The how and whyThe how and why

History Russian scientist Dmitri MendeleevDmitri Mendeleev

taught chemistry in terms of properties.

Mid 1800 - molar masses of elements were known.

Wrote down the elements in order of increasing mass (this is wrong!!!!)

Found a pattern of repeating properties.

Mendeleev’s Table Grouped elements in columns by similar

properties in order of increasing atomic mass.

Found some inconsistencies - felt that the properties were more important than the mass, so switched order.

Found some gaps. Must be undiscovered elements. Predicted their properties before they

were found.

The modern table Elements are still grouped by

properties. Similar properties are in the same

column. Order is in increasing atomic number. Added a column of elements Mendeleev

didn’t know about. The noble gases weren’t found because

they didn’t react with anything.

Horizontal rows are called Periods There are 7 periods

Vertical columns are called Groups.

Elements are placed in columns by similar properties.

Also called families

1A

2A 3A 4A 5A 6A7A

8A0

The elements in the A groups are called the representative elements

The group B are called the transition elements

These are called the inner transition elements and they belong here

Periodicity Explained Valence electron cloud Outside orbitals The orbitals fill up in a regular

pattern The outside orbital electron

configuration repeats The properties of atoms repeat when

placed in order of Atomic Number

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10 5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s2

4f145d106p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

He2

Ne10

Ar18

Kr36

Xe54

Rn86

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s22s22p63s23p64s23d104p65s24d105p6

1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p6

Alkali metals all end in s1

Alkaline earth metals all end in s2

really have to include He but it fits better later.

He has the properties of the noble gases.

s2s1 s- block

Transition Metals -d block

d1 d2 d3 d5 d5 d6 d7 d8 d9 d10

The p-block p1 p2 p3 p4 p5 p6

f - block inner transition elements

f1 f5f2 f3 f4

f6 f7 f8 f9 f10 f11 f12 f14

f13

Atomic Size First problem where do you start

measuring. The electron cloud doesn’t have a

definite edge. They get around this by measuring

more than 1 atom at a time.

Atomic Size

Atomic Radius = half the distance between two nuclei of a diatomic molecule.

}Radius

Trends in Atomic Size Influenced by two factors. Energy Level Higher energy level - electrons

are further away from nucleus. Charge on nucleus – Zeff

More charge (# of protons) pulls electrons in closer.

Group trends As we go down a

group Each atom has

another energy level

So the atoms get bigger.

HLi

Na

K

Rb

Periodic Trends As you go across a period the radius

gets smaller. Same energy level. More nuclear charge = more protons. Outermost electrons are closer.

Na Mg Al Si P S Cl Ar

Atomic Number

Ato

mic

Rad

ius

(nm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

Ionization Energy The amount of energy required to

completely remove an electron from a gaseous atom.

Removing one electron makes a

+1 ion. The energy required is called the first

ionization energy.

Ionization Energy The second ionization energy is the

energy required to remove the second electron.

Always greater than first IE. The third IE is the energy required to

remove a third electron.

IE1<IE2<IE3

Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11810 14840 3569 4619 4577 5301 6045 6276

Symbol First Second ThirdHHeLiBeBCNO F Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11810 14840 3569 4619 4577 5301 6045 6276

What determines IE nuclear charge = IE distance from nucleus = IE Filled and half filled orbitals have

lower energy, so achieving them is easier, lower IE.

Shielding

Periodic trends: Ionization Energy (IE)

All the atoms in the same period have the same energy level.

Same energy level = same shielding, but each atom gains a proton, therefore….

Increasing nuclear charge …helps pull e- in tighter, therefore it is harder to remove.

So IE generally increases from left to right. Exceptions at full and 1/2 fill orbitals.

e-

e-

Periodic trends: Ionization Energy (IE)

As you go down a group first IE decreases because….

The electron is further away…and There are energy levels between

the nucleus and the e- thus…. Shielding the e- from + nucleus.

e-

e-

Shielding The electron on the

outside energy level has to look through all the other energy levels to see the nucleus

Shielding The electron on the

outside energy level has to look through all the other energy levels to see the nucleus.

A second electron has the same shielding.

Effective Nuclear Charge, Zeff

AtomZeff Experienced by Electrons in Valence Orbitals

Li +1.28 Be ------- B +2.58 C +3.22 N +3.85 O +4.49 F +5.13

Increase in Increase in Z* across a Z* across a periodperiod

Firs

t Ion

izat

ion

ener

gy

Atomic number

He

He has a greater IE than H.

same shielding greater nuclear

charge

H

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li has lower IE than H

more shielding further away outweighs greater

nuclear charge

Li

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be has higher IE than Li

same shielding greater nuclear

charge

Li

Be

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He B has lower IE than

Be same shielding greater nuclear

charge By removing an

electron we make s orbital half filled Li

Be

B

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

Breaks the pattern because removing an electron gets to 1/2 filled p orbital

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Ne has a lower

IE than He Both are full, Ne has more

shielding Greater distance

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Na has a lower

IE than Li Both are s1

Na has more shielding

Greater distance

Na

Firs

t Ion

izat

ion

ener

gy

Atomic number

Electron Affinity The energy change associated with

adding an electron to a gaseous atom.

Easiest to add to group 7A. Gets them to full energy level. Increase from left to right atoms

become smaller, with greater nuclear charge.

Decrease as we go down a group.

Periodic trends: Electron Affinity(EA)e-

e-

• From left to right atoms become smaller, with greater nuclear charge.

e - are attracted by the + charged nucleus. •therefore EA will increase from left to right

Periodic trends: Electron Affinity(EA)e-

e-

• Down a group atoms become larger, and have greater nuclear charge. • e - are attracted by the increased +charge but shielding also increases …which has a greater influence!! • therefore EA will decrease as you go down a group or family

• The energy change associated with adding an electron to a gaseous atom.

• e - are attracted by a + charge.– Therefore more protons more attraction– But remember “shielding”

• Easiest to add to group 7A.• Gets them to full energy level.• Increase from left to right atoms become smaller,

with greater nuclear charge.• Decrease as we go down a group.

Periodic trends: Electron Affinity(EA)

Ionization Energy & Electron AffinityEffective Nuclear Charge INCREASE

Atomic Radius & Shielding INCREASES

Size of Isoelectronic ions Iso - same Iso electronic ions have the same #

of electrons Al+3, Mg+2, Na+1, Ne , F-1, O-2 and N-3 all have 10 electrons all have the configuration 1s12s22p6

Size of Isoelectronic ions Positvie ions have more protons so

they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

• Na2O(s) + H2O(l) 2 NaOH(aq)

• CaO(s) + H2O(l) Ca(OH)2 (aq)

• MgO(s) + 2HCl(aq) MgCl2(aq) + H2O (l)

• NiO(s) + H2SO4 (aq) NiSO4 (aq) + H2O (l)

Metal Oxides are BASIC

• CO2(g) + H2O(l) H2CO3(aq)

• P4O10(s) + 6 H2O(l) 4 H3PO4 (aq)

• CO2(g) + 2 NaOH (aq) Na2CO3(aq) + H2O (l)

• SO3(g) + 2 KOH (aq) K2SO4 (aq) + H2O (l)

Non-Metal Oxides are ACIDIC