metals are shinny and ductile (stretchable)
 - Metals are malleable (can be pounded into thin
sheets.)
 - A chemical property of metal is its reaction with
water which results in corrosion.
Characteristics of Non-metals
 - Non-metals are poor conductors of heat and
electricity.
 - Non-metals are not ductile or malleable.
 - Solid non-metals are brittle and break easily.
 - They are dull.
 - Many non-metals are gases.
Characteristics of Metalloids
 Are elements that possess both metallic and
nonmetallic properties.
 Not strictly a group themselves, they are found on the
right side of the periodic table, on both sides of the
zigzag line that divides the metals and the nonmetals.
 - Metalloids have properties of both metals and non-
metals.
 - They are solids that can be shinny or dull.
 - They conduct heat and electricity better than nonmetals but not as well as metals.
 - They are ductile and malleable.
Rare Earth Metals
 Some are
Radioactive
 The rare
earths are
silver, silverywhite, or gray
metals.
 Conduct
electricity
Rows on the Periodic Table
 The groups of elements- the columns in the periodic
table -have similar physical and chemical properties.
 These horizontal rows of elements are called periods.
 The first period contains two elements: hydrogen and
helium.
 The second period contains eight elements, starting
with lithium and ending with neon.
 As you go from left to right within a row, the atomic
number increases and the elements gradually change
from metallic to nonmetallic.
Atomic radius
 Atomic radius – size of an atom
 Increases down a group – more shells, electrons
added further from nucleus
Ionic radius
 Increases down a group
 Decreases across a period
 Cations (positively-charged ions)
 Cations are smaller than their parent atom because the
same number of protons in the nucleus pulls on less
electrons.
 Metals commonly become cations.
 Anions (negative ions)
 Anions are larger than their parent atom because the
same number of protons in the nucleus pulls on more
electrons.
Ionic Radius / Atomic Radius
 atomic radius is the radius of the atom when it's in
neutral.
ionic radius is the radius of the atom when it losses or
gains electron ( become an ion)
So, for the very same element the number of electrons
in atom differs than the number of electrons of the ion
which might cause a change in the number of orbitals
and the radios.
Reactivity
 Metals
 Period - reactivity decreases as you go from left to right
across a period.
Group - reactivity increases as you go down a group
 Why? The farther to the left and down the periodic chart
you go, the easier it is for electrons to be given or taken
away, resulting in higher reactivity.
 Non-metals
 Period - reactivity increases as you go from the left to
the right across a period.
Group - reactivity decreases as you go down the group.
 Why? The farther right and up you go on the periodic
table, the higher the electronegativity, resulting in a
more vigorous exchange of electron.
Video
 Periodic Table
1st video
 Formulas Lesson 1: Writing Formulas For Binary Ionic
Compounds - YouTube
Space Jam
Space station tour
Periodic Table
 [1] First side of your sheet, color and label the groups.
List two things you know about the groups. (10)
 Alkali Metals, Alkaline Earth Metals, Hydrogen, Transition
metals, Boron, Carbon, Nitrogen, Oxygen, Halogen, Noble
gases
 [2] Number the periods. List two things you know
about the periods.
 [1] Second side of your sheet, neatly shade the areas
of the table that contain metals, non-metals, and
metalloids.
 [2] Label each.
 [3] List the characteristics of each on the sheet
provided