1. Unit 3: Periodic Table

2. 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).

3. 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.

4. Moseley’s Periodic Table
In 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!

5. 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.

6. 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)

7. 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 ________)

8. Properties of Nonmetals
Dull
Brittle (nonmalleable)
Poor conductors of heat and electricity
Gases, liquids, or low-melting-point solids

9. 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)

10. Alkali Metals Group 1 One valence electron Very reactive – why?
What is the one exception in group 1 that is not an alkali metal?

11. 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

12. Transition Metals Groups 3-12
Do not follow patterns as well as groups 1, 2, and 13-18
# of valence electrons harder to predict

13. Halogens
Group 17
Seven valence electrons
Very reactive – why?

14. Noble Gases Group 18 Octet of valence electrons. (Full valence shell)
Inert – unreactive
Don’t form ions

15. Lanthanide and Actinide Series
Also called Rare Earth Elements
The Lan. Series is part of Period 6
The Act. Series is part of Period 7

16. Periodic trends

17. 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.

18. 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.

19. 4 factors that cause the trends
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.

20. 4 factors that cause the trends
3. Shielding electrons – all inner e- shield the valence electrons from nuclear pull

21. 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
Magnesium OR

22. 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
Draw it!
Atomic Size INCREAES

23. Atomic Radius Trend

24. Using the four factors that determine periodic trends, explain the sizes of the atoms.

25. 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)

26. Ion Size Metal 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.
If the two elements were to form ions, which element would have the largest ion size?
K or Rb
Which of the following elements has the largest atomic radii?
K or Rb

27. 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.

28. Ionization Energy DECREASES
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
Draw it!
Ionization Energy DECREASES

29. Using the four factors that determine periodic trends, explain why it takes more energy to remove an electron from a fluorine atom vs an oxygen atom.

30. Electronegativity DECREASES
The ability of an atom to take an electron from another atom.
Decreases down a group because there are more
electrons to shield the nucleus (which does the
pulling)
Increases across a period because of increased Z
(nuclear pull)
Electonegativity
INCREASES
Draw it!
Electronegativity DECREASES

31. 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 PT
are lower left (francium) and upper right (fluorine).
Noble gases not included- why?

32. Using the four factors that determine periodic trends, explain why electrons are more attractive to non-metals than metals.

33. 1. Which is the smallest atom. Explain. Na Li Be 2
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

34. Trend
What is it?
Trend across a period
Trend across a group
What causes trend?
Atomic Radius
Ionization Energy
Electronegativity