periodic table & trends - the colony high...
TRANSCRIPT
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Periodic Table & Trends
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History of the Periodic Table
• 1871 – Mendeleev arranged the elements according to: 1. Increasing atomic mass 2. Elements w/ similar properties were put in the same row
• 1913 – Moseley arranged the elements according to: 1. Increasing atomic number 2. Elements w/ similar properties were put in the same column
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Group Names Alkali +1
Alkaline Earth Metals +2
+3
-3
-2
Halogen -1
Noble Gases 0
H 1
He 2
Li 3
Be 4
B 5
C 6
N 7
O 8
F 9
Ne 10
Na 11
Mg 12
Al 13
Si 14
P 15
S 16
Cl 17
Ar 18
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S & P block – Representative Elements Metalloids (Semimetals, Semiconductors) – B,Si, Ge,
As, Sb, Te (properties of both metals & nonmetals)
Columns – groups or families Rows - periods
METALS
TRANSITION METALS
NONMETALS
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The Periodic Law
• Mendeleev understood the ‘Periodic Law’ which states:
• When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.
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The Periodic Law
• Atoms with similar properties appear in groups or families (vertical columns) on the periodic table.
• They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.
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Valence Electrons
• Do you remember how to tell the number of valence electrons for elements in the s- and p-blocks?
• How many valence electrons will the atoms in the d-block (transition metals) and the f-block (inner transition metals) have?
• Most have 2 valence e-, some only have 1.
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A Different Type of Grouping
• Besides the 4 blocks of the table, there is another way of classifying element:
• Metals • Nonmetals • Metalloids or Semi-metals. • The following slide shows where each
group is found.
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Metals, Nonmetals, Metalloids
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Metals, Nonmetals, Metalloids
• There is a zig-zag or staircase line that divides the table.
• Metals are on the left of the line, in blue.
• Nonmetals are on the right of the line, in orange.
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Metals, Nonmetals, Metalloids
• Elements that border the stair case, shown in purple are the metalloids or semi-metals.
• There is one important exception.
• Aluminum is more metallic than not.
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Metals, Nonmetals, Metalloids
• How can you identify a metal? • What are its properties? • What about the less common nonmetals? • What are their properties? • And what the heck is a metalloid?
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Metals
• Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity.
• They are mostly solids at room temp.
• What is one exception?
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Nonmetals
• Nonmetals are the opposite.
• They are dull, brittle, nonconductors (insulators).
• Some are solid, but many are gases, and Bromine is a liquid.
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Metalloids • Metalloids, aka semi-metals
are just that. • They have characteristics of
both metals and nonmetals. • They are shiny but brittle. • And they are
semiconductors. • What is our most important
semiconductor?
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Periodic Trends
• There are several important atomic characteristics that show predictable trends that you should know.
• The first and most important is atomic radius.
• Radius is the distance from the center of the nucleus to the “edge” of the electron cloud.
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Atomic Radius
• Since a cloud’s edge is difficult to define, scientists use define covalent radius, or half the distance between the nuclei of 2 bonded atoms.
• Atomic radii are usually measured in picometers (pm) or angstroms (Å). An angstrom is 1 x 10-10 m.
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Covalent Radius
• Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å.
2.86 Å
1.43 Å 1.43 Å
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Atomic Radius
• The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family.
• Why? • With each step down the family, we add an
entirely new PEL to the electron cloud, making the atoms larger with each step.
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Atomic Radius
• The trend across a horizontal period is less obvious.
• What happens to atomic structure as we step from left to right?
• Each step adds a proton and an electron (and 1 or 2 neutrons).
• Electrons are added to existing PELs or sublevels.
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Atomic Radius
• The effect is that the more positive nucleus has a greater pull on the electron cloud.
• The nucleus is more positive and the electron cloud is more negative.
• The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.
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Effective Nuclear Charge
• What keeps electrons from simply flying off into space?
• Effective nuclear charge is the pull that an electron “feels” from the nucleus.
• The closer an electron is to the nucleus, the more pull it feels.
• As effective nuclear charge increases, the electron cloud is pulled in tighter.
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Atomic Radius
• The overall trend in atomic radius looks like this.
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Atomic Radius
• Here is an animation to explain the trend. • On your help sheet, draw arrows like this:
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The Octet Rule • The “goal” of most atoms (except H, Li and
Be) is to have an octet or group of 8 electrons in their valence energy level.
• They may accomplish this by either giving electrons away or taking them.
• Metals generally give electrons, nonmetals take them from other atoms.
• Atoms that have gained or lost electrons are called ions.
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Ions
• When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion.
• In the same way that nonmetal atoms can gain electrons, metal atoms can lose electrons.
• They become positively charged cations.
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Ions
• Here is a simple way to remember which is the cation and which the anion:
This is a cat-ion. This is Ann Ion.
He’s a “plussy” cat! She’s unhappy and negative.
+ +
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Ionization Energy • This is the second important periodic trend. • If an electron is given enough energy (in the
form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely.
• The atom has been “ionized” or charged. • The number of protons and electrons is no
longer equal.
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Ionization Energy • The energy required to remove an electron
from an atom is ionization energy. (measured in kilojoules, kJ)
• The larger the atom is, the easier its electrons are to remove.
• Ionization energy and atomic radius are inversely proportional.
• Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.
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Ionization Energy
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Ionization Energy (Potential)
• Draw arrows on your help sheet like this:
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Electron Affinity
• What does the word ‘affinity’ mean? • Electron affinity is the energy change that
occurs when an atom gains an electron (also measured in kJ).
• Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.
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Electron Affinity
• Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy.
• If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic.
• This is true for the alkaline earth metals and the noble gases.
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Electron Affinity
• Your help sheet should look like this:
+ +
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Metallic Character
• This is simple a relative measure of how easily atoms lose or give up electrons.
• Your help sheet should look like this:
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Electronegativity • Electronegativity is a measure of an atom’s
attraction for another atom’s electrons. • It is an arbitrary scale that ranges from 0 to 4. • The units of electronegativity are Paulings. • Generally, metals are electron givers and have
low electronegativities. • Nonmetals are are electron takers and have
high electronegativities. • What about the noble gases?
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Electronegativity
• Your help sheet should look like this:
0
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Overall Reactivity
• This ties all the previous trends together in one package.
• However, we must treat metals and nonmetals separately.
• The most reactive metals are the largest since they are the best electron givers.
• The most reactive nonmetals are the smallest ones, the best electron takers.
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Overall Reactivity
• Your help sheet will look like this:
0
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Cation Formation
11p+
Na atom
1 valence electron
Valence e- lost in ion formation
Effective nuclear charge on remaining electrons increases.
Remaining e- are pulled in closer to the nucleus. Ionic size decreases.
Result: a smaller sodium cation, Na+
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Anion Formation
17p+
Chlorine atom with 7 valence e-
One e- is added to the outer shell.
Effective nuclear charge is reduced and the e- cloud expands.
A chloride ion is produced. It is larger than the original atom.