10/11/12Warm-up #1

List at least 5 ways that the periodic table is a valuable tool.

Daily Objective

Students will begin the unit on the periodic table and periodic law.

Agenda

Discuss Objectives Discuss Warm-up Notes Practice

Unit 4The Periodic

Table

Vocabulary

While taking notes, keep a list of words that you feel might be listed on the test as vocabulary.

How do I write names & symbols?

Spelling COUNTS!! Symbols MUST be written in BLOCK print!! Symbol’s first letter is always uppercase and

the second letter (if one) is always lowercase!!

For Example:Co Cobalt (an element)CO Carbon Monoxide

(compound of carbon and oxygen combined)

NO Scripty / Cursive A, N, H, etc...

l is a lowercase L not an uppercase I

Which elements do I have to know?

ELEMENT SYMBOLHydrogen HHelium HeLithium LiBeryllium BeBoron BCarbon CNitrogen NOxygen O

ELEMENT SYMBOL

Fluorine FNeon NeSodium NaMagnesium MgAluminum AlSilicon SiPhosphorus PSulfur S

Which elements do I have to know?

ELEMENT SYMBOL

Chlorine ClArgon ArPotassium KCalcium CaIron FeCopper CuZinc ZnBromine Br

ELEMENT SYMBOL

Silver AgTin SnIodine IGold AuMercury HgLead PbFrancium FrUranium U

Dmitri Mendeleev Russian Chemist In 1869, he published the first periodic

table. He organized the elements in a way that

would help his students learn them more easily.

He made a card game with the information known about each element listed on separate cards. They could then be arranged by the properties the elements had in common.

Mendeleev settled on an organization of elements that was based on the masses of the elements.

Mendeleev found that when the elements were arranged in order of atomic mass, many physical and chemical properties of the elements followed repeating patterns.

Worked for most elements, but not all. Three pairs of elements had to be

switched, but Mendeleev thought these masses were measured incorrectly.

Ekasilicon

Mendeleev was able to accurately predict the existence of elements not yet discovered. These showed up as gaps in his periodic table.

One such element gap, Mendeleev called ekasilicon. He predicted its mass, density, melting point and color based on its location in the periodic table.

Fifteen years after this prediction, a new element was discovered in Germany and given the name Germanium. Its properties matched the properties of ekasilicon.

H.G.J Moseley worked in Rutherford’s lab. Found that metals produce X-rays when

bombarded with energetic electrons and that the frequencies differed for each metal. These frequencies, came from differences in a fundamental property of each element; the amount of positive charge in the nucleus.

The amount of positive charge = the number of protons = the atomic number.

Modern Periodic Law

When the elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic, repeating pattern.

Periodic Table Arrangement

Periods – the horizontal rows of elements. The modern periodic table has 7 periods. The period tells the number of energy levels used.

This number is called the Principal Quantum Number.

Groups – the vertical columns of elements. The modern periodic table has 18 groups. The “A” groups tell the number of valence electrons. Groups are sometimes referred to as families.

Li = 1s22s1 Na = 1s22s22p63s1

K = 1s22s22p63s23p64s1

*Li, Na, K all have 1 valence e-*All are found in group 1 or IA*Li in pd. 2 Na in pd. 3 K in pd. 4

B = 1s22s22p1 Al = 1s22s22p63s23p1

Ga = 1s22s22p63s23p64s23d104p1

*B, Al, Ga all have 3 valence e- *All are found in group 13 or IIIA

*B in pd. 2 Al in pd. 3 Ga in pd. 4

N = 1s22s22p3 P = 1s22s22p63s23p3

As = 1s22s22p63s23p64s23d104p3

*N, P, As all have 5 valence e-*All are found in group 15 or VA*N in pd. 2 P in pd. 3 As in pd. 4

F = 1s22s22p5 Cl = 1s22s22p63s23p5

Br = 1s22s22p63s23p64s23d104p5

*F, Cl, Br all have 7 valence e-*All are found in group 17 or VIIA*F in pd. 2 Cl in pd. 3 Br in pd. 4

Groups of the Periodic Table

Elements within a group on the periodic table have similar properties to each other.

This is due to the number of valence electrons. Having the same number of valence electrons makes them bond to similar atoms in the same ratios.

If you know the properties of one element in a group, you know the properties of all the elements in that group!!!

Trends of the Periodic Table

Atomic Number Atomic Mass Metal / Nonmetal Trend Atomic Radius Ionization Energy Electronegativity

Atomic Number

The number of protons in one atom of a given element.

Increases as you move down a group.

Increases as you move across a period.

Average Atomic Mass

The average mass of all the isotopes of a given element.

Increases as you move down a group.

Increases as you move across a period.

Atomic Number/Avg.Atomic Mass

increases

I n c r e a s e s

Metal / Nonmetal / Metalloid

Elements on the left side of the table are metals.

There are 88 metals on the periodic table.

Elements on the right side are nonmetals. There are 17 nonmetals on the periodic

table. Elements on the “staircase” between the

metals and nonmetals are metalloids. There are 7 metalloids on the periodic

table.

Properties of Metals and Nonmetals

Property Metals Nonmetals

Color Silver Not Silver

Luster Shiny Dull

Conductivity Conductor Insulator

Malleability Malleable Brittle

Reaction to Acid

Reaction to Acid

No Reaction to Acid

Electron donor or acceptor

Electron Donors

Electron Acceptors

Properties of Metalloids

Metalloids have properties of both metals and nonmetals.

For example: SiliconIs silver, shiny and a conductor like a metal.Is rough, brittle and has no rxn to acid like a nonmetal.

Activity Series Some metals are more reactive than others

and will replace less reactive metals during a reaction.

Li, K, Ca, Na, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, AgDECREASING REACTIVITY

Will Ca replace Zn in a reaction?Yes, Ca is more reactive than Zn

Will Zn replace Mg in a reaction?No, Zn is less reactive than Mg

Metal vs. Nonmetal Trend

Nonmetals Metals Metalloids

Most Active Nonmetal

Most Active Metal

Atomic Radius

The distance from the center of an atom’s nucleus to its outermost electron.

Atomic radius increases as you move down a group on the periodic table.

Because electron clouds are added.

Atomic radius decreases as you move across a period on the periodic table.

Full shells with paired-up electrons have less repulsion, so the take up less space.

Atomic Radius

Fr

Ionization Energy

The amount of energy required to remove a valence electron from an atom.

Decreases as you move down a group on the periodic table.

More “shielding” from full electron shells.

Increases as you move across a period on the periodic table.

More difficult to remove electrons from a full shell.

Electronegativity

The attraction for shared electrons in a chemical bond.

Decreases as you move down a group on the periodic table.

Valence electrons get farther from the nucleus because there are more full e- shells. The attraction between the + nucleus and – electrons decreases.

Increases as you move across a period on the periodic table.

More protons, stronger attraction to electrons.

Electronegativity/Ionization Energy

Look at the positions of Fluorine and Francium on the periodic table.

Francium doesn’t need another electron to become stable. It actually would be more stable by giving one away. (Has lowest ionization energy due to 6 full shells of shielding.)

Electronegativity/Ionization Energy

Fluorine needs one more electron to fill its valence shell. It has the strongest electronegativity. The closer the valence electrons are to the nucleus, the stronger the electronegativity. (The valence e- are closer to the positive nucleus.)

Classwork

What is not finished in class is due tomorrow

Chapter 6 in textbook Section 6.1 Assessment 1-6 Practice problems 8-10 Section 6.2 Assessment 11-15

Warm-up #2

Draw a small periodic table (no elements, just an outline) and label the trends for atomic radius, electronegativity and ionization energy.

Electronegativity and Ionization Energy

Atomic Radius

Objective

Students will continue to learn trends of the periodic table

Agenda

Warm-up Continue with notes about Trends Practice using Periodic Table Puzzle

and Trends WS Wrap-up

Alkali Metals

Group 1 or 1A. These elements have 1 valence

electron which they will give up when bonding to become more stable. (full valence shell) Therefore, they will become +1 ions.

Alkali means “ashes.” Sodium and potassium are present in the ashes of burned plants.

Alkali Metals Alkali metals are very reactive. Will react

with water and air! (Stored in oil.) They are soft enough to be cut by a knife. The members of this group are: (Not

Hydrogen.) Li [He]2s1

Na [Ne]3s1

K [Ar]4s1

Rb [Kr]5s1

Cs [Xe]6s1

Fr [Rn]7s1

Alkaline Earth Metals

Group 2 or 2A. These elements have 2 valence

electrons which they will give up when bonding to become more stable. (full valence shell) Therefore they will become +2 ions.

Obtained from alkaline earths. Earths were substances unchanged by fire.

Alkaline Earth Metals

Alkaline Earth Metals are also very reactive. (Not as reactive as group I or 1A)

They have higher densities and melting points than the alkali metals.

Found in mineral deposits. The members of this group are:

Be [He]2s2 Mg [Ne]3s2

Ca [Ar]4s2 Sr [Kr]5s2

Ba [Xe]6s2 Ra [Rn]7s2

Transition Metals

Groups 3 – 12 Play an important role in living

organisms, are extremely valuable as strong structurally useful materials.

Vary greatly in properties and abundance.

Most have high densities and high melting points.

Number of valence electrons varies, therefore ionic charge varies.

Inner Transition Metals

Lanthanide and Actinide Series elements (“f” block elements)

Properties similar to Transition Metals Many are man-made Many are naturally radioactive (Large,

unstable nuclei. Ratio of neutrons to protons is high)

Have two valence electrons. (after 6s2, 7s2)

The Boron Group Group 13 or 3A. These elements have 3 valence

electrons which they will give when bonding to become more stable. (full valence shell) Therefore they will become +3 ions.

The members of this group are:B [He]2s22p1 Al [Ne]3s23p1

Ga [Ar]4s24p1 In [Kr]5s25p1

Tl [Xe]6s26p1

The Carbon Group Group 14 or 4A. These elements have 4 valence

electrons which they can give up, share or accept to become more stable. (full valence shell) Therefore they can become +4 or –4 ions.

The members of this group are:C [He]2s22p2 Si [Ne]3s23p2

Ge [Ar]4s24p2 Sn [Kr]5s25p2

Pb [Xe]6s26p2

The Nitrogen Group Group 15 or 5A. These elements have 5 valence

electrons. They will accept 3 more to become stable. (full valence shell) Therefore they become –3 ions.

The members of this group are:N [He]2s22p3 P [Ne]3s23p3

As [Ar]4s24p3 Sb [Kr]5s25p3

Bi [Xe]6s26p3

The Oxygen Group Group 16 or 6A. These elements have 6 valence

electrons. They will accept 2 more to become more stable. (full valence shell) Therefore they become –2 ions.

The members of this group are:O [He]2s22p4 S [Ne]3s23p4

Se [Ar]4s24p4 Te [Kr]5s25p4

Po [Xe]6s26p4

The Halogen Group

Group 17 or 7A. These elements have 7 valence

electrons. They will accept 1 more to become more stable. (full valence shell) Therefore they become –1 ions.

Comes from the Greek word which means “salt former.”

The Halogen Group The Halogens are very reactive

nonmetals and exist in elemental form as diatomic molecules.

(F2, Cl2, Br2, I2, At2)

The members of this group are:F [He]2s22p5 Cl [Ne]3s23p5

Br [Ar]4s24p5 I [Kr]5s25p5

At [Xe]6s26p5

The Noble Gases

Group 18 or 8A. These elements have 8 valence

electrons. (Except helium, which has 2) They are already stable, so they do not accept or receive electrons. They do not typically form ions.

These elements are very stable.

The Noble Gases Members of this group include:

Helium [He] 1s2

Neon [Ne] 1s2 2s22p6

Argon [Ar] 1s2 2s22p6 3s23p6

Krypton [Kr] 1s2 2s22p6 3s23p64s23d104p6

Xenon [Xe] 1s2 2s22p6 3s23p64s23d104p6 5s24d105p6

Radon [Rn] 1s2 2s22p6 3s23p64s23d104p6 5s24d105p66s24f145d106p6

Hydrogen (The one and only!)

Has one valence electron, but is not an Alkali metal.

Hydrogen is a nonmetal that exists as a gas under normal conditions.

Colorless, odorless, and composed of H2 molecules. (Diatomic, like the halogens)

Most of the Earth’s hydrogen is combined with oxygen as water and carbon in organic compounds called hydrocarbons.

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Fe Cu Zn Br

Ag Sn I

Au Hg Pb

U

Alkali Metals 1 IA 2 IIA

Alkaline Earth Metals

3 4 75 86 9 10

11

12

Transition Metals

13 IIIA

Boron

Group

14 IVA

15 VA

16 VIACarbon Group

Nitrogen Group

Oxygen Group

17 VIIA

18 VIIIA

Halogens

Noble Gases

Lanthanides

Actinides

Fr

1

2

3

4

5

6

7

1

2

3

4

5

6

7