Transcript Ch 7 Rimkas Group

Tadas Rimkus
Krishna Trehan
Rachel Won
Soo Jeon
What is the SWBAT????
 Objective
 Know a brief history of the periodic table
 Know the effective nuclear charge
 Examine periodic tends in the atomic size, ionization
energy, and electron affinity
 Examine the sizes of ions and their electron
configurations
 Explore some of the difference in the physical and
chemical properties of metals, nonmetals, and
metalloids.
 Discuss some periodic trends.
7.1 The Periodic Table
 Mendeleev, Meyer, Moseley strived to investigate the





possibilities of classifying elements in useful ways.
Mendeleev and Meyer - developed the periodic table on the
basis of the similarity in chemical and physical properties.
He predicted Ga, Ge, and Sc.
Moseley- established that each element has a unique atom
number.
- found out that each element produces X rays of a
unique frequency; the frequency increases as the atomic
mass increases
Rutherford- proposed the nuclear model of the atom.
Element in the same column of the periodic table have the
same number of electrons
Periodic Table
7.2 Sizes of Atoms and Ions
 Electron shells in atoms
- As we move down a column of the P.T, we change the
principal quantum number, n, of the valence orbital of
the atoms.
-All the orbitals with the same value of n
= shell
 Before Bohr, Gilbert N. Lewis had suggested that
electrons in atom arranged in spherical shells around
the nucleus.
Distance from nuclueus
 The quantity plotted on the
vertical axis is called radial
electron density.
 Red line = Ne
 Blue line = He
 The reason that 1s shell in Ne so
much closer to the nucleus than
the shell in He is..
nuclear charge of Ne is +10, but
He is +2.
 Because the 1s electrons are the
innermost electrons of the atom,
they are not shielded from the
nucleus very effectively by other
electrons.
Atomic sizes
 Ionic size – the size of an ion plays an important role in
determining the structure and stability of ionic solids.
 Periodic trend – metals tend to lose electrons. The size
of the atom becomes smaller with the loss of each.
 Bonding atomic radius – is based on the distances
separating atoms when they are chemically bonded to
one another. The radius of the atom is determined by
the sizes of the orbitals occupied by the outermost
electrons.
Sample Problem
 Arrange these atoms and ions in order of decreasing
size:
Si, Si+2 and C+2.
• Solution- Cations are smaller than their parent atoms.
So,
The Si+2 ion is smaller than the Si atom. Because Si is
below C in group of 4A of the periodic table, Si is
larger than C+2. ( Si > Si+2 > C )
The distinction between non-bonding and
bonding atomic radius
Bonding
atomic
radius
Non-bonding
atomic radius
7.3 Ionization Energy
 the amount of energy need to remove an electron from
a specific atom or ion in its ground state
 1st Ionization Energy
 energy needed to remove the 1st electron from an atom
Ionization Energy… CONTINUE
 ACROSS- increase
 As you go across the periodic table, the electrons
are closer to the nucleus increasing the energy
necessary to remove an electron
 DOWN- decrease
 As you go down the periodic table, the electrons
are farther from the nucleus
• The ionization energies of
the transition elements
increase slowly as we proceed
from left to right in a period.
• Alkali metals have the lowest
ionization energy in each row.
•Noble gases are the highest.
Sample problem
• of the following, which has the highest and lowest
first
ionization energy?
C and Al
 Answer: Highest- C
Lowest- AL
7.4 Electron Affinities
 Ionization Energy: energy required to remove an
valence electron
 Cl(g)  Cl+(g) + e- ΔE = 1251 kJ/mol
 See how E is positive?
 This is the energy which you must add, in order to
remove an electron from Chlorine
Electron Affinity
 Did you know that atoms can gain electrons to gain a
negative charge?!
 This energy change that occurs is called ELECTRON
AFFINITY because it measures the ATTRACTION
between the newly added electron and the atom
 It shows how much the atom WANTS the electron!
Electron Affinity
 So……. FLUORINE
wants an electron
MORE than
LITHIIUM, meaning it
would have a LARGER
electron affinity!
http://www.chemguide.co.uk/inorganic/group7/eadiag.gif
Continue?
 The elements on the right, are one electron away from
GETTING A FULL OCTET!
 They will have a high affinity, so they can gain the last
electron.
 (The more negative the electron affinity, the greater the
attraction for an electron)
Continue!







Nuclear Charge
the change in the nucleus or the number of protons
ACROSS- increase
DOWN- increase
Atomic Radius
one half the distance from center to center of like atom down-increases
As you go down the periodic table, a new energy level is added
increasing the size of the atoms
 Ex) of the following, which has the highest and lowest first ionization
energy?

Answer: Highest- C

Lowest- AL
Periodic Trends Of Electron Affinity
 Electron Affinity tends to decrease as we go from left to
right!
 BUT WHY?!?!?!?!?!?!?!?!?!?!?!
 This is because the elements on LEFT side of the
Periodic Table of LESS valence electrons.
7.5 metals, nonmetals, and metalloids
 Metals are the shiny things we have learned to love
and treasure
 They are ductile, conduct heat and electricity, and
malleable.
 They are all solid (except Mercury!) at room
temperature.
 They can have high melting points, for example
Chromium has a melting point at 1900 C.
Metals
 Metals have low ionization energies (energy required
to remove an valence electron).
 WHICH MEANS, they tend to form positive ions
EASILY!!!!!!!
Continue…
 Periodic Trend - Nonmetals tend to gain electrons. The
size of the atom becomes larger with the addition of
each consecutive electron. The nuclear charge has less
of a pull on the electrons with the gain of each electron
due to greater electron repulsion!