Download Topic 3 Periodic Trends Notes 14-15

Topic 3: Periodicity
3.1: The periodic table
Essential Idea: The arrangement of elements in the
periodic table helps to predict their configurations
Nature of Science: Obtain evidence of scientific theories
by making and testing predictions based on them –
scientists organize subjects based on structure and
function; the periodic table is a key example of this.
Early models of the periodic table from Mendeleev, and
later Moseley, allowed for the prediction of properties of
elements that had not yet been discovered. (1.9)
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Topic 3: Periodicity
3.1: The Periodic Table
Understandings:
1. The Periodic Table is arranged into four blocks
associated with the four sub-levels: s, p, d, and g
2. The Periodic Table consists of groups (vertical columns)
and periods (horizontal rows)
3. The period number (n) is the outer energy level that is
occupied by electrons
4. The number of the principal energy level and the
number of the valence electrons in an atom can be
deduced from its position on the periodic table.
5. The periodic table shows the positions of metals, nonmetals, and metalloids.
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Topic 3: Periodicity
3.1: The Periodic Table
Applications and Skills:
1. Deduction of the electron configuration of an atom from
the element’s position on the Periodic Table and vice
versa.
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3.1 Background Information
Development of the Periodic Table
• Johan Dobereiner
Grouped similar elements into
groups of 3 (triads) such as
chlorine, bromine, and iodine.
(1817-1829).
• John Newlands
Found every eighth element
(arranged by atomic weight)
showed similar properties. Law
of Octaves (1863).
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3.1 Background Information
Development of the Periodic Table
• Dmitri Mendeleev
Arranged elements by similar
properties but left blanks for
undiscovered elements (1869).
Crash Course Video:
https://www.youtube.com/watch?
v=0RRVV4Diomg
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3.1 Background Information
Development of the Periodic Table
• Henry Mosley
Arranged the elements by
increasing atomic number instead
of mass (1913)
• Glen Seaborg
Discovered the transuranium
elements (93-102) and added the
actinide and lanthanide series
(1945)
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3.1 U2 Groups and Periods
Elements arranged by increasing atomic number into
• periods (rows) 1-7, which relate to energy levels
• groups or families (columns), which share similar
properties
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3.1 U5. Metals, Non-metals, Metalloids
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www1.whsd.net
3.1 U5. Metals, Non-metals, Metalloids
Metals
– Left side of the periodic table
(except hydrogen).
– Good conductivity of heat and
electricity
– Luster (shiny)
– Ductile (drawn into wires)
– Malleable (hammered into
sheets)
– Lose electrons in chemical
reactions (oxidized)
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3.1 U5. Metals, Non-metals, Metalloids
• Alkali metals: Group 1
Alkaline earth metals: Group
2
• Transition metals: Group 312, lanthanide & actinide series
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3.1 U5. Metals, Non-metals, Metalloids
Nonmetals
– Right side of the periodic table
– Poor conductors of heat and
electricity
– Non-lustrous
– Gain electrons in chemical
reactions (reduced)
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3.1 U5. Metals, Non-metals, Metalloids
– Halogens: Group 17
– Noble gases: Group 18
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3.1 U5. Metals, Non-metals, Metalloids
Metalloids
– Between metals and nonmetals,
along the stair step line
– Properties intermediate between
metals and nonmetals
- Some are semi-conductors
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3.1 U5. Metals, Non-metals, Metalloids
Metalloids
– Boron (B), Silicon (Si),
Germanium (Ge), Arsenic (As),
Antimony (Sb), Tellurium (Te),
Astatine (At)
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3.1 The Periodic Table
Amazing Periodic Table Resource
http://www.rsc.org/periodic-table
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3.1 U4. Valence Electrons
Valence Electrons: electrons in
the outermost (highest) energy
level of s and p sublevels
–
–
–
–
–
Group 1 elements s1 = 1
Group 2 elements s2 = 2
Group 13 elements s2 p1 = 3
So on and so forth
Group 18 s2 p6 = 8 (except for
helium, which has 2)
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3.1 U4. Valence Electrons
Drawing Valence Electrons:
(Lewis Dot Diagrams)
Write the element symbol,
starting at the top, add dots 1 at
a time clockwise, separate to 1
per side up to 4, then double up.
No more than 8 dots total.
Group 1 = 1 dot
Group 2 = 2 dots
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3.1 U4. Valence Electrons
Lewis (electron) dot diagrams
Group 1: 1 dot
X
Group 15: 5 dots
X
Group 2: 2 dots
X
Group 16: 6 dots
X
Group 3: 13 dots
X
Group 17: 7 dots
X
Group 4: 14 dots
X
Group 18: 8 dots (except He)
X
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IB Topic 3: Periodicity
3.2: Physical properties
• Essential Idea: Elements show trends in their physical
and chemical properties across periods and down
groups.
• Nature of Science: Looking for patterns – the position
of an element in the periodic table allows scientists to
make accurate predictions of its physical and chemical
properties. This gives scientists the ability to synthesize
new substances based on the expected reactivity of
elements (3.1)
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IB Topic 3: Periodicity
3.2: Physical properties
Understandings:
1.Vertical and horizontal trends in the periodic table exist
for atomic radius, ionic radius, ionization energy, electron
affinity, and electronegativity.
1.Trends in metallic and non-metallic behaviour are due to
the trends above.
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IB Topic 3: Periodicity
3.2: Physical properties
Applications and Skills:
1.Prediction and explanation of the metallic and nonmetallic behaviour of an element based on its position in
the periodic table.
1.Discussion of the similarities and differences in the
properties of elements in the same group, with reference to
alkali metals (group 1) and halogens (group 17).
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IB Topic 3: Periodicity
3.2: Physical properties
Important Terms:
Core Electrons: the inner non-valence electrons of the
atom
Nuclear Charge: the number of protons in the nucleus of
the atom
Shielding (screening): the core electrons shield (block)
the valence electrons from the nucleus, reducing the
nuclear charge
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3.2 U1. Atomic Radius
Atomic Radii: size of the radius of the atom
• Group trend
– Atomic size increases as you move down a group of
the periodic table.
– Reason: each row going down adds an energy level,
increasing the size of the atom
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3.2 U1. Atomic Radius
Atomic Radii: size of the radius of the atom
• Periodic trend
– Atomic size decreases as you move across a period.
– Reason: The increase in nuclear charge (more protons in
the nucleus) increases the attraction to the outer shell so
the outer energy level progressively becomes closer to the
nucleus decreasing the size of the atom
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Atomic Radii- Your graph should look similar
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3.2 U1. Ionic Radius
Ionic Radius: size of the ion’s radius
• Positive ions (cations) are smaller than their
atoms.
– Fewer electrons so nucleus attracts remaining electrons
more strongly
– One fewer energy level since valence electrons removed.
• Negative ions (anions) are larger than their atoms
– More electrons so nucleus has less attraction for them
– Greater electron-electron repulsion
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3.2 U1. Ionic Radius
Ionic Radius: size of the ion’s radius
• Group trend
– Ions get larger down a group
– Reason: energy levels are added, electrons are farther
away from nucleus
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3.2 U1. Ionic Radius
Ionic Radius: size of the ion’s radius
• Periodic trend
– Ions decrease as you move across a period.
– Reason: This increase in nuclear charge increases the
attraction to the outer shell so the outer energy level
progressively becomes closer to the nucleus
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3.2 U1. First Ionization Energy
First Ionization Energy: The energy required to remove the first
electron from a gaseous atom.
*Second ionization removes the second electron and so on. Can be
used to predict ionic charges.
• Group trend
– Generally decreases as you move down a group in the
periodic table
– Reason: Since atomic radius increases down a group, the
outermost electron is farther away from the nucleus and is
easier to remove. The shielding effect of the core electrons
also increases.
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3.2 U1. First Ionization Energy
First Ionization Energy: The energy required to remove
the first electron from a gaseous atom.
• Periodic Trend
– Increases as you move from left to right across a
period.
– Reason: effect of increasing nuclear charge (more protons)
makes it harder to remove an electron, stronger attraction
between the nucleus and the outer electrons, the atomic
radius also decreases, electrons are closer to the nucleus,
more difficult to remove
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Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell
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3.2 U1. Electronegativity
Electronegativity: The relative attraction an atom has for
the shared pair of electrons in a covalent bond.
*Helps predict the type of bonding (ionic/covalent).
• Group trend
– Generally decreases as you move down a group in the
periodic table.
– Reason: atomic radius increases and shielding effect of
core electrons, electrons are farther away, less likely to
attract more
– For metals, the lower the number the more reactive.
– For nonmetals, the higher the number the more reactive.
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3.2 U1. Electronegativity
Electronegativity: The relative attraction an atom has for
the shared pair of electrons in a covalent bond.
• Periodic Trend
– Increases as you move from left to right across a
period.
– Reason: nuclear charge increases, atomic radius increases
creating greater attraction of electrons
– Nonmetals have a greater attraction for electrons than
metals.
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Electronegativity
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3.2 U1. Electron Affinity
Electron Affinity: the energy required to detach an
electron from a singly charged negative ion in the gas
phase.
• Group Trend
– Generally become less negative as you move down a
group
– Patterns vary by group, do not show a clear trend down a
group
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3.2 U1. Electron Affinity
Electron Affinity: the energy required to detach an
electron from a singly charged negative ion in the gas
phase.
• Period Trend
– Generally become more negative as you move across
a period from left to right
– Reason: gaining electrons makes negative ions more
stable
– Trends are not as well highlighted as other trends
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3.2 A2 Similarities and Differences in Groups
•
•
•
•
•
Group 1: Alkali Metals
Have 1 valence electron
Shiny, silvery, soft metals
React with water & halogens
Oxidize easily (lose electrons)
Reactivity increases down the
group
– Atomic radius increases and
ionization energy decreases
going down group, easier to lose
an e-.
Images.fineartamericacom
www.Britannica.com
www.elementsales.com
www.greatmining.com
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3.2 A2 Similarities and Differences in Groups
•
•
•
•
Group 17: Halogens
Have 7 valence electrons
Colored gas (F2, Cl2); liquid
(Br2);
Solid (I2)
Oxidizer (gain electrons)
Reactivity decreases down the
group
– Atomic radius increases, more
difficult to gain an e41