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The Periodic Table
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Section 1: History of the Periodic
Table
Many scientist noticed patterns amongst elements
when organized by their properties:
A) Johann Dobereiner (early 1800’s)– recognized
“triads” – groups of three elements with similar
properties. Ex/ Ca, Sr, Ba;
S, Se, Te; and Cl, Br, I
B) John Newlands (1865) recognized rows of
elements with similar properties repeated every 8th
element; pattern called Law of Octaves (oct = 8)
C) Dmitri Mendeleev (1869)


“father” of the periodic table because his organization
of elements (by atomic mass) at the time left gaps in
the table for undiscovered elements at the time.
Ex/ ekaaluminum (eka = “one beyond”) temporary
name for an element that Mendeleev predicted would
be below aluminum (now known as Ga)
 D) Henry Moseley (1914) His work led to
organization of the periodic table by
atomic number
The Periodic Law
States that an element’s physical and chemical properties
are a function of their atomic number.
Therefore, when elements are placed in order of atomic
number, the elements with similar properties occur at
regular intervals (groups!)
**Note: Elements in a group are similar because of their
similar valence electrons.
Valence electrons found in the highest occupied energy
level (outermost shell) of an atom; these electrons
determine chemical reactivity.
Copy each question in notes
1)How many valence electrons are in the following
elements: a) calcium
b) sulfur
c) bromine
Hint: write the electron configuration.
1) a) Calcium: 1s22s22p63s23p64s2
Answer: 2 valence electron
1) b) sulfur: 1s22s22p63s23p4
Ans. 6 valence electrons
1) c) Ans. 7
4) In which period and group is the element whose electron
configuration is [Ne] 3s23p1 ?
Ans. Period 3 group 13
5) Given 1s22s22p63s23p64s23d10 4p5 Give the period and group
numbers.
Ans. Period 4 and group 17
6) Write the outer electron configuration for a group 2 period 4
element.
Ans. 4s2
Class work: copy each question on
back of Periodic Table & keep for
discussion
1)How many valence electrons does potassium have?
2)How many valence electrons does chlorine have?
3)How many valence electrons does Kr have?
4) In which period and group is the element whose electron
configuration is [Ar]4s2 ?
5) Given 1s22s22p63s23p64s23d9 . Give the period and group
numbers.
6) Write the outer electron configuration for a group 1 period 7
element.
The Modern Periodic Table
Periodic Table: arrangement of elements in order of atomic
number so that elements with similar properties fall in
the same column, or group.
Using your text as a reference list notes in your notebook
about the following groups and areas of the periodic
table
Noble Gases, transition metals, inner-transition metals,
representative elements, alkali metals, alkaline earth
metals, halogens, metals, nonmetals, metalloids,
s-block, p-block, d-block, f-block, lanthanides, actinides,
Define each term in your notes
& note location on Periodic Table
Periodic Table
Noble Gases,
transition metals,
inner-transition metals,
representative elements,
alkali metals,
alkaline earth metals,
halogens,
metals,
nonmetals,
metalloids,
s-block, p-block, d-block, f-block,
Lanthanides & actinides
Periods
What determines the length of each period on the periodic table?
Answer: Determined by the number of electrons that can occupy the
sublevels being filled in that period.
sublevels in order
Period #
# of elements in period
of filling:
1
2
1s
2
8
2s 2p
3
8
3s 3p
4
18
4s 3d 4p
5
18
5s 4d 5p
6
32
6s 4f 5d 6p
7
32
7s 5f 6d 7p
The Modern Periodic Table
Periodic Table: arrangement of elements in order of atomic
number so that elements with similar properties fall in
the same column, or group.
The s-Block (group 1)
Group 1 are called Alkali metals
 Extremely reactive metals (hence stored in oil)
 Outermost energy level contains 1 electron in an s orbital
 Soft; cuts easily with a knife
 Not found in nature in elemental state, only in chemical
bonds with other elements.
 React with water to form hydrogen gas & a metallic
hydroxide (alkalis)
The s-Block (group 2)
Group 2 are called the alkaline-earths
 Contain a pair of electrons in the s orbital (2)
 Harder, denser, & stonger than group 1
 Less reactive than alkali metals
 Too reactive to be found as free elements
d-Block Elements: Groups 3-12
d-block are also known as transition metals
 d sublevel first appears when n = 3
 10 electrons fit in d-orbitals
 Good conductors of electricity, high luster
 Reactivity varies; but some are so unreactive
that they exist free in nature (ex/ Platinum &
Gold)
p-Block Elements: Groups 13-18
 Consists of all elements in groups 13-18,
except Helium
 Electrons add to a p sublevel only after the s
sublevel in the energy level is filled.
 p-block contains all 3 element types (metal,
nonmetal, and all 6 metalloids – B, Si, Ge, As,
Sb, & Te)
p-Block
Halogens - Group 17
 Most reactive nonmetals.
 React vigorously with metals to form salts ex/ sodium
chloride NaCl – table salt
 7 valence electrons
p-Block
Noble Gases: group 18
 1894-1900 all noble gases identified
 Inert (non-reactive)
 All have an octet (exception is He with a duet – 2
electrons) of electrons in their valence level (octet means
8 valence electrons)
f-block (below the P.T.)
Lanthanides – Begins after La (Lanthanum);
14 elements below the table (part of f-block)
 Atomic numbers 58 Cerium to 71 Lutetium
 Period 6 (no group numbers assigned – between groups 3 & 4)
Actinides – Begins after Ac (Actinium); below the table too (also
part of f-block); all radioactive
 Atomic numbers 90 Thorium to 103 Lawrencium
 Period 7 (no group numbers assigned – between groups 3 & 4)
Main-group or Representative
Elements




s and p block together elements represent the main-group
Predictive chemistry for these blocks
Valence electrons are easily determined by their group number
Group number
valence electrons








1
2
13
14
15
16
17
18
1
2
3
4
5
6
7
8
Periodicity – repeating patterns
Periodicity with respect to atomic number can be observed in
any group of elements.
 For example:
 Group 2
Group 18
 Atomic number: Difference:
at. #:
diff.
4
2
12
8
10
8
20
8
18
8
38
18
36
18
56
18
54
18
88
32
86
32
Section 3 Electron Configuration
and Periodic Properties(Trends)
Atomic radius one-half the distance of adjacent nuclei
of identical atoms that are bonded together
(a measure of size of atoms)
 refer to fig. 6.15 pg. 175
 Period Trend: Atomic radius decreases L to R*
 **The trend to smaller atoms across a period is
caused by the increasing positive charge on the
nucleus.
 Group Trend: Atomic radius increases down a group
Example questions:
1. Which of the following elements has the
largest atomic radius: Li, O, C, F?
Ans. Li
2. Which has the smallest atomic radius:
Li, K, Cs?
Ans. Li
Atoms can either lose or gain
electrons to become ions.
ion – an atom (or group of bonded atoms) that have a
positive or negative charge.
Metals lose electrons to form positively charged ions called
cations.
 Ex/ Sodium atom: 11 protons & 11 electrons (neutral) Na
 1s22s22p63s1
 Sodium ion: 11 protons & 10 electrons (positive charge)
Na+ 1s22s22p6 (stable electron configuration; like Ne)
Nonmetals gain electrons to form negatively charged
ions called anions.
 Ex/ Chlorine atom: 17 protons & 17 electrons
(neutral) Cl
 1s22s22p63s23p5
 Chloride ion (name changes for neg ions add –ide):
 17 protons & 18 electrons (negative charge) Cl 1s22s22p63s23p6 (stable configuration like Ar)
Ionization Energy (IE)
 IE – energy required to remove an electron from a neutral
atom (kJ/mol) (or First Ionization Energy, IE1 )
 fig. 6.20 in textbook
 Period Trend: IE’s of main-group elements increases L to R
across a period
 Group Trend: IE’s of main-group elements decrease down a
group
2nd and 3rd IE’s increase because each successive electron
removed from an ion feels increasing stronger effective
nuclear charge (pos. nucleus attracting neg. cloud) see p. 177
In each of the following pairs of elements, choose the
element that has the higher first ionization energy.
a) Ca and Ba
b) Ca and Br
c) Ca and K
d) Ca and Mg
Ans. a) Ca b) Br c) Ca d) Mg
Trend: Ionic Radii – size of ions
*same as Atomic Radius trends
 fig. 6.22 pg. 179
Period Trend decreasing L to R
across a period
Group Trend increasing down a
group
Electron Affinity
 EA – The energy change that occurs when an electron is
acquired by a neutral atom (kJ/mol)
 Period Trend - EA increases L to R across a period.
 Group Trend - EA decreases down a group
• Element with highest electron affinity is Fluorine
(-339.9 kJ/mol). The higher neg. value the higher the
affinity for electrons.
Electronegativity
 EN - Measure of the ability of an atom in a
chemical compound to attract electrons in a
chemical bond.
 0 to 4.0 scale (no units) Fluorine is the most
electronegative!
 table 6.2 pg 181
 Period Trend – EN inc. L to R across a period
 Group Trend – EN dec. down a group or
remain about the same.