periodic table and periodicity - belton independent … · moseley’s contribution henry moseley...
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Periodic Table and
Periodicity
BHS Chemistry 2010- 2011
In 1869, Dmitri Mendeleev, a Russian chemist noticed patterns in certain elements.
He discovered a way to arrange the elements so that they were organized by their chemical and physical properties.
Moseley’s Contribution
Henry Moseley is
credited for further
arranging the elements
on the periodic table in
order of the number of
protons they contained.
Circle Periodic Table
Attention:
New Additions to Periodic Table
WOMANIUM (Wo)
Physical properties: Boils at nothing and may freeze any time.
Chemical properties: Very active and highly unstable. Possesses
strong affinity with gold, silver, platinum, and precious stones. Violent
when left alone. Turns slightly green when placed next to a better
specimen.
Usage: An extremely good catalyst for dispersion of wealth.
MANIUM (Xy)
Physical properties: Solid at room temperature but gets bent out of
shape easily.
Chemical properties: Becomes explosive when mixed with Childrium
for prolonged period of time.
Usage: Possibly good methane source.
Caution: In the absence of WO, this element rapidly decomposes and
begins to smell.
Periods
The table is arranged in horizontal
rows (going across) called periods.
There are 7 periods.
The period tells you how many
electron energy levels the atom
has.
1
5 4
3
2
7
6
Group
The table is also arranged in vertical
columns (going down) called groups.
There are 18 groups.
Members of each group have similar
physical and chemical properties.
1
2
3
2
4
2
5 6
2
7 8 9 10 11 12
2
13 14 15 16
18
17
III. Properties of Metals, Nonmetals and Metalloids
A. Properties of Metals
1. They are malleable, and have luster.
2. They reflect heat and light
3. Good conductors of heat and electricity.
4. Typically solids at STP
5. High melting points
6. They lose electrons in chemical reactions to become cations (+)
B. Properties of Nonmetals
1. They are dull and brittle.
2. Don’t conduct heat and electricity well.
3. Low boiling and freezing points.
4. They exist in all three phases at STP, but most are gases.
5. They gain electrons in chemical reactions to become anions (negative ions)
C. Properties of Metalloids
(B, Si, Ge, As, Sb, Te)
1. They possess
intermediate properties
between metals and
nonmetals.
2. They are
semiconductors at
higher than room
temperatures.
3. They are all solids at
STP.
IV. Special Groups
A. Alkali Metals – Group 1 (IA)
on the periodic table.
1. They are soft and easily
cut.
2. They are highly reactive. (especially with H2O)
Highly Reactive Video
3. All have an electron
configuration ending in s1
4. Gives up 1 electron in
bonding (+1)
5. Has 1 valence electron
B. Alkaline Earth Metals – Group 2 (IIA) on the periodic table.
1. They are less reactive than the alkali metals.
2. They have an electron configuration ending in s2
3. Gives up 2 electrons to form a (2+) charge
4. Has 2 valence electrons
C. Halogens – Group 17 (VIIA) on the periodic table.
1. They are highly reactive and react violently with hot metals.
2. They form diatomic molecules
3. They have an electron configuration ending in s2p5
4. Accepts 1 electron to have a (-1) charge
5. Has 7 valence electrons
D. Noble Gases – Group 18 (VIIIA) on the periodic table.
1. Extremely stable and unreactive
2. They exist as single atoms
3. They have an electron configuration ending in s2p6
4. Has 8 valence electrons
E. Transition Metals – Groups 3 - 12 (B groups) on the periodic table.
1. They possess characteristics of active metals to varying degrees.
2. They form compounds that are usually brightly colored.
F. Inner Transition Metals – two rows at the bottom of the periodic table.
1. Many of the inner transition metals are radioactive.
2. Many of the actinides are synthetic.
“The Periodic Table”
The Periodic Law says: When elements are arranged in order
of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Horizontal rows = periods
Vertical column = group (or family)
• Similar physical & chemical prop.
ALL Periodic Table Trends Influenced by three factors:
1. Energy Level
• Higher energy levels are further
away from the nucleus.
2. Charge on nucleus (# protons)
• More charge pulls electrons in
closer. (+ and – attract each other)
3. Shielding effect
Shielding The electron on the outermost
energy level has to look through all
the other energy levels to see the
nucleus.
This effect decreases
the attraction of the
nucleus for the outer
electrons..
What do they influence?
Energy levels and Shielding have an
effect on the GROUP ( )
Nuclear charge has an effect
on a PERIOD ( )
Atomic Size
Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.
} Radius
#1. Atomic Size - Group trends
As we increase
the atomic
number (or go
down a group). . .
each atom has
another energy
level,
so the atoms get
bigger.
H
Li
Na
K
Rb
#1. Atomic Size - Period Trends
Going from left to right across a period,
the size gets smaller.
Electrons are in the same energy level.
But, there is more nuclear charge.
Outermost electrons are pulled closer.
Na Mg Al Si P S Cl Ar
Ions
Some compounds are composed of
particles called “ions”
• An ion is an atom (or group of atoms)
that has a positive or negative charge
Atoms are neutral because the number
of protons equals electrons
• Positive and negative ions are formed
when electrons are transferred (lost or
gained) between atoms
Ions Metals tend to LOSE electrons,
from their outer energy level
Nonmetals tend to GAIN one or
more electrons
Ions
Here is a simple way to remember
which is the cation and which the
anion:
This is a cat-ion. This is Anion.
She’s unhappy and
negative.
+ +
Ionic Radius
The size of an ion
Cations are smaller (lost e-) and
anions are larger than the atoms they
cam from.
Ionic radius Group trends
Each step down a
group is adding an
energy level
Ions therefore get
bigger as you go
down, because of
the additional
energy level.
Li1+
Na1+
K1+
Rb1+
Cs1+
Ionic radius Period Trends Across the period from left to
right, the nuclear charge
increases - so they get smaller.
Notice the energy level changes
between anions and cations.
Li1+
Be2+
B3+
C4+
N3- O2-
F1-
Trends in Ionization Energy
Ionization energy is the amount of energy required to completely remove an electron.
Ionization Energy - Group trends
As you go down a group, the first IE decreases because...
• The electron is further away from the attraction of the nucleus, and
• There is more shielding.
Ionization Energy - Period trends
IE generally increases from left
to right.
Same shielding.
But, increasing nuclear charge
The arrows indicate the trend:
Ionization INCREASE in these
directions
Trends in Electronegativity
Electronegativity is the tendency
for an atom to attract electrons to
itself.
An element with a big
electronegativity means it pulls the
electron towards itself strongly!
Electronegativity Group Trend The further down a group, the farther
the electron is away from the nucleus, plus the more electrons an atom has.
Thus, more willing to share.
Going down a group, EN decreased
Electronegativity Period Trend Metals
• They want to lose electrons
• Low electronegativity
Nonmetals.
• They want more electrons.
• Going across a period, the EN
increases
The arrows indicate the trend:
Electronegativity INCREASE in these
directions
Textbook:
Atomic Radius: pg. 141
Ionization Energy: pg. 143
Ionic Radius: pg. 149
Electronegativity: pg. 151