ch. 5: periodic table · 1. without looking at the periodic table, identify the group, period, and...
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Ch. 5: Periodic
Table
C. Goodman, Doral Academy Preparatory High School, 2011-2013
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Essential Question: Section 5.1
1. What is the history of the development of the Periodic Table?
2. What is the periodic law, and how can it be used to predict physical and chemical properties of elements?
3. What is the overall organization of the modern Periodic Table? 1. Three types of elements
2. Named groups
3. Other families
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Section 5.1 Vocabulary
•Mendeleev
• Moseley
• Periodic Law
• Period
• Group
•Main group elements
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Periodic Table - Definition
Periodic Table -- an arrangement of the elements in order of their atomic numbers so that elements with the same chemical properties are in the same group (family). Examples: halogens, noble gases, alkali metals.
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Why is it cool?
• http://www.youtube.com/watch?v=u2ogMUDBaf4&playnext=1&list=PLAC3A0775D813045F&feature=results_main
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History of the Periodic Table I
• Mendeleev: 1869
– Atomic mass
– Repeating (periodic) patterns of reactivity
•In his favor: predicted
the discovery of Gallium,
which was isolated in his
lifetime
•Certain characteristic
properties of elements
can be foretold from their
atomic weights
•Problem:
Iodine and Tellurium
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History of the Periodic Table II
Moseley: 1914 Atomic number
each element has a unique atomic number; resequenced the table by electronic charge (=atomic #) rather than atomic weight.
Periodic Law
•In his favor:
solved the “Iodine and
Tellurium” problem
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Moseley’s Periodic Law - Definition
• Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic numbers.
http://www.youtube.com/watch?v=OduTDU
GeAXEFind
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How to use the periodic table…
Atomic number: # of protons in the nucleus of an atom
Symbol Basically the abbreviation for the element
Average atomic mass # of Protons + # of Neutrons (amu) Weighted average mass of isotopes of the element
Remember nuclear notation for isotopes? Notice that the atomic mass is a whole number – it’s not an average Also notice the different locations of the atomic mass and atomic number.
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Groups (families)
The Columns
Elements in
groups have
similar chemical
properties
Periods
The Rows
Elements properties vary
across periods
The length of each period is
determined by the number of
electrons that can occupy the
sublevels being filled in that
period
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Important terms
• Main Group Elements s-block + p-block elements
• Transition metals d-block elements
• Lanthanides and actinides f-block elements
• Metalloids, metals, non-metals (see below)
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2 Main Sections in Periodic Table
Metals Nonmetals
- Majority of elements -Good Electrical & Heat Conductors - Room temperature = most solids -Contain properties
- Malleability - Ductility - High tensile strength
- Poor Electrical Conductors - Poor Heat Conductors - Room temperature = most gases - One is a liquid at r.t. = Bromine - Solid nonmetals generally brittle
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Names of groups
• Group 1a – Alkali metals
• Group 2a – Alkali earth metals
• Group 7a (17) – halogens
• Group 8a (18) – Noble gases
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Names of families– Transition metals
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Names of families– Semiconductors
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Section 2: Electron Configuration
1. What is the relationship between the location of atoms in a group, their electron configuration, and their chemical and physical properties?
2. What are the s-, p-, d-, and f-blocks, and how can their electron configurations of their elements be determined?
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Section 5.2 vocabulary
• Ion
• Valence
• Valence electrons
• s-, p-, d- and f-block elements
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Valence Electrons • Valence = outermost
energy level in which contains electrons (in unexcited state).
• Valence electrons are the electrons on the outermost energy level of the element.
• The number of valence electrons determines the type of chemical reactions available to the element!
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What does this have to do with groups?
• All main group elements in a particular group have the same number of valence electrons.
• Prove it?
• Heh heh heh – that’s your job!
• Write electron configurations, noble gas notation, of the s and p block elements in the first 5 rows. Write valence electrons (s/b only) in contrasting color
• Elements hydrogen – xenon, for columns #1a, 2a, 3a, 4a, 5a, 6a, 7a, 8a
• Huhhh? See next slide
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Valence electrons &energy levels
• Purpose: to determine the relationship between the group number and number of valence electrons.
• Procedures 1. For each of the elements in the first 5 periods of the following
groups… group (1a, 2a, 3a, 4a, 5a, 6a, 7a, and 8a)
2. Write the name of the element
3. Write the type of element
4. Write the electron configuration, using noble gas notation
• Conclusion (Answer the following question)s: – 1. How does the group # relate to the number of valence
electrons?
– 2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?
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Valence electrons & Energy Levels
• Conclusion (Answer the following questions): – 1. How does the group # relate to the number of valence
electrons?
– 2. How do you think the chemical reactivity of the elements in a particular group, relates to this number of valence electrons?
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Group 14
Element Electron configuration – noble gas notation
Number of valence electrons
Carbon [He]2s22p2 4
Silicon [Ne]3s23p2 4
Germanium [Ar]4s23d104p2 4
Tin [Kr]5s24d105p2 4
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Valence Electrons
• Main group elements have characteristic numbers of valence electrons.
• Group 1 – 1 valence electron
• Group 2 – 2 valence electrons
• Groups 13-18
– # valence electrons = Group # - 10
– Example: Group 13 elements have 13-10 = 3 valence electrons
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Valence Electrons • Main group elements have characteristic
numbers of valence electrons.
• S block
– Group 1 – 1 valence electron
– Group 2 – 2 valence electrons
• P block
– Groups 13-18
• Group # - 10
• Example: Group 13 elements have 13-10 = 3 valence electrons
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Relationship Between Periodicity and Electron Configurations
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Sample problem A
1. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s2 is located.
2. Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)?
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Sample problem B
An element has the electron configuration [Kr] 5s24d5. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.
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Sample problem C
Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, use your periodic table to name the element, and identify it as a metal, nonmetal, or metalloid.
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In book, p. 135
1: Identify period, block, group, element
[Kr]5s2
2. write configuration of…
a. Group 2 elements
b. the group 2 element in the fourth period.
c. the element in the 3rd period, group 15
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Sample Problem C Solution
• p-block (group # >12
• 14-10 = 4 electrons in s, p
• 2 e- in s, 2e- in p
• The outer electron configuration is 2s22p2.
• The element is carbon, C, which is a nonmetal.
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More practice problems!
Name the block and group in which each of the following elements is located in the periodic table. Use the periodic table to name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity.
1. [Xe] 6s24f145d8
2. [Ne]3s23p2
3. [Ne]3s23p5
4. [Xe]4f66s1
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How do I know which groups are more
reactive than others?
For main group elements, look at the
number of valence (s/p) electrons
8 valence electrons (noble gases)
= not reactive
Less reactive 4<3<2<1 More reactive
Less reactive 4<5<6<7 More reactive
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• Sample Problem D Solution a. The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially
filled with nine electrons. Therefore, this element is in the d block.
• The element is the transition metal platinum, Pt, which is in Group 10 and has a low reactivity.
• b. The incompletely filled p sublevel shows that this element is in the p block.
• A total of seven electrons are in the ns and np sublevels, so this element is in Group 17, the halogens.
• The element is chlorine, Cl, and is highly reactive.
Section 2 Electron Configuration
and the Periodic Table Chapter 5
Periods and Blocks of the Periodic Table, continued
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• Sample Problem D Solution, continued • c. This element has a noble-gas configuration and thus is in Group 18
in the p block.
• The element is argon, Ar, which is an unreactive nonmetal and a noble gas.
• d. The incomplete 4f sublevel shows that the element is in the f block and is a lanthanide.
• Group numbers are not assigned to the f block.
• The element is samarium, Sm. All of the lanthanides are reactive metals.
Section 2 Electron Configuration
and the Periodic Table Chapter 5
Periods and Blocks of the Periodic Table, continued
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Section 3: Periodic Trends Essential Questions
1. Compare the periodic trends of atomic radii,
ionization energy, electronegativity, and state the
reasons for these variations.
2. What are valence electrons, and how many are
present in atoms of each main-group element?
3. Compare the atomic radii, ionization energies, and
electronegativities of the d-block elements with
those of the main-group elements.
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Section 5.3 Vocabulary
• Atomic radius
• Ion
• Ionization energy
• Cation
• Anion
• Electron affinity
• Electronegativity
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Atomic Radii: ½ the distance between
the nuclei of identical atoms bonded
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DECREASES across periods
Because of increasing positive
charge of the nucleus
Holds electrons more tightly
INCREASES down groups
Higher principle quantum
number
Valence electrons in higher
main energy levels
Located farther from the
nucleus
Periodic Trends: Atomic Radii
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IONS
Watch it kid,
I’ve got my ion you.
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Ion – Definition An Ion is…
An atom or group of bonded atoms, which has a positive (+) or negative (-) charge.
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There are two kinds of ions…
Cation • CRUNCH (subtract) an electron
• This results in a positive charge
• When an electron is removed, the atom loses bulk (like a muscle which shrinks when it atrophies)
• So, the radius of a cation is smaller than the atomic radius
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Anions
• ADD an electron
• This results in a NEGATIVE charge
• When an electron is added, the atom gains bulk (like a muscle which grows when you work out)
• So, the radius of a anion is larger than the atomic radius
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Two sodium atoms bumped into each other.
One said: "Why do you look so sad?“ The other responded:
"I lost an electron.“
The first one asked "Are you sure?“
The other replied "I'm positive."
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Which elements form
which type of ion?
Metal on left tend to form cations (+)
Nonmetals at the upper right tend to form
anions (-)
Hydrogen is a non-metal, but it forms a
cation (+)
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Periodic Trends: Ionic Radii – very similar to trends for atomic radii
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Ionization Energy (IE) Definition
Ionization energy is…
The energy required to remove one electron from a neutral atom of an element, forming a cation.
Note: does not apply to formation of anions!
Unit of measure: kJ/mol
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It takes energy to “steal” electrons
Ionization energy is
(almost always)
positive (J)
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Look p. 145, Figure 3.4
Low IE - lose electrons easily
High IE – it’s harder to lose electrons
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Ionization Energy:
Atom (neutral) + energy
Cation (A+)+ (e-) (removed)
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Ionization Energy (IE)
IE increases across periods because of increasing nuclear charge.
Why?
Higher nuclear positive charge attracts e- in same energy level more
strongly
Trend down groups decrease because of more electrons between the
nucleus and furthest out electrons (lower nuclear charge)
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Watch out!!!
Ionization energy only applies to the
formation of cations.
If you want to talk anion formation,
you need electron affinity.
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Electron Affinity The energy change that occurs when an
electron is acquired by a neutral atom
Most release energy when they acquire an
electron: Take a neutral atom (A) add an electron (e-)
get an anion (A-) + energy is released
Some must be “forced” to gain an electron
by adding energy: A + e- + energy A-
Common unit used is kJ/mol
Some positive affinities are difficult to
determine with any accuracy
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Here is a mnemonic for electron affinity:
If the ion is negatively charged (anion), the electron affinity is more
strongly negative.
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Why do some elements give up their electrons more easily?
Q: Which group of elements do you think hold on to their
electrons the most strongly?
A: The noble gases! Their valence is full, so it’s very hard to
remove their electrons!
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Electron Affinity Period Trends: Halogens (VII A) gain
electrons the most readily
Group Trends: Electrons add with greater
difficulty down a group
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Electron Affinity
• Question Based on this trend what elements will most likely form cations?
• Which will most likely form anions?
• Which will most likely have small ionic radii?
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Adding Electrons to Negative Ions
Difficult to add a second electron
to an already negatively charged ion
Second electron affinities are
therefore all positive
Halogens become negative ions by
adding one electron (i.e. Cl-)
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Electronegativity
• A measure of the ability of an atom in a
chemical compound to attract
electrons from another atom in the
compound
• Most electronegative element is
Fluorine
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Electronegativity
Trend across periods increase (some exceptions)
More or less, trend down groups decrease
Similar to IE
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Additional Help Website
http://www.hcc.mnscu.edu/chem/stacks.php