PPNOTES – CHAPTER 6: THE PERIODIC TABLE 6.1 – Organizing the Elements

6.2 – Classifying the Elements

6.1 – ORGANIZING THE ELEMENTS Scientists knew of thirteen elements up until the 1700s. _________________ improved – better mining tools, refining (____________) Industrial revolution – more exploration, need _____________________ As more elements were discovered, people wanted to know…

6.3 – Periodic Trends Questions to answer in this section • How did chemists begin to organize elements? • How did Mendeleev organize elements? • How is the modern periodic table organized?

• What about earth, air, fire, water theory?

• Why so many elements?

• Were some still undiscovered?

• What makes some similar and some different?

J. W. DOBEREINER – published Law of ___________, showing trends in 3’s. Elements with similar properties, like Cl, Br, I, had one with an atomic mass midway between the other two. Not perfect, but a good start… DMITRI MENDELEEV – given credit as ____________ to develop periodic table in 1869. Made separate element cards with ____________ & ___________; tried for best arrangement. Placed elements in order by increasing _________________. Lothar Meyer – also published PT. Why did Mendeleev get credit? What was so special about Mendeleev’s table? 1. Published first & could explain his _______________ best. 2. Left blank spaces for 3 __________________ elements.

PREDICTED PROPERTIES Mendeleev predicted __________ and _____________ properties of gallium, scandium, germanium. Used avg. values (________________, atomic mass, ______________, melting point, etc) of elements in same group as undiscovered element to make predictions. Scientists knew properties to look for. His predictions turned out to be very close! HENRY MOSELEY – Mendeleev’s table was a breakthrough, but not perfect! There were a few discrepancies when ordered by _____________________ Some properties didn’t line up.

Argon (39.9), Potassium (39.1)

Tellurium (127.6), Iodine (126.9)

1913, Moseley rearranged Mendeleev’s periodic table, basing it on _______________________, not atomic mass. We still use this version. In 1915, killed in WW I at 27. Why didn’t Mendeleev (1869) arrange by atomic number in the first place?

MODERN PERIODIC LAW: When elements are arranged in order of _____________________ atomic number, there is a periodic repetition of their physical and chemical properties. _______________________: The tendency for something to repeat at regular intervals.

6.2 – CLASSIFYING THE ELEMENTS METALS – 80% of periodic table Shiny, ductile (wires)

Questions to answer in this section • What are three broad classes of elements? • What info is displayed on periodic table? • How are groups related to electron configuration?

Good conductors of _________/electricity Mostly _____________, malleable

___________ melting point NONMETALS – ___________ conductors of heat/electricity Mostly _________, some solids, one liquid at RT

Brittle

Lower melting points

METALLOIDS – Properties like metals & nonmetals Different conditions make different ________________

Semiconductors (Si)

•

To locate metals, nonmetals and metalloids, find the _______________.

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Metalloids touch zig-zag, metals ___________, nonmetals ______________

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___________/Families = columns

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Various labeling systems for 18 groups

• _______________ = horizontal rows European System: 1A, 2A, 3A, 4A……1B, 2B, 3B….

U.S. System: 1A, 2A, 3B, 4B, 5B……1B, 2B, 3A, 4A, 5A… We will use 1-2

IUPAC: 1985 decided on #1-18

3-8

GROUP NAMES 1

8 2

3

4

5

6

7

1

8 SECTION NAMES

2

3

4

5

6

7

REPRESENTATIVE ELEMENTS: s- & p-block elements show a _______________________ of properties. Includes metals, nonmetals, metalloids, solids, liquids, gases, reactive and unreactive. Group gets unique properties from electron configuration. Elements in a ___________ all have the same outer e− configuration & the same # of ________________. H – 1s

# of VE = group # O – 1s 2s22p4

1

2

Li – 1s22s1

S – 1s22s22p63s23p4

Na – 1s22s22p63s1

Se – 1s22s22p63s23p64s23d104p4

TRANSITION METALS: d-block elements show some variation due to different numbers of “d” electrons. All have 2 outermost “s” electrons.

V – [Ar]4s23d3

Nb – [Kr]5s24d3

Ta – [Xe]6s24f145d3

INNER TRANSITION METALS: f-block elements show very little variation; have s2 with varying number of “f” electrons.

Nd – [Xe]6s24f4

U – [Rn]7s25f4

6.3 – PERIODIC TRENDS TWO REASONS WHY PROPERTIES CHANGE 1. More

of electrons – as you move DOWN a group,

more inner energy levels are full. They can block or shield the ________________ of the nucleus (+) on outermost e−: _______________________ 2. More protons, same energy level – as you move across a period, more ________ in nucleus and more ____________ in same outermost energy level. Leads to greater ___________ by nucleus on outermost e−.

Questions to answer in this section • What are the trends in atomic size? • How do ions form? What is the trend in ionic size? • What are the trends in first ionization energy and electronegativity? • What causes these trends?

1. ATOMIC RADIUS Atomic radius is half the distance between two bonded nuclei. Atoms get ________________ as you move right across a period. Atoms get ________________ as you move down a group. •

Across Same Period: more p+ makes a stronger pull from the nucleus, e− in ________________ E level are pulled in closer to nucleus.

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Down Same Group: Each element contains electrons in the next higher E level; more inner layers, outermost is __________________ from the nucleus.

2. SIZE OF IONS – SIZE: ATOMS vs. IONS

ATOMIC RADIUS

Generally, ionic sizes follow the same trends as atomic radius. Depends mostly on _________________of electrons •

Loss of electron = _____________ ion - sometimes wipe out an entire _________when valence e− are lost

•

Addition of electron = ______________ ion - More e− without p+ to pull in close

3. FIRST IONIZATION ENERGY Ionization energy (IE) is the amount of ____________ it takes to ____________ an electron. 1st IE is energy to remove 1st electron, 2nd IE is energy to then remove the 2nd electron, and so on. •

Across Same Period: more p+ = stronger pull from nucleus, e− in same E level are pulled in closer to nucleus. 1st IE ___________________ as you move across a period.

•

Down Same Group: Each element contains electrons in the next higher E level; more inner layers, outermost further from nucleus. Pull from nucleus gets weaker. 1st IE _________________as you move down a group.

1ST IONIZATION ENERGY

4. ELECTRONEGATIVITY (ELECTRON AFFINITY) Electronegativity (EN) is a measure of how strongly an atom can ______________ additional electrons when bonded. “_____________________.” We use the Pauling scale of 0 – 4.0, where 4.0 is highest (F). •

Across Same Period: more p+ = stronger pull from nucleus, e− in same E level are pulled in closer to nucleus. EN _______________ as you move across a period.

•

Down Same Group: Each element contains electrons in the next higher E level; more inner layers, outermost further from the nucleus. Pull from

ELECTRONEGATIVITY

nucleus gets weaker. EN __________________ as you move down a group. Highest value wins the “Tug of War.” Electrons end up __________ to element with highest EN. •

Why would Noble Gases not get electronegativity values?

SUMMARY Answer these questions •

How was Mendeleev’s periodic table organized?

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How is the modern periodic table organized?

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On the periodic table what separates metals, nonmetals and metalloids?

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What electron configuration do all halogens end in?

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Which is bigger, an atom of Mg or an atom of Cl?

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What happens to radius when an atom loses an electron?

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Which will be smaller, S atom or S2− ion? K+ or Ca2+?

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What will have a higher 1st ionization energy, Na or Cs?

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Which will have higher electronegativity N or O?

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Briefly say why atomic radius gets smaller as you move across a period.

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Briefly say why S has higher electronegativity than Te.