Download Ch 5

Ch 5
Sec 5-1
Around 1850, more than 60 elements had been discovered.
Chemists had to learn the properties of these elements as well
as the compounds they formed, which was very difficult.
At that time, there was no accurate way to determine the atomic
mass or the # of atoms of an element in a particular compound.
Different chemists used different atomic masses for the same
element, which made it impossible to compare results.
In 1860, a group of chemists met at the 1st International Congress of
Chemists in Germany.
•
They wanted to settle the issue of atomic mass and other matters.
•
Italian Chemist – Stanislao Cannizzaro – presented a convincing
method for accurately measuring the relative mass of an atom.
•
They agreed on his method as a standard and initiated a search for
relationships between the atomic mass and other properties of the
elements.
Dmitri Mendeleev (Russian teacher) heard about the new method
to calculate the atomic masses, he decided to include the new values
in a chemistry textbook he was writing.
 Arranged note cards with the name of the element, atomic mass,
physical and chemical properties.
 Started to arrange the note cards in various ways – he looked for
trends.
 Organized them in increasing atomic mass – then the phy. & chem.
properties repeated periodically.
(See pg 124 for the 1st periodic table.)
** Mendeleev left holes in the table but he predicted the existence &
properties of 3 elements that would fit in the spaces. Therefore, he is
credited with the periodic law.
A few questions remained:
1. Why could most of the elements be arranged in the order of
increasing atomic mass, but a few could not?
2.
What was the reason for the chemical periodicity?
The first question wasn’t answered until about 40 years later after the
first published table.
Henry Moseley (1887 – 1915) – worked with Rutherford.
• He examined the spectra of 38 different metals.
• Discovered a previously unrecognized pattern.
• The elements fit better if they were arranged by the # of p+.
• This led to the modern definition of the atomic mass and that the
periodic table should be based on the atomic # not the atomic mass.
Periodic Law – The physical and chemical properties of the elements
are periodic functions of their atomic #.
The periodic table continues to undergo slight changes, but it is now
arranged so that elements with similar properties fall in the same
column.
Sec 5-2
In the periodic table, we have special names for different sections.
1.
2.
3.
4.
5.
6.
7.
8.
9.
Alkali Metals – group 1 (except Hydrogen)
Alkaline – Earth – group 2
Transition Metals – groups 3 – 12
Main group elements – groups 13 – 18
Halogens – group 17
Noble gases – group 18
Lanthanides
Actinides
Metalloids
Alkali Metals –
• Silvery appearance
• Soft – could cut with a knife
• So reactive they are not found as free elements in nature
• React violently with most non-metals
• React violently with water to produce hydrogen gas
• Usually stored in kerosene
Alkaline – earth metals –
• Harder, denser, stronger, and have higher melting points than the Alkali
metals.
• Less reactive than the Alkali metals, but are still reactive enough that
they are not found free in nature.
Transition Metals
•As all metals, they are good conductors of electricity and heat. They also
have a high luster.
•They are typically less reactive than the alkali and alkali-earth metals.
Therefore some are found in nature as pure elements. Gold for example.
 Main group elements
 Properties of the elements vary greatly.
 Includes all of the non-metals except H.
 Contain all the metalloids
 Halogens
 Most reactive non-metals
 React vigorously with most metals to form “salts”.
 Gases at room temperature
 Noble Gases
 Unreactive because they contain 8 outershell e-.
Section 5 - 3
Electron Configuration and Periodic Properties
Atomic Radii
 One way to express an atom’s radius is to measure the
distance between two identical atoms that are chemically
bonded together, then divide the distance by two.
 Atomic Radius— ½ the distance between the nuclei of
identical atoms that are bonded together
 What would the atomic radius be of the Chlorine atom
below:
198 pm
Periodic Trends
 As you travel from left to right across the periodic table, the
atomic radii decreases
 Caused by the increased pull of the e- by the nucleus
Group Trends
 As your travel from the top of the periodic table to the bottom,
the atomic radii increases
Periodic Table of Atomic Radii
Decrease in Atomic Radii
Increase in Atomic Radii
Examples
Ex 1: Of the elements magnesium, Mg, chlorine, Cl, sodium,
Na, and phosphorus, P, which has the largest atomic
radius? Explain.
Ex 2: Of the elements calcium, Ca, beryllium, Be, barium, Ba,
and strontium, Sr, which has the largest atomic radius?
Explain.
Ionization Energy
 Ionization Energy—the amount of energy needed to
remove an e Can be expressed as:
A + energy  A+ + eAtom
Ion with a 1+ charge
 Ex:
Na + energy  Na+ + eCa + energy  Ca2+ + e-
Period Trends
 Group 1 metals have the lowest ionization energythey lose
e- the easiest
 Group 18 noble gases have a very high ionization energy
the do not lose e- very easily
 Increase across a period is caused by the increase in nuclear
charge
 In general, nonmetals have a higher ionization energy than
metals
Group Trends
 As you travel down the periodic table the e- in the outer shell
get further away from the nucleus
 Can be removed much more easily
 Valence e- can very easily be taken off of the outer shell.
However, after that it takes an immense amount of energy
Examples
 Consider two main-group elements A and B. Element A has
a first ionization energy of 419 kJ/mol. Element B has a first
ionization energy of 1000 kJ/mol. For each element, decide
if it is more likely to be in the s block or p block. Which
element is more likely to form a positive ion?
Examples
 Consider the four hypothetical main-group elements Q, R, T,
and X with the outer electron configurations indication below.
Then answer the following questions:
Q: 3s2 3p5
a)
b)
c)
d)
R: 3s1
T: 5s2 4d10 5p5
X: 5s2 4d10 5p1
Identify the block location of each main-group element.
Which of these elements are in the same period? Which are in the
same group?
Which elements would you expect to have the highest first
ionization energy? Which would you expect to have the lowest?
Which of the elements is most likely to for a 1+ ion?
 Metals (on the left) need to lose e- to form an octect, or
eight e- in their outer shell.
 Makes them (+) charged—Called a cation
 Ex: Li+1, Na+1, and K+1 all need to lose 1 e- to be happy
 Ex: Be2+, Mg2+ and Ca2+ all need to lose 2 e- to be happy
 Metals are such good conductors because they want to lose e-
so easily
 Nonmetals don’t want to lose any e-, they would rather gain.
Electron Affinity
 Electron Affinity—the energy that occurs when an e- is
added to an atom.
 Adding an e- to an atom would make the atom (-) charged—
this is called an anion
 Ex: F-1, Cl-1, and Br-1 all need to gain 1 e- to be happy
 Ex: O-2, S-2, and Se-2 all need to gain 2 e Cations, like Na+1 will combing with anions, like Cl-1 so both
are happy
 Will form NaCl (table salt)
Ionic Radii
 Cation—Positive ion
 Formed by the loss of one or more e Results in the decrease in atomic radius
 Anion—Negative ion
 Formed by the addition of one or more e Results in the increase in atomic radius
Period Trends—Ionic Radii
 Metals on left form cations
 Radii decreases across a period b/c e- cloud shrinks
 Nonmetals at upper right form anions
Group Trends—Ionic Radii
 Gradual increase in Ionic Radii as you travel down a group
Valence Electrons
 The electrons that interact to form chemical compounds are
those in the highest energy levels
 Valence Electrons—The electrons available to be lost,
gained, or shared in the formation of chemical compounds
 Located in incompletely filled main-energy levels
 Ex: When Na loses an electron it becomes Na+. The e- that is
lost is known as a valence e Occupy the s and p orbitals
Examples
How many valence e- does the following have:
1)
2)
3)
4)
5)
6)
Mg
Cs
C
O
Ar
Al
Electronegativity
 Valence e- hold chemical compounds together
 The negative charge is concentrated closer to one atom than
another
 Linus Pauling—devised an electronegativity scale
 Electronegativity—a measure of the ability if an atom in a
chemical compound to attract electrons
 Most electronegative element is Fluorine Given an
electronegativity value of 4
Period Trends
 Tend to increase across each period
 Alkali and Alkaline-Earth Metals are the least electronegative
 Nitrogen, Oxygen, and halogens are the most electronegative
 Atoms attract e- strongly in compounds
Group Trends
• Tend to decrease down a group
• Noble Gases are unique in that some are unreactive and have no
electronegativity
Example
Among the elements Ga, Br, and Ca which has the highest
electronegativity. Why?