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Unit 3: Periodic Tablehttp://ed.ted.com/periodic-videos
How did chemists begin to organize known elements?
Cu, Au and Ag have been known for thousands of years…. but by 1700 only 13 elements had been identified.
In 1829, Dobereiner first used chemical and physical properties to sort elements into groups (triads).
Mendeleev’s Periodic Table
In 1869, Mendeleev proposed arranging elements by atomic mass and repeating chemical and physical properties.
Mendeleev left gaps on his PT and predicted new elements would be discovered to fill the spaces.
Moseley’s Periodic TableIn 1913, Moseley re-arranged elements by atomic number. Moseley used x-rays to determine the exact atomic number of each element establishing Periodic Law. (physical and chemical properties of elements repeat when elements arranged by increasing atomic number-elements with similar properties appear periodically)
The gaps left in Moseley’s tables had atomic numbers of 43, 61, 72, 75, 86, 87, and 91. These elements are all radioactive!
Organization of the Periodic Table
Group – vertical columns; elements in the same group have same # of electrons in their valence shell and share common characteristics.
Period – horizontal rows; the filling of each energy level with electrons corresponds to a row on the table; do not share common characteristics.
The 3 broad classes of elements
Metals are to the left of the zig-zag line (except H)
Non-metals are to the right of the zig-zag line
Elements touching the line are called “metalloids” (except Al)
Properties of Metals
Shiny
Malleable (can be bent or hammered flat)
Ductile (can be drawn into wire)
Good conductors of heat and electricity
Solids at room temperature (except for ________)
Properties of Nonmetals
Dull
Brittle (nonmalleable)
Poor conductors of heat and electricity
Gases, liquids, or low-melting-point solids
Properties of Metalloids
B, Si, Ge, As, Sb, Te, Po
Have properties of metals and nonmetals
Commonly used in electronics as a semiconductor (ex: Si wafers)
Alkali Metals
Group 1
One valence electron
Very reactive – why?
What is the one exception in group 1 that is not an alkali metal?
Alkaline Earth Metals
Group 2
Two valence electrons
Less reactive than group 1
How to remember names?
Alkali – one word, group 1 Alkaline Earth – two words, group 2
Transition Metals
Groups 3-12
Do not follow patterns as well as groups 1, 2, and 13-18
# of valence electrons harder to predict
Halogens
Group 17
Seven valence electrons
Very reactive –why?
Noble Gases
Group 18
Octet of valence electrons. (Full valence shell)
Inert – unreactive
Don’t form ions
Lanthanide and Actinide Series
Also called Rare Earth Elements
The Lan. Series is part of Period 6The Act. Series is part of Period 7
PERIODIC TRENDS
Periodic Trends
Because of the established Periodic Law, scientists began to notice tendencies of certain elemental characteristics to increase or decrease along a row or column of the periodic table of elements. These tendencies are called Periodic Trends.
4 factors that cause the trends
1. Nuclear Pull (Z) – the number of protons
The protons pull on the outer electrons. The more protons, the more pull exerted by the nucleus on the outer electrons.
2. Electron repulsion –size of e- cloud
The more electrons in
an atom’s electron
cloud, the more they
are pushed away from
each other (due to having
the same charge), making
a bigger cloud.
4 factors that cause the trends
3. Shielding electrons –all inner e- shield the valence electrons from nuclear pull
4 factors that cause the trends
4. Zeff – the “effective” nuclear pull on outer electrons. The nuclear pull taking into account the shielding electrons which are taking most of the force.
Zeff= # protons - # non-valence electrons
Which element has more effective nuclear pull?
Fluorine MagnesiumOR
4 factors that cause the trends
Atomic Radius Trend
• Increases down a column because the valence electrons are in a farther energy level (higher energy)
• Decreases across a period because the nuclear pull is increasing and pulling the energy levels in.
Atomic Size DECREASES
Ato
mic
Siz
e I
NC
REA
ES
Atomic Radius Trend
Using the four factors that determine periodic trends, explain the sizes of the atoms.
IONS
• An atom or group of atoms that has a positive or a negative charge
– Atoms are electrically neutral (protons=electrons)
– Atoms can lose electrons to form cations (+ charge)
– Atoms can gain electrons to form anions (- charge)
Ion SizeMetal ions (cations) are smaller than their atoms
– metal ions LOSE electrons, causing less electron repulsion, and smaller size.
Nonmetal ions (anions) are larger than their atom
– Nonmetal ions GAIN electrons, causing more electron repulsion, and larger size.
Which of the following elements has the largest atomic radii?
K or Rb
If the two elements were to form ions, which element would have the largest ion size?
K or Rb
Ionization Energy
• The energy needed to remove an electron from an atom.
• The greater the ionization energy, the more difficult it is to remove an electron.
Ionization Energy
– Decreases down a group because there are more shielding electrons, so it takes LESS ENERGY to “steal” an electron.
– Increases across a period because the nuclear pull on those electrons is increased with no extra shielding, so it takes MORE ENERGY to get the electrons away.
Ionization Energy INCREASES
Ion
izat
ion
En
erg
y D
ECR
EASE
S
Using the four factors that determine periodic trends, explain why it takes more energy to remove an electron from a fluorine atom vsan oxygen atom.
Electronegativity
• The ability of an atom to take an electron from another atom. Decreases down a group because there are moreelectrons to shield the nucleus (which does thepulling)Increases across a period because of increased Z(nuclear pull)
ElectonegativityINCREASES
Ele
ctro
neg
ativ
ity
DEC
REA
SES
Electron Affinity
• The energy change that occurs when an atom acquires an electron.
• If an atom begins as very unstable (think chaos, lots of energy) and becomes stable by gaining an electron, there was a LARGE ENERGY CHANGE
• Trend correlates with electronegativity, because both involve ability to take electrons
• Decreases down a group and increases across a period
Also important: the most reactive corners of PTare lower left (francium) and upper right (fluorine).Noble gases not included- why?
Using the four factors that determine periodic trends, explain why electrons are more attractive to non-metals than metals.
1. Which is the smallest atom? Explain.Na Li Be
2. Which has the highest electronegativity? Explain.As Sn S
3. In the following pairs, which has the larger atomic radius? Explain.
Mg or Ba Cu or Cu2+ S or S2-
4. In the following pairs, which has the higher ionization energy? Explain.
Li or Cs Ca or Br
Trend What is it?Trend across a
periodTrend across a
groupWhat causes
trend?
Atomic Radius
Ionization Energy
Electronegativity