Chemistry
Chapter 13 Chemical Periodicity
I.
Introduction
- In the 19th century chemists began to categorize the elements according to similarities in their
physical and chemical properties
- The end result of this was the modern periodic table
- The periodic table is very useful for understanding and predicting the properties of elements
II.
13.1 Development of the Periodic Table
- Newland an English chemist published list of elements arranged in order according to their
increasing atomic mass.
- He stated that the elements properties repeated when they were arranged according to increasing
atomic mass in groups of eight
- He called this the arrangement the law of octaves
- Similar to musical scale that repeats every eighth note
- Law only works up to Ca
- Mendeleev Russian chemist refined and added to the arrange of elements in a table according to
their atomic masses
- With this arrangement he noticed a regular (periodic) recurrence of their physical and chemical
properties
- Mosely rearranged the periodic table according to the atomic number of the elements, which is
how the modern periodic table is arranged today
- The periodic table is a valuable organizational tool for chemists
3.2 The Modern Periodic Table
- The most commonly used modern periodic table, sometimes called the long form (your table)
- The long form table list many properties of the elements so that the chemist can check them at a
glance
- The horizontal rows of the periodic table are called periods – there are 7 periods in the periodic
table
- The vertical columns are called groups or families identified by number and a letter
- Groups 1A through 7A and group 0 make up the representative elements (wide variety of
properties)
- Group B elements are the transition metals
- Two rows of elements below the periodic table are the lanthanides and actinides
- The periodic law states that when elements are arranged in order of increasing atomic number,
there is a periodic pattern in their physical and chemical properties
13.3 Electron Configurations and Periodicity (Valence electrons)
- The electron configuration of an element plays the greatest part in determining it’s physical and
chemical properties
- Most elements within the same group have the identical electron configurations in their outer
most energy level (valence level)
- Elements are classified into 4 different categories according to their valence (outer) electron
configuration
1. Noble gases (group 0) are elements in which the outermost s and p sublevels are filled also called
inert gases because they do not react with other elements because they are stable – all have 8 but
He has 2 – Ne, Ar, Kr, Xe, Rn
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III.
IV.
2. Representative elements (Group A)
- Elements whose outermost s or p sublevels are only partially
filled
- Group 1A are known as the alkali metals 1 electron in outermost energy level – very reactive
(only in compounds in nature)
- Group2A are known as the alkaline earth metals 2 electrons in outermost energy level – also
reactive but not as reactive as 1A (only in compounds in nature)
- Group 7A are known as the halogens 7 electrons in outermost energy level – nonmetals also
highly reactive – also called salt formers
- For any representative element the group number is equal to the number of electrons in the
outermost energy level
- Look at periodic table 354 –355
3. Transition Metals (Group B)
- Elements whose outermost s sublevel and nearby d sublevel contain electrons – the d sublevels
overlap with s sublevels and this is why they are transition elements
- Characterized by having electrons added to the d orbital
- Not as reactive as Group A elements
4. Inner Transition Metals
- Elements whose outermost s sublevel and nearby f sublevel generally contain electrons
- Characterized by the filling of the f orbitals
- The periodic table can be divided into sections, which correspond to the sublevels that are filled
with electrons (on your table) (blocks)
- Group1A and 2A are in the s block – valence level = period #
- Group 3A, 4A, 5A, 6A, 7A, and 0 belong to p block, except He it is s block – valence level =
period #
- Transition belong to d block – exceptions d # is one less than period
- Inner transition belong to f block – exception F # 2 less than period
- The valence electron configurations can be determined by using the block diagram in figure 13.4
– on your periodic table
V.
13.4 Periodic Trends in Atomic Size
- An atom does not have a sharply defined boundary to set the limit of its size
- The radius of an atom can not be measured directly, but can be measured indirectly or estimated –
x-ray diffraction.
- Half the distance between two like atoms (diatomic) is the atomic radius
- Group trends - atomic size generally increases as you move down a group in the periodic table
- The size increases because electrons are added to higher principle energy levels – added charge of
nucleus pulls electrons inward – but the net effect is an increase in size – because electrons are
further from nucleus
- Period trends – atomic size generally decreases as you move from left to right across a period
- The size decreases because electrons are added to the same principle energy level - added charge
of nucleus pulls electrons inward – but net effect is a decrease in size – attraction is stronger
because electron is closer to nucleus
- This trend is less pronounced in periods where there are more electrons in the occupied principle
energy levels between the nucleus and the outermost electrons
- This is referred to as the shielding effect
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VI.
13.5 Periodic Trends in Ionization Energy
- When an atom gains or losses an electron it becomes an ion
- The energy that is required to overcome the attraction of the nuclear charge and remove an
electron from a gaseous atom is called the ionization energy
- The first ionization energy is the amount needed to remove the first outermost electron
- The second ionization energy is the amount needed to remove the next outermost electron
- The third ionization energy is the amount to remove the third and so on (Table 13.1 page 362)
- Ionization energies can be used to predict how many electrons an atom will gain or lose in a
chemical reaction (1A and 2A)
- Two factors affect ionization energy: nuclear charge and distance from the nucleus
- Group Trends – increase atomic number within a group produces larger atoms, because of the
electrons greater distance from the nucleus it is more easily removed
- Therefore ionization energy decreases as you move down a group
- Period trends – increase atomic number within a period produces smaller atoms, because
electrons are closer to the nucleus it is harder to remove them
- Therefore ionization energy increases as you move from left to right across the period
Symbol
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
VII.
1st IE kj/mol
1312
2371
520
900
800
1086
1402
1314
1681
2080
496
738
577
786
1012
1000
1255
1520
419
590
2nd IE kj/mol
3rd IE kj/mol
5247
7297
1757
2430
2352
2857
3391
3375
3963
4565
1450
1816
1577
1896
2260
2297
2665
3069
1146
11810
14840
3659
4619
4577
5301
6045
6276
6912
7732
2744
3229
2910
3380
3850
3947
4600
4941
13.6 Periodic Trends in Ionic Size
- Positive ions (cations) are always smaller than the atoms from which they are formed
- Loss of outer electron increases the attraction of the nucleus for the few remaining electrons
- Negative ions (anions) are always larger than the atoms from which they are formed
- The effective nuclear attraction is less for an increased number of electrons
- Periodic relationship: going left to right across a period there is a gradual decrease in size for
positive ions and negative ions
- The ionic radius of both anions and cations increases as you go down each group
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VIII.
13.7 Periodic Trends in Electronegativity
- The electronegativity of an element is the tendency for the atoms of the element to attract
electrons when they are chemically combined with another element
- The Pauling scale uses arbitrary units to express the electronegativity of the all elements (except
noble gases)
- The Pauling scale is based on ionization energies and electron affinities
- In general a high electronegativity number means that an element has a tendency to attract
electrons when forming chemical bonds
- F – 3.98 is the most electronegative element, it always attracts electrons and forms a negative ion
- Fr - .70 is the least electronegative element, it loses electrons and forms positive ions
- Table 13.2 page 366
- As you go across a period from left to right, the electronegativity of the representative elements
increases
- Metallic elements far left have low electronegativities
- Nonmetallic elements far right have high electronegativities
- Electronegativities generally decrease as you go down a group
- Transition metals do not show as regular trend of electronegativities
IX.
Electronegativity and Chemical Bonding
 Valence electrons are the electrons that are in the highest occupied energy level of an atom of an element
 Valence electrons are usually the only electrons that are involved when a chemical bond is formed,
therefore the valence electrons determine the chemical properties of the element
 Electron (Lewis) dot structures are used to show the valence electrons of an atom – the number of
valence electrons for the representative elements is equal to the group number – Follow Pauli exclusion
and Hund’s rule – start with one in s fill s – then 1 in each p – then go back and double up p
.
.
.
.
.
..
..
H
Be:
B:
.C:
.N:
:O:
:F:
:Ne:
.
.
.
..







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Atoms that have a filled outer shell are stable (Noble gases) all other atoms become stable by forming
bonds to fill their outer shell
Atoms fill their outer shell by gaining, losing, or sharing electrons with one or more atoms
Ionic bonds are formed when one atom takes one or more electron from another atom
The atom that takes the electron becomes negatively charged and is called an anion
The atom that gave up the electron becomes positively charged and is called a cation
The positive and negative charges on the ions causes them to bond together – and the net charge is zero
Covalent bonds are formed between atoms of an element when they become stable by sharing outer shell
electrons
When atoms of elements share electrons equally the bond is called nonpolar and the net charge is zero
When there is an unequal sharing of electrons the bond is called polar and one end of the molecule is
positively charged and the other is negatively charged (Water)
 http://wiki.answers.com/Q/What_makes_water_the_universal_solvent
 http://www.hbci.com/~wenonah/hydro/h2o.htm

By subtracting the electronegativities of two elements the type of bond they will form when they
combine chemically can be predicted
EN Difference
Type of Bond
Example
4
0 to .4
.4 to 1.0
1.0 to 2.0
˃ 2.0
Covalent (nonpolar)
Covalent (Sw polar)
Covalent (very Polar)
Ionic
H to H – (0)
H – Br – (.86)
H – F - (1.9)
Na+ - Cl- - (2.23)
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