topic 3: periodicity. history during the 19 th century, it was becoming apparent that some elements...

TOPIC 3:PERIODICITY

History

During the 19th century, it was becoming apparent that some elements displayed similar properties, and some patterns were beginning to emerge.

In 1869, Dmitri Mendeleev of Russia and Lothar Meyer in Germany published nearly identical tables.

Mendeleev is given credit for the modern periodic table since he advanced his ideas more vigorously.

History (cont.)

He even proposed that certain elements were missing, and even predicted what the properties of these missing elements would be! After the discovery of germanium and gallium, it was found that Mendeleev was correct!

IB Core Objective

3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.

Describe: Give a detailed account.

3.1.1 Describe the arrangement of elements in the periodic table in order of increasing atomic number.

Elements are arranged from left to right in order of increasing atomic number.

IB Core Objective

3.1.2 Distinguish between the terms group and period.

Distinguish: Give the differences between two or more different items.

Periods

Gro

up

s o

r Fam

ilie

s



Group is a vertical column consisting of elements with the same number of valence electrons.

Period is a horizontal row where atomic number increases and chemical properties gradually change from left to right: From metals (except hydrogen), to metalloids, then non metals, and on the far right are the noble gases.

Review: Valence electrons are electrons in the outermost energy level (orbital) of an atom.

IB Core Objective

3.1.3 Apply the relationship between the electron configuration of elements and their position in the periodic table up to Z = 20.

Apply: Use an idea, equation, principle, theory or law in a new situation.


What do you remember from Topic 2 on electron configuration as you move across a period on the periodic table?

Each period begins with the first element having one electron in a new main energy level.

Period 1 fills the n=1 level, period 2 fills the n=2 level, etc.

What about the group number?


The group number gives the number of electrons in the valence shell.

For example, electron configuration in group 1:

H: 1 Li: 2,1 Na: 2,8,1 K: 2,8,8,1

They all have one valence electron! Group 1=1 valence electron Group 2? Group 2=2 valence electrons


Why is knowing the electron configuration (and valence electrons) important?

Because I said it’s important.

Also, it is the valence electrons that take part in chemical reactions and determine the physical and chemical properties for an element.

IB Core Objective

3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

Apply: Use an idea, equation, principle, theory or law in a new situation.

3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.

Look at the periodic table for trends regarding number of electrons in the highest occupied energy level.

3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its

position in the periodic table.

What energy level is sulfur (S) in?A: n=3How many valence electrons does it have?A: 6

Speed GameValence Electrons and Naming

IB Core Objective

3.2.1 Define the terms first ionization energy and electronegativity.

Define: Give the precise meaning of a word, phrase or physical quantity.


Ionization energy: Taught to you by the friendly folks from HL!

ReviewIonization Energy

Energy required to rip off one electron

from an element in its gaseous state.X(g) X+

(g) + e-


ElectronegativityHow strongly the atom attracts electrons

in a covalent bond.Mr. F can really attract the electrons. Mr. K, not so much

IB Core Objective

3.2.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across period 3.



Use the Excel spreadsheets you have created on atomic radii and ionization energy to help determine these trends.

Break into small groups. HL students—using the graphs, explain the trends in ionization energy to the SL students.

Why is sodium the lowest in period 3? Why is aluminum lower than magnesium in period

3? Why is the ionization energy so high for argon in

period 3?


Look at the atomic radii graphs you have made. What trends do you notice?

Does the atomic radii go up as the atomic number goes up?

Why do you think these trends exist?


When going across a period on the periodic table, the number of protons in the nucleus increases.

What happens to the charge on the nucleus when protons increase (not taking into account the electrons)?

A: It increases When the nucleus is more positively

charged, the electrons become more attracted.

Attraction moves them closer, decreasing the atomic radii.

+ -

When you’re

positive, I feel so much

closer to you.


ElectronegativityGoing across period 3, electronegativity

increases.

The elements further to the right have a higher affinity for electrons—the ones on the left tend to lose electrons.

IB Core Objective

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li → Cs) and the halogens (F → I).

Describe: Give a detailed account. Explain: Give a detailed account of

causes, reasons or mechanisms.


What do you think would happen if an electron was lost, like in one of the alkali metals?

A: It would lose a shell, so the ionic radius would decrease.

What would happen if an electron was gained in one of the halogens?

A: The extra electron would cause repulsion, making the ionic radius bigger.


QuestionsWhat happens to the

atomic radii as you move down the alkali metals or halogens?

A: It increases.


QuestionsWhat happens to the ionic radii when the

alkali metals become ions?A: They become smaller.

What happens to the ionic radii when the halogens become ions?

A: They become bigger.


QuestionsWhat happens to the first ionization

energy as you go down the alkali metal or halogen group?

A: It goes down


ElectronegativityWhen going down the periodic table,

electronegativity decreases.


Melting PointsWhat causes melting?The attractive forces holding the particles

in the structure of a solid are overcome, and the particles are then free to move around.

The melting point depends on the strength of these attractive forces and the way in which particles are arranged in the solid state.


Melting PointMelting point decreases going down the alkali

metals on the periodic table.Melting point decreases because as the atoms get larger, the forces of attraction between them decrease.

Melting point increases going down the halogens on the periodic table. Why is this?

The solids of halogens are non-polar diatomic molecules, and are thus weakly attracted to each other by van der Waals’ forces. The forces of attraction in van der Waals’ increase as the mass of the molecules increase.

IB Core Objective

3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.

Compare: Give an account of similarities and difference between two (or more) items, referring to both (all) of them throughout.


Which is more electronegative? Sodium or potassium? A: Sodium

Which is more electronegative? Chlorine or bromine? A: Chlorine

Which is more electronegative? Bromine or nitrogen? A: Nitrogen This is a little harder to figure out using just the trends,

but there’s a key word you can learn….


FONClBrISCHP !!A fun way to remember the order of electronegativity!

Which is more electronegative? Bromine or sulfur?

A: Bromine!

IB Core Objective

3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.

Discuss: Give an account including, where possible, a range of arguments for and against the relative importance of various factors, or comparisons of alternative hypotheses.


Alkali metals• Low melting points.• Soft and malleable.• Low densities (why are the densities low?)• A: They are the largest atoms in their period of

the periodic table.• Alkali metals are chemically very reactive.

Why?• A: The outer electron is very easily lost, and

combine easily with reactive non-metals.• So what do you think might react easily with

alkali metals?


All of the metals react with water to form a metal hydroxide and hydrogen gas.

2M(s) + 2H2O(l) → 2M+(aq) + 2OH-

(aq) + H2(g)

Lithium: The reaction occurs slowly and steadily. Sodium: The reaction is vigorous, producing

enough heat to melt the sodium (98°C) Potassium: Reaction is violent and the heat

produced is enough to ignite the hydrogen gas involved.

The solution becomes strongly alkaline (basic) after these reactions.

Metal Sodium (Na) in Water

Metal Potassium (K) in Water

Large Quantity of NaLarge Amount of Na Larger amount of Na


Halogens• Very reactive non-metals.• How many electrons do they require to

fill their valence shell?• A: one• They exist as diatomic molecules. What

are the diatomic molecules in the halogen group?

• A: F2, Cl2, Br2, I2


What do you remember about fluorine (or Mr. F) when it comes to attracting electrons?

It has a strong electron affinity, or it pulls in electrons strongly. (High electronegativity)

Fluorine is the strongest oxidizing agent known (another way to say it easily takes electrons).

It is so reactive that its use is not allowed in most school laboratories.


The ability to oxidize decreases as you go down Group 7.

For example, if one reacted aqueous chlorine with sodium bromide, the chlorine would take the electron:

Cl2(aq) + 2Br-(aq) → 2Cl-(aq) + Br2(aq)

Chlorine is also very reactive, and was used as a poisonous gas in World War I.


Halogens are only slightly soluble in water since they are non-polar.

In aqueous solutions halogens dissociate slightly to form an acidic solution:Cl2(aq) + H2O(l) → H+

(aq) + Cl-(aq) + HOCl(aq)


Halogens will bond with metals to give ionically bonded salts.

2K(s) + Br2(l) → 2KBr(s)

2Na(s) + Cl2(g) → 2NaCl(s)

These salts are usually white and soluble in water (exceptions are lead and silver halide salts, which are insoluble).

IB Core Objective

3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.

Discuss: Give an account including, where possible, a range of arguments for and against the relative importance of various factors, or comparisons of alternative hypotheses.


Look at the first two elements (Na & Mg). What is the ionization energy compared to elements further to the right?

A: They have relatively low ionization energies.

What kinds of bonds would these elements typically form?

A: Ionic What is an oxide? A compound with oxygen bonded with at

least one other element.


So oxides with metals are ionic bonds. Since the oxide ion can form a bond with

hydrogen ions, they act as bases in water. Na2O(s) + H2O(l) → 2Na+

(aq) + 2OH-(aq)

The aluminum oxide is amphoteric. What do you think amphoteric means?

It can act as either an acid or a base. Al2O3(s) + 6H+

(aq) → 2Al3+(aq) + 3H2O(l)

Al2O3(s) + 2OH-(aq) + 3H2O(l) → 2Al(OH)4

-(aq)


Moving to the right, we reach silicon. Silicon is a metalloid. What is a metalloid?

A: Possesses properties of a metal and a non-metal.

The ionization energy for silicon becomes too great for it to effectively form an ion. So what kind of bond would it form?

A: Covalent It is slightly acidic in solution.


The remaining elements form covalent bonds (except argon), and form acidic solutions in water.

For example, sulfur trioxide reacts to form sulfuric acid, a strong acid:

SO3(s) + H2O(l) → H+(aq) + HSO4

-(aq)


Two other equations you need to know… Magnesium oxide…what do you think the

equation would be if it reacts with water?MgO(s) + H2O(l) → Mg(OH)2(aq)/(s)

Acidic or basic in solution?A: Basic

What about phosphorous pentoxide in water?P4O10(s) + 6H2O(l) → 4H3PO4(aq)

Acidic in solution


To summarize the trend across period 3: Ionic Covalent Very Basic weaker basic amphoteric

weakly acidic Very acidic Na2O + H2O 2NaOH (react with water) Na2O + HCl Na+ + Cl- + H2O (react with acid)

Al2O3 is amphoteric

IB HL Objective

13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and structure.

Explain: Give a detailed account of causes, reasons or mechanisms.


Can you deduce what the oxides would be for the first four elements across period 3?

A: Na2O, MgO, Al2O3, SiO2

What about phosphorous? Phosphorous is….different. You will learn down the

road that phosphorous can form five covalent bonds, and can have different oxidation states (Topic 9).

The two oxides for phosphorous you will need to know is P4O6 and P4O10.

The rest are SO2, SO3, Cl2O and Cl2O7. What is the general trend for number of oxides

bonding as you move from left to right on period 3?


As number of valence electrons increases, there is an increase in the number of electrons available for bond formation.

Thus the number of oxides goes up.

Think about the electronegativity differences across period 3 compared to oxygen. What is the trend?


The metals on the left (Na, Mg, and Al) all have a larger electronegativity difference with oxygen.

Oxygen likes to attract their electrons. So oxygen becomes negatively charged,

and the metals become positively charged.

So they form an ionic bond. This causes them to have high melting

points, and they can conduct electricity when molten.


SiO2 (silicon dioxide) also has a high melting point, because it forms a complex lattice with strong covalent bonds.

Silicon dioxide is also commonly called silica. Do you know what this is used for?

http://en.wikipedia.org/wiki/File:Glass_tetrahedon.png

http://en.wikipedia.org/wiki/File:Silica.svg


Oxides of phosphorous, sulfur and chlorine form covalent bonds because the electronegativity difference is small.

So they need to share. They have a low melting point. Low boiling point. And don’t conduct electricity.


ChloridesHow would you write the first three elements

bonded with chlorine?A: NaCl, MgCl2, Al2Cl6NaCl and MgCl2 are ionically bonded, solids at

room temperature and have high melting points.When molten or added to water, they conduct

electricity.Al2Cl6 is a poor conductor of electricity and has a

much lower melting point than NaCl and MgCl2.


How would you write the chlorides for the next three elements across period 3?

A: SiCl4, PCl5, PCl3, and Cl2 Chlorine usually forms only one bond, so

the same extended lattice with Si is not the same as with oxides.

All of these molecules are covalently bonded and have low melting and boiling points.

IB HL Objective

13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.



NaCl added to water gives a neutral solution with a pH of 7.

MgCl gives a slightly acidic solution with water.

When AlCl3 is added to water, a very exothermic reaction takes place, and hydrochloric acid is formed.

2AlCl3(s) + 3H2O(l) → Al2O3(s) + 6HCl(aq)


The other chlorides also react vigorously with water to produce acidic solutions of hydrochloric acid.

SiCl4(l) + 4H2O(l) → Si(OH)4(aq) + 4HCl(aq)

IB HL Objective

13.2.1 List the characteristic properties of transition elements.

List: Give a sequence of names or other brief answers with no explanation.


Which subshell block are the transition elements located?

A: d-block Transition elements are those in which

the element has a partially filled d-sublevel.


d-block contain all dense, hard metallic elements.

The oxidation state can be variable. Because of their small size, transition

metals can attract species that are rich in electrons, and form complex ions.

Compounds of transition metals are colored.

Transition elements have catalytic properties.

IB HL Objective

13.2.2 Explain why Sc and Zn are not considered to be transition elements.


13.2.2 Explain why Sc and Zn are not considered to be transition elements.

To be a transition metal, it must have a partially filled d suborbital in one of the common oxidation states (ion)

Sc: [Ar] 3d 4s2 Sc3+ [Ar] (no d-suborbital)

Zn: [Ar] 3d10 4s2 Zn+2 [Ar] 3d10 (d-suborbital is not partially full)

IB HL Objective

13.2.3 Explain the existence of variable oxidation number in ions of transition elements.



The difference in energy levels between 3d and 4s is small Typically d-block will form +2 ions due to

removal of e- from the 4s instead of 3d. Oxidation numbers can also be predicted

based on losing all the electrons in 4s and 3d.

Sc: [Ar] 4s23d1 Sc3+

Ti: [Ar] 4s23d2 Ti4+

Fe: [Ar] 4s23d6 Fe3+ and Fe2+

Stable due to half filled d-

subshell


QuestionIf iron (III) formed an oxide, what would

the formula be?A: Fe2O3


Question: What are the 2 exceptions to the electron configuration rules for non-ions in the D-block?

A: Chromium and copper. Chromium: [Ar] 4s13d5

Copper: [Ar]4s13d10

It is more favorable for them to half-fill and completely fill the 3d sub-level than spin-pair two in the 4s orbital.


QuestionIf copper lost the 4s electron, what would

the oxidation state be?A: Cu+1

What would the formula be if copper(I) formed an oxide?

A: Cu2O


Metals with 4+ oxidation stateMetals with a 4+ oxidation state are rare.For example, Mn4+ is small and highly charged…it

excites the oxygen ligands into a vacant d orbital into the manganese.

Therefore, the bonding is closer to covalent!MnO2

This complex ion will then ionically bond with a metal, such as potassium, creating….

KMnO4


Why is the oxidation state of Mn4+ rare? Look at the electronic configuration! [Ar]3d3

Not a very stable formation! What do you think another oxidation state for

Mn would be? Mn7+….lose all of the electrons! High

oxidation number equals not very stable state. Bonds with 4 oxygen to make MnO4

-. This complex ion then bonds with potassium to make KMnO4


What if chromium lost all of it’s electrons? Then it would be Cr6+. Another oxidation state for chromium is Cr3+. If chromium (III) bonded with chloride, what

would the formula be? A: CrCl3 If chromium (VI) bonded with oxygen, what

would the formula be? A: Cr2O7

2-

IB HL Objective

13.2.4 Define the term ligand.

Define: Give the precise meaning of a word, phrase or physical quantity.


Ligands: neutral molecules or negative ions that contain a non-bonding pair of electrons.

In an ordinary covalent bond, each atom will share an electron.

Because of their small size, transition metal ions attract species that are rich in electrons.

Therefore, a ligand and the metal form a dative covalent bond where the ligand provides both of the electrons.

H

H

NH

Lone pair of e-

2+Cu

H

HNH

H

H

N H

H

H NH

2+


IB HL Objective

13.2.5 Describe and explain the formation of complexes of d-block elements.




Coordination Numbers # of ligands formed around a central atom

6 ligands: Form octahedral structure 4 ligands: Form tetrahedral 2 ligands: Form linear structure

Complex ions can form (+) Cations (NH4)+

(-) Anions (OH)-

http://en.wikibooks.org/w/index.php?title=File:Ferricyanide-3D-balls.png&filetimestamp=20070423162831


Water is a common ligand, and will form hexahydrates with most transition metals:

[Fe(H2O)6]3+

Cyanide (CN-), also forms octahedral structures with transition metals:

[Fe(CN)6]3-

What is the oxidation state for the iron in both of these complex ions?


Chloride is another example of a ligand. It can form a tetrahedral structure with

copper(II). So what is the formula for this complex ion?

A: [CuCl4]2-

Ammonia is a neutral ligand that can form linear structures with silver(I). What would the formula be for this complex ion?

A: [Ag(NH3)2]+

13.2.5 Describe and explain the formation of complexes of d-block elements. Ligands can often replace other ligands: Ex: [CuCl4]2- + 4H2O ↔ [Cu(H2O)4]2+ +

4Cl- The color will then go from yellow/green

to pale blue. [Cu(H2O)4]2+ + 4NH3 ↔ [Cu(NH3)4]2+ +

4H2O The color will go from pale blue to deep

blue.

IB HL Objective

13.2.6 Explain why some complexes of d-block elements are coloured.



If an atom is all alone, all orbitals have equal energy….BUT

If the atom is surrounded by charged or polar molecules (ligands), the effect of the electric field from these ligands have a different effect on the d-orbitals.

For example, if six ligands bond to a transition metal ion to form an octahedral complex, the 3d sub-level will be split into two.


Cu

2+

H

H

O

HH

O

HH

O

H

H

O

H

H O

H

HO

Equal energy

Energy Gap corresponds

to visible light

Un-equal energy


What causes color? Wavelengths of light. Some of the

wavelengths are being absorbed from white light (promoting the electron to a higher energy level).

The colors that you see are the wavelengths which are not being absorbed (they are transmitted back).

If there are no electrons in the d-orbitals, then it will appear colorless (i.e. Sc3+ and Ti4+)

topic 3: periodicity. history during the 19 th century, it was becoming apparent that some elements...

Documents

group number

number of electrons

number of valence electrons

arrangement of elements

atomic number

missing elements

modern periodic table

valence electrons important