Unit 7 Periodic Table for website.pptx

Periodic Table

AP Chemistry, Unit 7

7.1: explain how the trends in the Periodic Table

are influenced by effective nuclear charge,

including: atomic/ionic radii, ionization energy, and

electron affinity.

7.2: identify the groups/ regions of the P.T. with

Development of the Periodic

Table

• _{In 1869 Dmitri Mendeleev originally arranged the periodic}

table according to atomic mass and element properties.

• _{Strangely enough, it’s been shown that Lothar Meyer}

conceived of the periodic table around the EXACT same time period that Mendeleev did…

• _{Mendeleev was able to make predictions about missing}

(undiscovered) elements and their properties.

• _{Modern P Table is now organized by:}

-Atomic Number (# of P+)

-Electron configurations (s, p, d, f –blocks)

•

1869

Development of Periodic Table

Mendeleev, for instance, predicted the

Development of Periodic Table

•

Mendeleev and

Meyer’s periodic

tables had only a

couple dozen

elements.

•

The remaining

elements were

discovered by

Must see...

• https://www.youtube.com/watch?v=AcS3NOQnsQM

• _{A Video that is a right of passage for all chemistry}

Effective Nuclear Charge

• Almost all periodic table properties and trends correlate to the ENC

experienced by an elements valence electrons.

• In a multi-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. • The nuclear charge that an

electron experiences (Effective Nuclear Charge) depends on two factors:

-How many protons in the nucleus -How many core electrons are

Effective Nuclear Charge

• _{The effective nuclear charge, Zeff,}

can be mathematically approximated as:

Z_eff = Z − S

• _where

Z is the atomic number = # protons

Now you try…



• Calculate the effective nuclear charge of a valence electron in sodium, magnesium, phosphorous, chlorine, and argon.

Z_eff = Z − S

Na = 11 – 10 = +1 Mg = 12 – 10 = +2 P = 15 – 10 = +5 Cl = 17 – 10 = +7 Ar = 18 – 10 = +8

So, ENC increases from left to right across periodic table and increases slightly from bottom to top.

Notice that the # of

core electrons remains constant (at -10) as

the # of protons

increases across a row.

NOTE: You would not notice a mathematical change in zeff down a column but it does

Group

Properties of Metal, Nonmetals,

and Metalloids

Valence e- in metals

experience low ENC so they can give up

electrons easily

Valence e- in

non-metals experience high ENC so the electrons

cannot leave easily.

They both want an

octet... So metals tend to give electrons to

Metals versus Nonmetals

Thus,

• _{Metals tend to form cations.}

Metals

• _{Lustrous (shiny)}

• Malleable (will bend on impact, not break)

• _{Ductile (can be made into wire)} • _{Good conductors of heat and}

electricity.

• Solid at room temp (except Hg)

• _{Metal oxides tend to be basic}

Nonmetals

Tend to be:

• Dull (not shiny)

• Brittle (break/crack/crumble on impact)

• Poor conductors of heat and electricity

• Tend to be liquid/gas at room temp

Metalloids

• _{Have some}

characteristics of metals, some of nonmetals.

• _{For instance, silicon looks}

Group IA: Alkali Metals

• _{Name comes from Arabic word}

for ashes.

• EXTREMELY REACTIVE

• _{Found only as compounds in}

nature. (too reactive to exist)

• _{Have low densities and melting}

points.

• _{Also have low ENC valence}

Alkali Metals

Their reactions with water are famously exothermic and also produce flammable hydrogen gas…

Excess heat + hydrogen = BOOM!

http://

Alkali Metals

• _{Alkali metals (except Li) react with}

oxygen to form peroxides.

• K, Rb, and Cs also form superoxides: K + O₂  KO₂

Group IIA: Alkaline Earth Metals

• _{Harder metals (but still softer than transition metals)} • _{Have higher densities and melting points than alkali}

metals.

Alkaline Earth Metals

water, Mg reacts only with steam, but others react readily with water.

• _{Will also oxidize when}

heated.

• _{Reactivity tends to}

Group VIIA: Halogens

• _{Prototypical nonmetals (display non-metal properties}

well)

Group VIIA: Halogens

• EXTREMELY REACTIVE

• LARGE ENC

• _{Therefore, tend to oxidize (take}

e- from) other elements easily

• React directly with metals to form metal halides

• “Putrid” colors of green,

brown, purple, and so forth.

• With 7 valence e-, forms

Group VIIIA: Noble Gases

• _{HIGHEST ENC in valence electron (+8)} • _{Complete octet = very inert (unreactive)} • Monatomic gases at room temp

• _{Generate visible spectra light when}

Group VIIIA: Noble Gases

compounds:

• _XeF2

• _{XeF4 (at right)} • _XeF6

• Kr forms only one stable compound:

• _KrF2

• _{The unstable HArF}

Sizes of Atoms

•

Determining the exact

radius of an atoms

electron cloud can be

challenging…

•

The bonding atomic

radius is defined as

one-half of the distance

between covalently

Sizes of Atoms

Bonding atomic radius tends to…

• Decrease (get smaller) from left to right across a row

• _{Due to increasing Zeff, e-}

cloud is pulled in closer.

• Decreases from bottom to top of a column. Again, due to slight increase in Z_eff

• _{E.g. the higher the Zeff the}

Sizes of Ions

• _{Ionic size depends upon:}

• _{Nuclear charge.}

• Number of electrons.

• _{Orbitals in which electrons}

Sizes of Ions

•

Cations are smaller

than their parent

atoms.

•

The outermost

electrons are removed.

•

The valence energy

level is gone

•

The core electron level

Sizes of Ions

•

Anions are larger than

their parent atoms.

• Electrons are added to valence shell.

• Repulsion of valence shell with core shell dominates.

Sizes of Ions

•

Ions increase in size as

you go down a column.

•

Due to increasing

numbers of energy levels.

•

Same as non-ions

• _{Isoelectric series: a series of ions with the same number of}

electrons.

Ionization Energy

•

Amount of energy required to remove an

electron from the ground state of a

gaseous atom or ion.

•

First ionization energy is the energy required

to

remove first electron from the valence

energy level.

•

Second ionization energy is that energy

Ionization Energy

• It requires more energy to remove each successive electron.

• _{When all valence electrons have been}

Trends in First Ionization Energies

•

As one goes down a

column, less energy is

required to remove the

first electron.

•

For atoms in the same

group,

Z

is essentially

the same, but the

valence electrons are

farther from the

Trends in First Ionization Energies

•

Generally, as one

goes across a row, it

gets harder to

remove an electron.

•

As you go from left

to right,

Z

increases so the

valence e- are held

tighter by the

Trends in First Ionization Energies

However, there are

two apparent

discontinuities in

each of these

Trends in First Ionization Energies

•

The first occurs

between Groups IIA

and IIIA. Why?

•

Electron removed

from

p

-orbital rather

than

s

-orbital

•

Electron farther from

nucleus on average

•

Small amount of

repulsion by

s

Trends in First Ionization Energies

•

The second occurs

between Groups VA

and VIA. Why?

•

Electron removed

comes from doubly

occupied orbital.

•

Repulsion from

Trends in Electron Affinity

• _{Energy RELEASED}

accompanying

addition of electron to gaseous atom

• _{Cl(g) + e}−  Cl (g)

-• _{In general, electron}

affinity becomes

more exothermic as you go from left to right across a row. • So they release

more energy and are more stable which makes

Trends in Electron Affinity

There are

again,

however, two

discontinuitie

s in this

Trends in Electron Affinity

• The first occurs between

Groups IA and IIA.

• Added electron must go

in p-orbital, not s-orbital.

• Electrons must pair in

S-orbital for group 2 metals. This is less stable due to 1st e-

repelling.

P orbitals

Trends in Electron Affinity

between Groups IVA and VA.

• Group VA has no empty orbitals.

• _{Extra electron must}

go into occupied orbital, creating repulsion.

P orbitals

Easy way(s) to know trends!!!

• _{Know the two}

“F”s

• _{Remember Flie}

https://www.youtube.com/watch?v=rJHaSL6d2Mc