Chapter 6 The Periodic Table Why are the elements cobalt phosphorus and sulfur essential for...

Chapter 6The Periodic Table

Why are the elements cobalt phosphorus and sulfur essential for maintaining law

and order?

Because when they are put together, they’re CoPS

6.1 Organizing the Elements OBJECTIVES:

• Explain how elements are organized in a periodic table

• Compare early and modern periodic tables

• Identify three broad classes of elements

Section 6.1Organizing the Elements

A few elements, such as gold and copper, known for thousands of years

Yet only about 13 had been identified by the year 1700.

As more elements discovered, chemists realized a way to organize the elements was needed.

6.1 Organizing the Elements Chemists used properties of

elements to sort them into groups. In 1829 J. W. Dobereiner arranged

elements into triads – groups of three elements with similar properties• One element in each triad had

properties intermediate of other two• Worked well for many elements, but not

all

Mendeleev’s Periodic Table By mid-1800s, about 70 elements known Dmitri Mendeleev, a Russian chemist and

teacher, arranged elements in order of increasing atomic mass

Mendeleev noticed that if he grouped elements a certain way, certain chemical characteristics appeared periodically on the table

Thus, the Periodic Table

Mendeleev Mendeleev left “blanks” on his table

where he thought undiscovered elements should fit• When some of these elements were

discovered, many of his predictions proved to be spot on

But, there were problems:• Co and Ni; Ar and K; Te and I• He’d switched them on his table because

he didn’t know about isotopes

A better arrangement Mendeleev’s mistake: arranging elements

based on atomic mass instead of z• so he’d switched Co and Ni because the atomic

mass of Co is slightly higher than Ni, due to Co having heavier isotopes (more neutrons)

In 1913, Henry Moseley, a British physicist, arranged elements according to increasing atomic number

The symbol, atomic number & mass are basic items included on most tables

Spiral Spiral Periodic TablePeriodic Table

The Periodic Law says: When elements are arranged in order of

increasing atomic number, a periodic repetition of physical and chemical properties occurs

Horizontal rows = periods• There are 7 periods

Vertical columns = groups (or families)• Similar physical & chemical props

– Especially main block elements (groups 1A - 8A)

Classes of Elements

1. Metals: electrical conductors, have luster, ductile, malleable

2. Nonmetals: brittle non-lustrous solids, or gases, one liquid (Br), all poor conductors of heat & electricity

3. Metalloids: border zig-zag line on both sides and have properties intermediate between metals and nonmetals

Checking Understanding What did Mendeleev base his table of elements upon? What caused Mendeleev to mistakenly switch Co and

Ni; Ar and K; Te and I on his original table? Complete the following table (Note charges on ions):

Element Metal? Z Mass # neutrons electrons72Ge Metalloid 32 72 40 32

40Ca2+

16O127I−

59Co2+

40Ar

6.2 Classifying the Elements

OBJECTIVES:• Describe the information provided in a

periodic table• Classify elements based on electron

configurations• Distinguish representative elements

(main block) and transition metals

Squares in a Periodic Table The periodic table displays symbols

and names of elements, along with information about the structure of their atoms:• Atomic number and atomic mass• Symbols usually color coded

• solid; gas; liquid

Based on the average atomic mass and the atomic number, what is likely the most common isotope of Na?

Groups of Elements - Families Group 1A – alkali metals

• Forms a “base” (or alkali) when reacting (very violently!) with water

• Na + 2 H2O Na+ + 2 OH− + H2

Group 2A – alkaline earth metals• Also form bases with water• do not dissolve well and melt at very high T,

hence usually solid (“earth”) Group 7A – halo-gens (= “salt-forming”)

Electron Configurations in Groups Elements can be sorted into 4

different groupings based on their electron configurations:

1) Noble gases2) Representative elements3) Transition metals4) Inner transition metals

Let’s now take a closer look at these.

Electron Configurations in Groups1. Main Block Elements = Groups 1A - 8Aa) Representative Elements Groups 1A - 7A

i. Display wide range of properties, thus a good “representative” sample of all matter

ii. There are metals, nonmetals, and metalloids; Many are solid, others are gases or liquids

iii. Their outermost electrons are in s and p subshells that are NOT filled

b) Noble Gases: so called because they rarely take part in a reaction; very stable (exist as free atoms in nature)

i. Noble gases have their outer s and p sublevels completely full (the most stable arrangement)

Electron Configurations in Groups2) Transition (d-Block) metals in the “B”

columns of the U.S. style periodic table• Electron configuration has the outer s

sublevel full (usually), and all or part of d subshell filled

• Block “transitions” between the metal and nonmetal area of table

• Examples include familiar metals: gold, copper, silver, iron

Electron Configurations in Groups3) Inner Transition Metals located below

main body of the table, in two rows• Electron configuration has the outer s

sublevel full, and f sublevel is partly to completely filled

• Formerly called “rare-earth” elements, but this is not true because some are quite abundant

• We won’t be concerned with these much in this class

1A

2A 3A 4A 5A 6A7A

8A Main Block in color Groups 1A-7A are

representative elements

The group B are called the transition elements

These are called the inner transition elements, and they belong here

Group 1A are the alkali metals (but NOT H)

Group 2A are the alkaline earth metalsH

Group 8A are the noble gases Group 7A is called the halogens

1s1

1s22s1

1s22s22p63s1

1s22s22p63s23p64s1

1s22s22p63s23p64s23d104p65s1

1s22s22p63s23p64s23d104p65s24d10

5p66s1

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

H1

Li3

Na11

K19

Rb37

Cs55

Fr87

Do you notice any similarity in the configurations of the alkali metals?

He2

Ne10

Ar18

Kr36

Xe54

Rn86

1s2

1s22s22p6

1s22s22p63s23p6

1s22s22p63s23p64s23d104p6

1s22s22p63s23p64s23d104p65s24d105p6

1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p6

Do you notice any similarity in the configurations of the noble gases?

Alkali metals all end in s1

Alkaline earth metals all end in s2

• really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons.

s2s1 Elements in the s - blocks

He

Transition Metals - d block

d1 d2 d3s1 d5 d5 d6 d7 d8

s1 d10 d10

Note exceptions to electron configuration rules

The P-block

p1 p2 p3 p4 p5 p6

Each row (or period) is the energy level for s and p orbitals (d block is one level lower in each row).

1

2

3

4

5

6

7

Period Number

3d4d

f orbitals start filling at 4f, and are 2 less than the period number

1

2

3

4

5

6

7 4f

5f

6.3 Periodic Trends OBJECTIVES:

• Describe and explain trends among the elements for atomic size

• Explain how ions form• Describe and explain periodic trends

for first ionization energy, ionic size• Describe the periodic trend for

electronegativity

Trends in Atomic Size

How do we measure?• electron clouds of atoms

don’t have a definite edge Chemists get around this by

measuring more than 1 atom at a time

Atomic Size

Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.

}Radius

3 Factors Affect Periodic Trends 1. Energy Level Electrons in higher energy levels are farther away

from the nucleus Each successive level builds on top of lower ones2. Nuclear Charge (# protons) More + charge pulls electrons in closer. (+ and –

attract each other)3. Shielding effect “core” electrons near nucleus shield outer

electrons from protons’ “pull” The higher the energy level greater shielding

#1. Atomic Size - Group trends As we go down a

group, each atom has another energy level (= more electrons, farther from nucleus)

so the atoms get bigger.

HLi

Na

K

Rb

#1. Atomic Size - Period Trends Going from left to right across a period,

the size gets smaller. Electrons are in the same energy level. But, there is more nuclear charge. Outermost electrons are pulled closer.

Na Mg Al Si P S Cl Ar

Atomic Number

Ato

mic

Rad

ius

(pm

)

H

Li

Ne

Ar

10

Na

K

Kr

Rb

3

Period 2

Ions Some compounds are composed of

particles called “ions”• An ion is a atom (or group of atoms)

that has a positive or negative charge Atoms are neutral because the number

of protons equals electrons• Positive and negative ions are formed

when electrons are transferred (lost or gained) between atoms

Ions Metals tend to LOSE electrons, from their

outer energy level• Sodium loses one: there are now more

protons (11) than electrons (10), and thus a positively charged particle is formed = “cation”

• Charge is written as a number followed by a + or – sign: Na1+ or Na+

• Now named a “sodium ion”

Ions Nonmetals tend to GAIN one or

more electrons• Oxygen often gains two electrons• Protons (8) no longer equals

electrons (10), so a charge of -2• O2– is re-named a “oxide ion”• Negative ions are called “anions”

#2. Trends in Ionization Energy Ionization energy is the amount of energy

required to completely remove an electron (from a gaseous atom)

Removing one electron makes a 1+ ion The energy required to remove only the first

electron is called the first ionization energy Energy needed to remove a second electron

= second ionization energy Remove a third = third ionization energy And so on…

Ionization Energy The second ionization energy is

always greater than first IE• There are more protons than

electrons after the first is removed, so the nucleus pulls the rest closer

The third IE (energy to remove a third electron) is even greater• Removing an electron from a 2+ ion

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

Table 6.1, p. 173

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

Why did these values increase so much?

What factors determine IE Greater nuclear charge = greater IE

• More protons, more difficult to remove e-

Greater distance from nucleus decreases IE• Farther from “pull” of protons + shielding

effect of e- in lower energy levels Filled and half-filled orbitals have lower

energy, so achieving them is easier, lower IE.

Shielding Nucleus has weak

attraction to electron in outermost energy level due to repulsive forces of all the other energy levels in between.

Removing a second e– from the same period requires only a little more energy because shielding is the same

Symbol First Second Third

NeNaMgAl

2080 496 738578

39634565 14501816

691277322744

Why did these values increase so much?

The big jumps in IE seen above occur because that second electron removed from Na and the third removed from Mg are being taken from a lower energy level (a core electron rather than a valence electron)

Ionization Energy - Group trends

As you go down a group, the first IE decreases because...• The electron is further away from

the attraction of the nucleus, and• There is more shielding.

Ionization Energy - Period trends All the atoms in the same period have

the same energy level• = same shielding• BUT, increasing nuclear charge

So IE generally increases from left to right

Exceptions occur (of course), but can be explained in light of electron configurations

Trends in Ionic Size: Cations Cations form by losing electrons Cations are smaller than the atom

they came from – not only do they lose electrons, they often lose an entire energy level

Metals form cations Cations of representative elements

have the noble gas configuration before them

Ionic size: Anions Anions form by gaining electrons. Anions are bigger than the atom they

came from – have the same energy level, but more electrons to attract with the same nuclear charge

Nonmetals form anions. Anions of representative elements

generally have a noble gas configuration after they ionize.

Configuration of Ions Main block ions have noble gas

configurations ( = a full outer level) Na atom is: 1s22s22p63s1 Forms a 1+ sodium ion: 1s22s22p6 Same configuration as neon. Metals form ions with the

configuration of the noble gas before them - they lose electrons.

Configuration of Ions Non-metals form ions by gaining

electrons to achieve a noble gas configuration

The ion’s noble gas configuration is the same as the noble gas that comes after them in the periodic table

F = 1s2 2s2 2p5

F– = 1s2 2s2 2p6

Ion Group trends Each step down a

group is adding an energy level

Ions get bigger as you go down, because of the additional energy level• Same as with atoms

Li1+

Na1+

K1+

Rb1+

Cs1+

Ion Period Trends Across the period from left to

right, the nuclear charge increases - so they get smaller.

Notice the energy level changes between anions and cations.

Li1+

Be2+

B3+

C4+

N3- O2- F1-

Size of Isoelectronic ions Iso- means “the same” Isoelectronic ions have the same

# of electrons Al3+ Mg2+ Na1+ Ne F1- O2- and N3-

• all have 10 electrons all have the same configuration:

1s22s22p6 (which is = noble gas neon)

Size of Isoelectronic ions? Positive ions that have more protons

would be smaller (more protons would pull the same # of electrons in closer)

Al3+

Mg2+

Na1+ Ne F1- O2- N3-

13 12 11 10 9 8 7

#3. Trends in Electronegativity Electronegativity is the tendency for

an atom to attract shared electrons when it is chemically combined with another element.

They share electrons, but how equally do they share them?

An element with a high electronegativity means it pulls the electron towards itself strongly!

Electronegativity Group Trend The further down a group, the

farther the electron is away from the nucleus, plus the more electrons an atom has

Valence electrons farther from the nucleus, so less tightly held

Thus, more “willing to share”• = Low electronegativity

Electronegativity Period Trend Metals let their electrons go easily

• Thus, low electronegativity As you move from left to right metallic

character decreases, until you get to the non-metals at the far right of the table

Adding electrons makes non-metals more stable, so they often take them away from other atoms• = High electronegativity

The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions

Atomic size increases in these directions:

Firs

t Ion

izat

ion

ener

gy

Atomic number

He He has a greater IE than H.

Both elements have the same shielding since electrons are only in the first level

But He has a 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 These outweigh

the 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 similar shielding greater nuclear charge But 2p e– removed from B

is higher energy than the 2s removed from Be

This makes up for the greater nuclear charge of B, and thus the IE is slightly lower than Be

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

Pattern broken again, because electron removed from O comes out of a filled p orbital

The added repulsion of another electron in the same orbital makes it easier to remove one

Firs

t Ion

izat

ion

ener

gy

Atomic number

O

F

H

He

Li

Be

B

C

N

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